science of everyday things vol. 1 chemistry

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SCIENCE EVERYDAY

THINGS

OF

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SCIENCE EVERYDAY

THINGS

OF

volume 1: REAL-LIFE CHEMISTRY

edited by NEIL SCHLAGERwritten by JUDSON KNIGHT

A SCHLAGER INFORMATION GROUP BOOK

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SCIENCE OF EVERYDAY THINGSVOLUME 1 Real-Life chemistry

A Schlager Information Group BookNeil Schlager, EditorWritten by Judson Knight

Gale Group StaffKimberley A. McGrath, Senior Editor

Maria Franklin, Permissions ManagerMargaret A. Chamberlain, Permissions SpecialistShalice Shah-Caldwell, Permissions Associate

Mary Beth Trimper, Manager, Composition and Electronic PrepressEvi Seoud, Assistant Manager, Composition and Electronic Prepress

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While every effort has been made to ensure the reliability of the information presented in this publication, Gale Group does notguarantee the accuracy of the data contained herein. Gale accepts no payment for listing, and inclusion in the publication of anyorganization, agency, institution, publication, service, or individual does not imply endorsement of the editors and publisher.Errors brought to the attention of the publisher and verified to the satisfaction of the publisher will be corrected in future editions.

The paper used in the publication meets the minimum requirements of American National Standard for InformationSciences—Permanence Paper for Printed Library Materials, ANSI Z39.48-1984.

This publication is a creative work fully protected by all applicable copyright laws, as well as by misappropriation, trade secret,unfair competition, and other applicable laws. The authors and editors of this work have added value to the underlying factualmaterial herein through one or more of the following: unique and original selection, coordination, expression, arrangement,and classification of the information.

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Copyright © 2002Gale Group, 27500 Drake Road, Farmington Hills, Michigan 48331-3535

No part of this book may be reproduced in any form without permission in writing from the publisher, except by a reviewerwho wishes to quote brief passages or entries in connection with a review written for inclusion in a magazine or newspaper.

ISBN 0-7876-5631-3 (set)0-7876-5632-1 (vol. 1) 0-7876-5634-8 (vol. 3)0-7876-5633-X (vol. 2) 0-7876-5635-6 (vol. 4)

Printed in the United States of America10 9 8 7 6 5 4 3 2 1

Library of Congress Cataloging-in-Publication Data

Knight, Judson.Science of everyday things / written by Judson Knight, Neil Schlager, editor.

p. cm.Includes bibliographical references and indexes.Contents: v. 1. Real-life chemistry – v. 2 Real-life physics.ISBN 0-7876-5631-3 (set : hardcover) – ISBN 0-7876-5632-1 (v. 1) – ISBN

0-7876-5633-X (v. 2)1. Science–Popular works. I. Schlager, Neil, 1966-II. Title.

Q162.K678 2001500–dc21 2001050121

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iiiSC IENCE OF EVERYDAY TH INGS VOLUME 1: REAL-LIFE CHEMISTRY

Introduction.............................................v

Advisory Board ......................................vii

MEASUREMENT

Measurement ......................................................3Temperature and Heat......................................11

Mass, Density, and Volume ..............................23

MATTER

Properties of Matter .........................................32

Gases..................................................................48

ATOMS AND MOLECULES

Atoms ................................................................63

Atomic Mass......................................................76

Electrons............................................................84

Isotopes .............................................................92

Ions and Ionization........................................101

Molecules........................................................109

ELEMENTS

Elements .........................................................119

Periodic Table of Elements ............................127

Families of Elements......................................140

METALS

Metals..............................................................149

Alkali Metals...................................................162

Alkaline Earth Metals ....................................171

Transition Metals ...........................................181

Actinides .........................................................196

Lanthanides ....................................................205

NONMETALS AND METALLOIDS

Nonmetals ......................................................213

Metalloids .......................................................222

Halogens .........................................................229

Noble Gases ....................................................237

Carbon............................................................243

Hydrogen........................................................252

BONDING AND REACTIONS

Chemical Bonding .........................................263

Compounds....................................................273

Chemical Reactions........................................281

Oxidation-Reduction Reactions....................289

Chemical Equilibrium ...................................297

Catalysts..........................................................304

Acids and Bases ..............................................310

Acid-Base Reactions.......................................319

SOLUTIONS AND MIXTURES

Mixtures..........................................................329

Solutions.........................................................338

Osmosis ..........................................................347

Distillation and Filtration..............................354

ORGANIC CHEMISTRY

Organic Chemistry ........................................363

Polymers .........................................................372

General Subject Index ......................381

CONTENTS

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vSC IENCE OF EVERYDAY TH INGS VOLUME 1: REAL-LIFE CHEMISTRY

I N TRODUCT ION

Overview of the Series

Welcome to Science of Everyday Things. Our aimis to explain how scientific phenomena can beunderstood by observing common, real-worldevents. From luminescence to echolocation tobuoyancy, the series will illustrate the chief prin-ciples that underlay these phenomena andexplore their application in everyday life. Toencourage cross-disciplinary study, the entrieswill draw on applications from a wide variety offields and endeavors.

Science of Everyday Things initially compris-es four volumes:

Volume 1: Real-Life ChemistryVolume 2: Real-Life PhysicsVolume 3: Real-Life BiologyVolume 4: Real-Life Earth Science

Future supplements to the series will expandcoverage of these four areas and explore newareas, such as mathematics.

Arrangement of Real-LifePhysics

This volume contains 40 entries, each covering adifferent scientific phenomenon or principle.The entries are grouped together under commoncategories, with the categories arranged, in gen-eral, from the most basic to the most complex.Readers searching for a specific topic should con-sult the table of contents or the general subjectindex.

Within each entry, readers will find the fol-lowing rubrics:

• Concept Defines the scientific principle ortheory around which the entry is focused.

• How It Works Explains the principle or the-ory in straightforward, step-by-step lan-guage.

• Real-Life Applications Describes how thephenomenon can be seen in everydayevents.

• Where to Learn More Includes books, arti-cles, and Internet sites that contain furtherinformation about the topic.

Each entry also includes a “Key Terms” sec-tion that defines important concepts discussed inthe text. Finally, each volume includes numerousillustrations, graphs, tables, and photographs.

In addition, readers will find the compre-hensive general subject index valuable in access-ing the data.

About the Editor, Author,and Advisory Board

Neil Schlager and Judson Knight would like tothank the members of the advisory board fortheir assistance with this volume. The advisorswere instrumental in defining the list of topics,and reviewed each entry in the volume for scien-tific accuracy and reading level. The advisorsinclude university-level academics as well as highschool teachers; their names and affiliations arelisted elsewhere in the volume.

NEIL SCHLAGER is the president ofSchlager Information Group Inc., an editorialservices company. Among his publications areWhen Technology Fails (Gale, 1994); HowProducts Are Made (Gale, 1994); the St. JamesPress Gay and Lesbian Almanac (St. James Press,1998); Best Literature By and About Blacks (Gale,

set_fm_v1 9/26/01 11:50 AM Page v

Introduction 2000); Contemporary Novelists, 7th ed. (St. JamesPress, 2000); and Science and Its Times (7 vols.,Gale, 2000-2001). His publications have wonnumerous awards, including three RUSA awardsfrom the American Library Association, twoReference Books Bulletin/Booklist Editors’Choice awards, two New York Public LibraryOutstanding Reference awards, and a CHOICEaward for best academic book.

Judson Knight is a freelance writer, andauthor of numerous books on subjects rangingfrom science to history to music. His work on science titles includes Science, Technology, andSociety, 2000 B.C.-A.D. 1799 (U*X*L, 2002),as well as extensive contributions to Gale’s seven-volume Science and Its Times (2000-2001).As a writer on history, Knight has publishedMiddle Ages Reference Library (2000), Ancient

Civilizations (1999), and a volume in U*X*L’s

African American Biography series (1998).

Knight’s publications in the realm of music

include Parents Aren’t Supposed to Like It (2001),

an overview of contemporary performers and

genres, as well as Abbey Road to Zapple Records: A

Beatles Encyclopedia (Taylor, 1999). His wife,

Deidre Knight, is a literary agent and president of

the Knight Agency. They live in Atlanta with their

daughter Tyler, born in November 1998.

Comments and Suggestions

Your comments on this series and suggestions for

future editions are welcome. Please write: The

Editor, Science of Everyday Things, Gale Group,

27500 Drake Road, Farmington Hills, MI 48331.

vi SC IENCE OF EVERYDAY TH INGSVOLUME 1: REAL-LIFE CHEMISTRY

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viiSC IENCE OF EVERYDAY TH INGS VOLUME 1: REAL-LIFE CHEMISTRY

T I T LEADV ISORY BOARD

William E. Acree, Jr.Professor of Chemistry, University of North Texas

Russell J. ClarkResearch Physicist, Carnegie Mellon University

Maura C. FlanneryProfessor of Biology, St. John’s University, New

York

John GoudieScience Instructor, Kalamazoo (MI) Area

Mathematics and Science Center

Cheryl HachScience Instructor, Kalamazoo (MI) Area


Michael SinclairPhysics instructor, Kalamazoo (MI) Area


Rashmi VenkateswaranSenior Instructor and Lab Coordinator,

University of OttawaOttawa, Ontario, Canada

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1

SC IENCE OF EVERYDAY TH INGSr e a l - l i f e c h em i s t r y

M EASUREMENTMEASUREMENT

MEASUREMENT

TEMPERATURE AND HEAT

MASS , DENS I T Y, AND VOLUME

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3SC IENCE OF EVERYDAY TH INGS VOLUME 1: REAL-LIFE CHEMISTRY

MEASUREMENTMeasurement

CONCEPTMeasurement seems like a simple subject, on thesurface at least; indeed, all measurements can bereduced to just two components: number andunit. Yet one might easily ask, “What numbers,and what units?”—a question that helps bringinto focus the complexities involved in designat-ing measurements. As it turns out, some forms ofnumbers are more useful for rendering valuesthan others; hence the importance of significantfigures and scientific notation in measurements.The same goes for units. First, one has to deter-mine what is being measured: mass, length, orsome other property (such as volume) that isultimately derived from mass and length. Indeed,the process of learning how to measure revealsnot only a fundamental component of chemistry,but an underlying—if arbitrary and manmade—order in the quantifiable world.

HOW IT WORKS

Numbers

In modern life, people take for granted the exis-tence of the base-10, of decimal numeration sys-tem—a name derived from the Latin worddecem, meaning “ten.” Yet there is nothing obvi-ous about this system, which has its roots in theten fingers used for basic counting. At othertimes in history, societies have adopted the twohands or arms of a person as their numericalframe of reference, and from this developed abase-2 system. There have also been base-5 sys-tems relating to the fingers on one hand, andbase-20 systems that took as their reference pointthe combined number of fingers and toes.

Obviously, there is an arbitrary qualityunderlying the modern numerical system, yet itworks extremely well. In particular, the use ofdecimal fractions (for example, 0.01 or 0.235) isparticularly helpful for rendering figures otherthan whole numbers. Yet decimal fractions are arelatively recent innovation in Western mathe-matics, dating only to the sixteenth century. Inorder to be workable, decimal fractions rely onan even more fundamental concept that was notalways part of Western mathematics: place-value.

Place-Value and NotationSystems

Place-value is the location of a number relative toothers in a sequence, a location that makes it pos-sible to determine the number’s value. Forinstance, in the number 347, the 3 is in the hun-dreds place, which immediately establishes avalue for the number in units of 100. Similarly, aperson can tell at a glance that there are 4 units of10, and 7 units of 1.

Of course, today this information appears tobe self-evident—so much so that an explanationof it seems tedious and perfunctory—to almostanyone who has completed elementary-schoolarithmetic. In fact, however, as with almosteverything about numbers and units, there isnothing obvious at all about place-value; other-wise, it would not have taken Western mathe-maticians thousands of years to adopt a place-value numerical system. And though they dideventually make use of such a system, Westernersdid not develop it themselves, as we shall see.

ROMAN NUMERALS. Numerationsystems of various kinds have existed since atleast 3000 B.C., but the most important number


Measurement instance, trying to multiply these two. With thenumber system in use today, it is not difficult tomultiply 3,000 by 438 in one’s head. The problemcan be reduced to a few simple steps: multiply 3by 400, 3 by 30, and 3 by 8; add these productstogether; then multiply the total by 1,000—a stepthat requires the placement of three zeroes at theend of the number obtained in the earlier steps.

But try doing this with Roman numerals: itis essentially impossible to perform this calcula-tion without resorting to the much more practi-cal place-value system to which we’re accus-tomed. No wonder, then, that Roman numeralshave been relegated to the sidelines, used in mod-ern life for very specific purposes: in outlines, forinstance; in ordinal titles (for example, HenryVIII); or in designating the year of a motion pic-ture’s release.

HINDU-ARABIC NUMERALS.

The system of counting used throughout muchof the world—1, 2, 3, and so on—is the Hindu-Arabic notation system. Sometimes mistakenlyreferred to as “Arabic numerals,” these are mostaccurately designated as Hindu or Indian numerals. They came from India, but becauseEuropeans discovered them in the Near East during the Crusades (1095-1291), they assumedthe Arabs had invented the notation system, andhence began referring to them as Arabic numerals.

Developed in India during the first millenni-um B.C., Hindu notation represented a vastimprovement over any method in use up to orindeed since that time. Of particular importancewas a number invented by Indian mathemati-cians: zero. Until then, no one had consideredzero worth representing since it was, after all,nothing. But clearly the zeroes in a number suchas 2,000,002 stand for something. They performa place-holding function: otherwise, it would beimpossible to differentiate between 2,000,002and 22.

Uses of Numbers in Science

SCIENTIFIC NOTATION. Chem-ists and other scientists often deal in very large orvery small numbers, and if they had to write outthese numbers every time they discussed them,their work would soon be encumbered bylengthy numerical expressions. For this purpose,they use scientific notation, a method for writingextremely large or small numbers by representing

4 SC IENCE OF EVERYDAY TH INGSVOLUME 1: REAL-LIFE CHEMISTRY

system in the history of Western civilization priorto the late Middle Ages was the one used by theRomans. Rome ruled much of the known worldin the period from about 200 B.C. to about A.D.200, and continued to have an influence onEurope long after the fall of the Western RomanEmpire in A.D. 476—an influence felt even today.Though the Roman Empire is long gone andLatin a dead language, the impact of Rome con-tinues: thus, for instance, Latin terms are used todesignate species in biology. It is therefore easy tounderstand how Europeans continued to use theRoman numeral system up until the thirteenthcentury A.D.—despite the fact that Romannumerals were enormously cumbersome.

The Roman notation system has no meansof representing place-value: thus a relatively largenumber such as 3,000 is shown as MMM, where-as a much smaller number might use many more“places”: 438, for instance, is rendered asCDXXXVIII. Performing any sort of calculationswith these numbers is a nightmare. Imagine, for

STANDARDIZATION IS CRUCIAL TO MAINTAINING STABILI-TY IN A SOCIETY. DURING THE GERMAN INFLATIONARY

CRISIS OF THE 1920S, HYPERINFLATION LED TO AN

ECONOMIC DEPRESSION AND THE RISE OF ADOLF

HITLER. HERE, TWO CHILDREN GAZE UP AT A STACK OF

100,000 GERMAN MARKS—THE EQUIVALENT AT THE

TIME TO ONE U.S. DOLLAR. (© Bettmann/Corbis.)


Measurement


them as a number between 1 and 10 multipliedby a power of 10.

Instead of writing 75,120,000, for instance,the preferred scientific notation is 7.512 • 107. Tointerpret the value of large multiples of 10, it ishelpful to remember that the value of 10 raised toany power n is the same as 1 followed by thatnumber of zeroes. Hence 1025, for instance, issimply 1 followed by 25 zeroes.

Scientific notation is just as useful—tochemists in particular—for rendering very smallnumbers. Suppose a sample of a chemical com-pound weighed 0.0007713 grams. The preferredscientific notation, then, is 7.713 • 10–4. Note thatfor numbers less than 1, the power of 10 is a neg-ative number: 10–1 is 0.1, 10–2 is 0.01, and so on.

Again, there is an easy rule of thumb forquickly assessing the number of decimal placeswhere scientific notation is used for numbers lessthan 1. Where 10 is raised to any power –n, thedecimal point is followed by n places. If 10 israised to the power of –8, for instance, we knowat a glance that the decimal is followed by 7zeroes and a 1.

SIGNIFICANT FIGURES. In mak-ing measurements, there will always be a degreeof uncertainty. Of course, when the standards of

calibration (discussed below) are very high, andthe measuring instrument has been properly cal-ibrated, the degree of uncertainty will be verysmall. Yet there is bound to be uncertainty tosome degree, and for this reason, scientists usesignificant figures—numbers included in ameasurement, using all certain numbers alongwith the first uncertain number.

Suppose the mass of a chemical sample ismeasured on a scale known to be accurate to 10-5 kg. This is equal to 1/100,000 of a kilo, or1/100 of a gram; or, to put it in terms of place-value, the scale is accurate to the fifth place in adecimal fraction. Suppose, then, that an item isplaced on the scale, and a reading of 2.13283697kg is obtained. All the numbers prior to the 6 aresignificant figures, because they have beenobtained with certainty. On the other hand, the 6and the numbers that follow are not significantfigures because the scale is not known to be accu-rate beyond 10-5 kg.

Thus the measure above should be renderedwith 7 significant figures: the whole number 2,and the first 6 decimal places. But if the value isgiven as 2.132836, this might lead to inaccuraciesat some point when the measurement is factoredinto other equations. The 6, in fact, should be“rounded off” to a 7. Simple rules apply to the

THE UNITED STATES NAVAL OBSERVATORY IN WASHINGTON, D.C., IS AMERICA’S PREEMINENT STANDARD FOR THE

EXACT TIME OF DAY. (Richard T. Nowitz/Corbis. Reproduced by permission.)


Measurement rounding off of significant figures: if the digit fol-lowing the first uncertain number is less than 5,there is no need to round off. Thus, if the meas-urement had been 2.13283627 kg (note that the 9was changed to a 2), there is no need to roundoff, and in this case, the figure of 2.132836 is cor-rect. But since the number following the 6 is infact a 9, the correct significant figure is 7; thus thetotal would be 2.132837.

Fundamental Standards of Measure

So much for numbers; now to the subject ofunits. But before addressing systems of measure-ment, what are the properties being measured?All forms of scientific measurement, in fact, canbe reduced to expressions of four fundamentalproperties: length, mass, time, and electric cur-rent. Everything can be expressed in terms ofthese properties: even the speed of an electronspinning around the nucleus of an atom can beshown as “length” (though in this case, the meas-urement of space is in the form of a circle or evenmore complex shapes) divided by time.

Of particular interest to the chemist arelength and mass: length is a component of vol-ume, and both length and mass are elements ofdensity. For this reason, a separate essay in thisbook is devoted to the subject of Mass, Density,and Volume. Note that “length,” as used in thismost basic sense, can refer to distance along anyplane, or in any of the three dimensions—com-monly known as length, width, and height—ofthe observable world. (Time is the fourth dimen-sion.) In addition, as noted above, “length” meas-urements can be circular, in which case the for-mula for measuring space requires use of thecoefficient π, roughly equal to 3.14.

REAL-L IFEAPPLICATIONS

Standardized Units of Measure: Who Needs Them?

People use units of measure so frequently in dailylife that they hardly think about what they aredoing. A motorist goes to the gas station andpumps 13 gallons (a measure of volume) into anautomobile. To pay for the gas, the motorist usesdollars—another unit of measure, economic

rather than scientific—in the form of papermoney, a debit card, or a credit card.

This is simple enough. But what if themotorist did not know how much gas was in agallon, or if the motorist had some idea of a gal-lon that differed from what the gas station man-agement determined it to be? And what if thevalue of a dollar were not established, such thatthe motorist and the gas station attendant had tohaggle over the cost of the gasoline just pur-chased? The result would be a horribly confusedsituation: the motorist might run out of gas, ormoney, or both, and if such confusion were mul-tiplied by millions of motorists and millions ofgas stations, society would be on the verge ofbreakdown.

THE VALUE OF STANDARD-

IZATION TO A SOCIETY. Actually,there have been times when the value of curren-cy was highly unstable, and the result was nearanarchy. In Germany during the early 1920s, forinstance, rampant inflation had so badly deplet-ed the value of the mark, Germany’s currency,that employees demanded to be paid every day sothat they could cash their paychecks before thevalue went down even further. People made jokesabout the situation: it was said, for instance, thatwhen a woman went into a store and left a basketcontaining several million marks out front,thieves ran by and stole the basket—but left themoney. Yet there was nothing funny about thissituation, and it paved the way for the nightmar-ish dictatorship of Adolf Hitler and the NaziParty.

It is understandable, then, that standardiza-tion of weights and measures has always been animportant function of government. When Ch’inShih-huang-ti (259-210 B.C.) united China forthe first time, becoming its first emperor, he setabout standardizing units of measure as a meansof providing greater unity to the country—thusmaking it easier to rule. On the other hand, theRussian Empire of the late nineteenth centuryfailed to adopt standardized systems that wouldhave tied it more closely to the industrializednations of Western Europe. The width of railroadtracks in Russia was different than in WesternEurope, and Russia used the old Julian calendar,as opposed to the Gregorian calendar adoptedthroughout much of Western Europe after 1582.These and other factors made economicexchanges between Russia and Western Europe



Measurementextremely difficult, and the Russian Empireremained cut off from the rapid progress of theWest. Like Germany a few decades later, itbecame ripe for the establishment of a dictator-ship—in this case under the Communists led byV. I. Lenin.

Aware of the important role that standardi-zation of weights and measures plays in the gov-erning of a society, the U.S. Congress in 1901established the Bureau of Standards. Today it isknown as the National Institute of Standards andTechnology (NIST), a nonregulatory agencywithin the Commerce Department. As will bediscussed at the conclusion of this essay, theNIST maintains a wide variety of standard defi-nitions regarding mass, length, temperature andso forth, against which other devices can be cali-brated.

THE VALUE OF STANDARD-

IZATION TO SCIENCE. What if anurse, rather than carefully measuring a quantityof medicine before administering it to a patient,simply gave the patient an amount that “lookedright”? Or what if a pilot, instead of calculatingfuel, distance, and other factors carefully beforetaking off from the runway, merely used a “bestestimate”? Obviously, in either case, disastrousresults would be likely to follow. Though neithernurses or pilots are considered scientists, bothuse science in their professions, and those disas-trous results serve to highlight the crucial matterof using standardized measurements in science.

Standardized measurements are necessary toa chemist or any scientist because, in order for anexperiment to be useful, it must be possible toduplicate the experiment. If the chemist does notknow exactly how much of a certain element heor she mixed with another to form a given com-pound, the results of the experiment are useless.In order to share information and communicatethe results of experiments, then, scientists need astandardized “vocabulary” of measures.

This “vocabulary” is the International Sys-tem of Units, known as SI for its French name,Système International d’Unités. By internationalagreement, the worldwide scientific communityadopted what came to be known as SI at the 9thGeneral Conference on Weights and Measures in1948. The system was refined at the 11th GeneralConference in 1960, and given its present name;but in fact most components of SI belong to amuch older system of weights and measures

developed in France during the late eighteenthcentury.

SI vs. the English System

The United States, as almost everyone knows, isthe wealthiest and most powerful nation onEarth. On the other hand, Brunei—a tiny nation-state on the island of Java in the Indonesianarchipelago—enjoys considerable oil wealth, butis hardly what anyone would describe as a super-power. Yemen, though it is located on the Arabi-an peninsula, does not even possess significantoil wealth, and is a poor, economically develop-ing nation. Finally, Burma in Southeast Asia canhardly be described even as a “developing”nation: ruled by an extremely repressive militaryregime, it is one of the poorest nations in theworld.

So what do these four have in common?They are the only nations on the planet that havefailed to adopt the metric system of weights andmeasures. The system used in the United States iscalled the English system, though it should moreproperly be called the American system, sinceEngland itself has joined the rest of the world in“going metric.” Meanwhile, Americans continueto think in terms of gallons, miles, and pounds;yet American scientists use the much more con-venient metric units that are part of SI.

HOW THE ENGLISH SYSTEM

WORKS (OR DOES NOT WORK).

Like methods of counting described above, mostsystems of measurement in premodern timeswere modeled on parts of the human body. Thefoot is an obvious example of this, while the inchoriginated from the measure of a king’s firstthumb joint. At one point, the yard was definedas the distance from the nose of England’s KingHenry I to the tip of his outstretched middle finger.

Obviously, these are capricious, downrightabsurd standards on which to base a system ofmeasure. They involve things that change,depending for instance on whose foot is beingused as a standard. Yet the English system devel-oped in this willy-nilly fashion over the centuries;today, there are literally hundreds of units—including three types of miles, four kinds ofounces, and five kinds of tons, each with a differ-ent value.

What makes the English system particularlycumbersome, however, is its lack of convenient



Measurement conversion factors. For length, there are 12 inch-es in a foot, but 3 feet in a yard, and 1,760 yardsin a mile. Where volume is concerned, there are16 ounces in a pound (assuming one is talkingabout an avoirdupois ounce), but 2,000 poundsin a ton. And, to further complicate matters,there are all sorts of other units of measure devel-oped to address a particular property: horsepow-er, for instance, or the British thermal unit (Btu).

THE CONVENIENCE OF THE

METRIC SYSTEM. Great Britain, thoughit has long since adopted the metric system, in1824 established the British Imperial System,aspects of which are reflected in the system stillused in America. This is ironic, given the desire ofearly Americans to distance themselves psycho-logically from the empire to which their nationhad once belonged. In any case, England’s greatworldwide influence during the nineteenth cen-tury brought about widespread adoption of theEnglish or British system in colonies such as Aus-tralia and Canada. This acceptance had every-thing to do with British power and tradition, andnothing to do with convenience. A much moreusable standard had actually been embraced 25years before in a land that was then among Eng-land’s greatest enemies: France.

During the period leading up to and follow-ing the French Revolution of 1789, French intel-lectuals believed that every aspect of existencecould and should be treated in highly rational,scientific terms. Out of these ideas arose muchfolly, particularly during the Reign of Terror in1793, but one of the more positive outcomes wasthe metric system. This system is decimal—thatis, based entirely on the number 10 and powersof 10, making it easy to relate one figure toanother. For instance, there are 100 centimetersin a meter and 1,000 meters in a kilometer.

PREFIXES FOR SIZES IN THE

METRIC SYSTEM. For designating small-er values of a given measure, the metric systemuses principles much simpler than those of theEnglish system, with its irregular divisions of (forinstance) gallons, quarts, pints, and cups. In themetric system, one need only use a simple Greekor Latin prefix to designate that the value is mul-tiplied by a given power of 10. In general, the pre-fixes for values greater than 1 are Greek, whileLatin is used for those less than 1. These prefixes,along with their abbreviations and respective val-

ues, are as follows. (The symbol µ for "micro" isthe Greek letter mu.)

The Most Commonly Used Prefixes in theMetric System

• giga (G) = 109 (1,000,000,000)

• mega (M) = 106 (1,000,000)

• kilo (k) == 103 (1,000)

• deci (d) = 10–1 (0.1)

• centi (c) = 10–2 (0.01)

• milli (m) = 10–3 (0.001)

• micro (µ) = 10–6 (0.000001)

• nano (n) = 10–9 (0.000000001)

The use of these prefixes can be illustratedby reference to the basic metric unit of length,the meter. For long distances, a kilometer (1,000 m) is used; on the other hand, very shortdistances may require a centimeter (0.01 m) or amillimeter (0.001 m) and so on, down to ananometer (0.000000001 m). Measurements oflength also provide a good example of why SIincludes units that are not part of the metric sys-tem, though they are convertible to metric units.Hard as it may be to believe, scientists oftenmeasure lengths even smaller than a nanome-ter—the width of an atom, for instance, or thewavelength of a light ray. For this purpose,they use the angstrom (Å or A), equal to 0.1nanometers.

Calibration and SI Units

THE SEVEN BASIC SI UNITS.

The SI uses seven basic units, representinglength, mass, time, temperature, amount of sub-stance, electric current, and luminous intensity.The first four parameters are a part of everydaylife, whereas the last three are of importance onlyto scientists. “Amount of substance” is the num-ber of elementary particles in matter. This ismeasured by the mole, a unit discussed in theessay on Mass, Density, and Volume. Luminousintensity, or the brightness of a light source, ismeasured in candelas, while the SI unit of electriccurrent is the ampere.

The other four basic units are the meter forlength, the kilogram for mass, the second fortime, and the degree Celsius for temperature. Thelast of these is discussed in the essay on Temper-ature; as for meters, kilograms, and seconds, theywill be examined below in terms of the meansused to define each.



MeasurementCALIBRATION. Calibration is theprocess of checking and correcting the perform-ance of a measuring instrument or device againstthe accepted standard. America’s preeminentstandard for the exact time of day, for instance, isthe United States Naval Observatory in Washing-ton, D.C. Thanks to the Internet, people all overthe country can easily check the exact time, andcalibrate their clocks accordingly—though, ofcourse, the resulting accuracy is subject to factorssuch as the speed of the Internet connection.

There are independent scientific laboratoriesresponsible for the calibration of certain instru-ments ranging from clocks to torque wrenches,and from thermometers to laser-beam poweranalyzers. In the United States, instruments ordevices with high-precision applications—thatis, those used in scientific studies, or by high-techindustries—are calibrated according to standardsestablished by the NIST.

The NIST keeps on hand definitions, asopposed to using a meter stick or other physicalmodel. This is in accordance with the methods ofcalibration accepted today by scientists: ratherthan use a standard that might vary—forinstance, the meter stick could be bent impercep-tibly—unvarying standards, based on specificbehaviors in nature, are used.

METERS AND KILOGRAMS. Ameter, equal to 3.281 feet, was at one timedefined in terms of Earth’s size. Using an imagi-nary line drawn from the Equator to the NorthPole through Paris, this distance was divided into10 million meters. Later, however, scientists cameto the realization that Earth is subject to geologi-cal changes, and hence any measurement cali-brated to the planet’s size could not ultimately bereliable. Today the length of a meter is calibratedaccording to the amount of time it takes light totravel through that distance in a vacuum (an areaof space devoid of air or other matter). The offi-cial definition of a meter, then, is the distancetraveled by light in the interval of 1/299,792,458of a second.

One kilogram is, on Earth at least, equal to2.21 pounds; but whereas the kilogram is a unitof mass, the pound is a unit of weight, so the cor-respondence between the units varies dependingon the gravitational field in which a pound ismeasured. Yet the kilogram, though it representsa much more fundamental property of the phys-ical world than a pound, is still a somewhat arbi-


trary form of measure in comparison to themeter as it is defined today.

Given the desire for an unvarying standardagainst which to calibrate measurements, itwould be helpful to find some usable butunchanging standard of mass; unfortunately, sci-entists have yet to locate such a standard. There-fore, the value of a kilogram is calibrated much asit was two centuries ago. The standard is a bar ofplatinum-iridium alloy, known as the Interna-tional Prototype Kilogram, housed near Sèvres inFrance.

SECONDS. A second, of course, is aunit of time as familiar to non-scientificallytrained Americans as it is to scientists and peopleschooled in the metric system. In fact, it hasnothing to do with either the metric system or SI.The means of measuring time on Earth are not“metric”: Earth revolves around the Sun approx-imately every 365.25 days, and there is no way toturn this into a multiple of 10 without creating a

CALIBRATION: The process of check-

ing and correcting the performance of a

measuring instrument or device against a

commonly accepted standard.

SCIENTIFIC NOTATION: A method

used by scientists for writing extremely

large or small numbers by representing

them as a number between 1 and 10 multi-

plied by a power of 10. Instead of writing

0.0007713, the preferred scientific notation

is 7.713 • 10–4.

SI: An abbreviation of the French term

Système International d’Unités, or Interna-

tional System of Units. Based on the metric

system, SI is the system of measurement

units in use by scientists worldwide.

SIGNIFICANT FIGURES: Numbers

included in a measurement, using all cer-

tain numbers along with the first uncertain

number.

KEY TERMS


Measurement situation even more cumbersome than the Eng-lish units of measure.

The week and the month are units based oncycles of the Moon, though they are no longerrelated to lunar cycles because a lunar year wouldsoon become out-of-phase with a year based onEarth’s rotation around the Sun. The continuinguse of weeks and months as units of time is basedon tradition—as well as the essential need of asociety to divide up a year in some way.

A day, of course, is based on Earth’s rotation,but the units into which the day is divided—hours, minutes, and seconds—are purely arbi-trary, and likewise based on traditions of longstanding. Yet scientists must have some unit oftime to use as a standard, and, for this purpose,the second was chosen as the most practical. TheSI definition of a second, however, is not simplyone-sixtieth of a minute or anything else sostrongly influenced by the variation of Earth’smovement.

Instead, the scientific community chose asits standard the atomic vibration of a particularisotope of the metal cesium, cesium-133. Thevibration of this atom is presumed to be un-varying, because the properties of elements—unlike the size of Earth or its movement—do not change. Today, a second is defined as theamount of time it takes for a cesium-133 atom

to vibrate 9,192,631,770 times. Expressed in

scientific notation, with significant figures, this is

9.19263177 • 109.

WHERE TO LEARN MORE

Gardner, Robert. Science Projects About Methods of Mea-suring. Berkeley Heights, N.J.: Enslow Publishers,2000.

Long, Lynette. Measurement Mania: Games and ActivitiesThat Make Math Easy and Fun. New York: Wiley,2001.

“Measurement” (Web site).<http://www.dist214.k12.il.us/users/asanders/meas.html> (May 7, 2001).

“Measurement in Chemistry” (Web site). <http://bradley.edu/~campbell/lectnotes/149ch2/tsld001.htm> (May7, 2001).

MegaConverter 2 (Web site). <http://www.megaconverter.com> (May 7, 2001).

Patilla, Peter. Measuring. Des Plaines, IL: HeinemannLibrary, 2000.

Richards, Jon. Units and Measurements. Brookfield, CT:Copper Beech Books, 2000.

Sammis, Fran. Measurements. New York: BenchmarkBooks, 1998.

Units of Measurement (Web site). <http://www.unc.edu/~rowlett/units/> (May 7, 2001).

Wilton High School Chemistry Coach (Web site).<http://www.chemistrycoach.com> (May 7, 2001).




T EMPER ATURE AND HEATTemperature and Heat

CONCEPTTemperature, heat, and related concepts belongto the world of physics rather than chemistry; yetit would be impossible for the chemist to workwithout an understanding of these properties.Thermometers, of course, measure temperatureaccording to one or both of two well-knownscales based on the freezing and boiling points ofwater, though scientists prefer a scale based onthe virtual freezing point of all matter. Also relat-ed to temperature are specific heat capacity, orthe amount of energy required to change thetemperature of a substance, and also calorimetry,the measurement of changes in heat as a result ofphysical or chemical changes. Although theseconcepts do not originate from chemistry butfrom physics, they are no less useful to thechemist.

HOW IT WORKS

Energy

The area of physics known as thermodynamics,discussed briefly below in terms of thermody-namics laws, is the study of the relationshipsbetween heat, work, and energy. Work is definedas the exertion of force over a given distance todisplace or move an object, and energy is theability to accomplish work. Energy appears innumerous manifestations, including thermalenergy, or the energy associated with heat.

Another type of energy—one of particularinterest to chemists—is chemical energy, relatedto the forces that attract atoms to one another inchemical bonds. Hydrogen and oxygen atoms inwater, for instance, are joined by chemical bond-

ing, and when those bonds are broken, the forcesjoining the atoms are released in the form ofchemical energy. Another example of chemicalenergy release is combustion, whereby chemicalbonds in fuel, as well as in oxygen molecules, arebroken and new chemical bonds are formed. Thetotal energy in the newly formed chemical bondsis less than the energy of the original bonds, butthe energy that makes up the difference is notlost; it has simply been released.

Energy, in fact, is never lost: a fundamentallaw of the universe is the conservation of energy,which states that in a system isolated from allother outside factors, the total amount of energyremains the same, though transformations ofenergy from one form to another take place.When a fire burns, then, some chemical energy is turned into thermal energy. Similar trans-formations occur between these and other mani-festations of energy, including electrical andmagnetic (sometimes these two are combined aselectromagnetic energy), sound, and nuclearenergy. If a chemical reaction makes a noise, forinstance, some of the energy in the substancesbeing mixed has been dissipated to make thatsound. The overall energy that existed before thereaction will be the same as before; however, theenergy will not necessarily be in the same place asbefore.

Note that chemical and other forms of ener-gy are described as “manifestations,” rather than“types,” of energy. In fact, all of these can bedescribed in terms of two basic types of energy:kinetic energy, or the energy associated withmovement, and potential energy, or the energyassociated with position. The two are inverselyrelated: thus, if a spring is pulled back to its max-imum point of tension, its potential energy is


Temperatureand Heat

coldness is a recognizable sensory experience inhuman life, in scientific terms, cold is simply theabsence of heat.

If you grasp a snowball in your hand, thehand of course gets cold. The mind perceives thisas a transfer of cold from the snowball, but in factexactly the opposite has happened: heat hasmoved from your hand to the snow, and ifenough heat enters the snowball, it will melt. Atthe same time, the departure of heat from yourhand results in a loss of internal energy near thesurface of the hand, experienced as a sensation ofcoldness.

Understanding Temperature

Just as heat does not mean the same thing in sci-entific terms as it does in ordinary language, so“temperature” requires a definition that sets itapart from its everyday meaning. Temperaturemay be defined as a measure of the average inter-nal energy in a system. Two systems in a state of thermal equilibrium have the same tempera-ture; on the other hand, differences in tempera-ture determine the direction of internal energyflow between two systems where heat is beingtransferred.

This can be illustrated through an experi-ence familiar to everyone: having one’s tempera-ture taken with a thermometer. If one has a fever,the mouth will be warmer than the thermometer,and therefore heat will be transferred to the ther-mometer from the mouth. The thermometer,discussed in more depth later in this essay, meas-ures the temperature difference between itselfand any object with which it is in contact.

Temperature and Thermodynamics

One might pour a kettle of boiling water into acold bathtub to heat it up; or one might put anice cube in a hot cup of coffee “to cool it down.”In everyday experience, these seem like two verydifferent events, but from the standpoint of ther-modynamics, they are exactly the same. In bothcases, a body of high temperature is placed incontact with a body of low temperature, and inboth cases, heat passes from the high-tempera-ture body to the low-temperature body.

The boiling water warms the tub of coolwater, and due to the high ratio of cool water toboiling water in the bathtub, the boiling water


also at a maximum, while its kinetic energy iszero. Once it is released and begins springingthrough the air to return to the position it main-tained before it was stretched, it begins gainingkinetic energy and losing potential energy.

Heat

Thermal energy is actually a form of kineticenergy generated by the movement of particles atthe atomic or molecular level: the greater themovement of these particles, the greater the ther-mal energy. When people use the word “heat” inordinary language, what they are really referringto is “the quality of hotness”—that is, the ther-mal energy internal to a system. In scientificterms, however, heat is internal thermal energythat flows from one body of matter to another—or, more specifically, from a system at a highertemperature to one at a lower temperature.

Two systems at the same temperature aresaid to be in a state of thermal equilibrium.When this state exists, there is no exchange ofheat. Though in everyday terms people speak of“heat” as an expression of relative warmth orcoldness, in scientific terms, heat exists only intransfer between two systems. Furthermore,there can never be a transfer of “cold”; although

DURING HIS LIFETIME, GALILEO CONSTRUCTED A THER-MOSCOPE, THE FIRST PRACTICAL TEMPERATURE-MEAS-URING DEVICE. (Archive Photos, Inc. Reproduced by permission.)


Temperatureand Heat


expends all its energy raising the temperature inthe bathtub as a whole. The greater the ratio ofvery hot water to cool water, of course, thewarmer the bathtub will be in the end. But evenafter the bath water is heated, it will continue tolose heat, assuming the air in the room is notwarmer than the water in the tub—a safeassumption. If the water in the tub is warmerthan the air, it will immediately begin transfer-ring thermal energy to the lower-temperature airuntil their temperatures are equalized.

As for the coffee and the ice cube, what hap-pens is opposite to the explanation ordinarilygiven. The ice does not “cool down” the coffee:the coffee warms up, and presumably melts, theice. However, it expends at least some of its ther-mal energy in doing so, and, as a result, the coffeebecomes cooler than it was.

THE LAWS OF THERMODY-

NAMICS. These situations illustrate the sec-ond of the three laws of thermodynamics. Notonly do these laws help to clarify the relationshipbetween heat, temperature, and energy, but theyalso set limits on what can be accomplished inthe world. Hence British writer and scientist C. P.Snow (1905-1980) once described the thermody-

namics laws as a set of rules governing an impos-

sible game.

The first law of thermodynamics is essential-

ly the same as the conservation of energy:

because the amount of energy in a system

remains constant, it is impossible to perform

work that results in an energy output greater

than the energy input. It could be said that the

conservation of energy shows that “the glass is

half full”: energy is never lost. By contrast, the

first law of thermodynamics shows that “the glass

is half empty”: no system can ever produce more

energy than was put into it. Snow therefore

summed up the first law as stating that the game

is impossible to win.

The second law of thermodynamics begins

from the fact that the natural flow of heat is

always from an area of higher temperature to an

area of lower temperature—just as was shown in

the bathtub and coffee cup examples above. Con-

sequently, it is impossible for any system to take

heat from a source and perform an equivalent

amount of work: some of the heat will always be

lost. In other words, no system can ever be per-

fectly efficient: there will always be a degree of

BECAUSE OF WATER’S HIGH SPECIFIC HEAT CAPACITY, CITIES LOCATED NEXT TO LARGE BODIES OF WATER TEND TO

STAY WARMER IN THE WINTER AND COOLER IN THE SUMMER. DURING THE EARLY SUMMER MONTHS, FOR INSTANCE,CHICAGO’S LAKEFRONT STAYS COOLER THAN AREAS FURTHER INLAND. THIS IS BECAUSE THE LAKE IS COOLED FROM

THE WINTER’S COLD TEMPERATURES AND SNOW RUNOFF. (Farrell Grehan/Corbis. Reproduced by permission.)


Temperatureand Heat

breakdown, evidence of a natural tendency calledentropy.

Snow summed up the second law of ther-modynamics, sometimes called “the law ofentropy,” thus: not only is it impossible to win, itis impossible to break even. In effect, the secondlaw compounds the “bad news” delivered by thefirst with some even worse news. Though it istrue that energy is never lost, the energy availablefor work output will never be as great as the ener-gy put into a system.

The third law of thermodynamics states thatat the temperature of absolute zero—a phenom-enon discussed later in this essay—entropy alsoapproaches zero. This might seem to counteractthe second law, but in fact the third states ineffect that absolute zero is impossible to reach.The French physicist and engineer Sadi Carnot(1796-1832) had shown that a perfectly efficientengine is one whose lowest temperature wasabsolute zero; but the second law of thermody-namics shows that a perfectly efficient engine (orany other perfect system) cannot exist. Hence, asSnow observed, not only is it impossible to winor break even; it is impossible to get out of thegame.


Evolution of the Thermometer

A thermometer is a device that gauges tempera-ture by measuring a temperature-dependentproperty, such as the expansion of a liquid in asealed tube. The Greco-Roman physician Galen(c. 129-c. 199) was among the first thinkers toenvision a scale for measuring temperature, butdevelopment of a practical temperature-measur-ing device—the thermoscope—did not occuruntil the sixteenth century.

The great physicist Galileo Galilei (1564-1642) may have invented the thermoscope; cer-tainly he constructed one. Galileo’s thermoscopeconsisted of a long glass tube planted in a con-tainer of liquid. Prior to inserting the tube intothe liquid—which was usually colored water,though Galileo’s thermoscope used wine—asmuch air as possible was removed from the tube.This created a vacuum (an area devoid of matter,including air), and as a result of pressure differ-

ences between the liquid and the interior of thethermoscope tube, some of the liquid went intothe tube.

But the liquid was not the thermometricmedium—that is, the substance whose tempera-ture-dependent property changes were measuredby the thermoscope. (Mercury, for instance, is thethermometric medium in many thermometerstoday; however, due to the toxic quality of mer-cury, an effort is underway to remove mercurythermometers from U.S. schools.) Instead, the airwas the medium whose changes the thermoscopemeasured: when it was warm, the air expanded,pushing down on the liquid; and when the aircooled, it contracted, allowing the liquid to rise.

EARLY THERMOMETERS: THE

SEARCH FOR A TEMPERATURE

SCALE. The first true thermometer, built byFerdinand II, Grand Duke of Tuscany (1610-1670) in 1641, used alcohol sealed in glass. Thelatter was marked with a temperature scale con-taining 50 units, but did not designate a value forzero. In 1664, English physicist Robert Hooke(1635-1703) created a thermometer with a scaledivided into units equal to about 1/500 of thevolume of the thermometric medium. For thezero point, Hooke chose the temperature atwhich water freezes, thus establishing a standardstill used today in the Fahrenheit and Celsiusscales.

Olaus Roemer (1644-1710), a Danishastronomer, introduced another important stan-dard. Roemer’s thermometer, built in 1702, wasbased not on one but two fixed points, which hedesignated as the temperature of snow orcrushed ice on the one hand, and the boilingpoint of water on the other. As with Hooke’s useof the freezing point, Roemer’s idea of designat-ing the freezing and boiling points of water as thetwo parameters for temperature measurementshas remained in use ever since.

Temperature Scales

THE FAHRENHEIT SCALE. Notonly did he develop the Fahrenheit scale, oldestof the temperature scales still used in Westernnations today, but in 1714, German physicistDaniel Fahrenheit (1686-1736) built the firstthermometer to contain mercury as a thermo-metric medium. Alcohol has a low boiling point,whereas mercury remains fluid at a wide range oftemperatures. In addition, it expands and con-



Temperatureand Heat

tracts at a very constant rate, and tends not tostick to glass. Furthermore, its silvery colormakes a mercury thermometer easy to read.

Fahrenheit also conceived the idea of using“degrees” to measure temperature. It is no mis-take that the same word refers to portions of acircle, or that exactly 180 degrees—half the num-ber of degrees in a circle—separate the freezingand boiling points for water on Fahrenheit’s ther-mometer. Ancient astronomers first divided acircle into 360 degrees, as a close approximationof the ratio between days and years, because 360has a large quantity of divisors. So, too, does180—a total of 16 whole-number divisors otherthan 1 and itself.

Though today it might seem obvious that 0should denote the freezing point of water, and180 its boiling point, such an idea was far fromobvious in the early eighteenth century. Fahren-heit considered a 0-to-180 scale, but also a 180-to-360 one, yet in the end he chose neither—orrather, he chose not to equate the freezing pointof water with zero on his scale. For zero, he chosethe coldest possible temperature he could createin his laboratory, using what he described as “amixture of sal ammoniac or sea salt, ice, andwater.” Salt lowers the melting point of ice (whichis why it is used in the northern United States tomelt snow and ice from the streets on cold win-ter days), and thus the mixture of salt and iceproduced an extremely cold liquid water whosetemperature he equated to zero.

On the Fahrenheit scale, the ordinary freez-ing point of water is 32°, and the boiling pointexactly 180° above it, at 212°. Just a few years afterFahrenheit introduced his scale, in 1730, a Frenchnaturalist and physicist named Rene AntoineFerchault de Reaumur (1683-1757) presented ascale for which 0° represented the freezing pointof water and 80° the boiling point. Although theReaumur scale never caught on to the sameextent as Fahrenheit’s, it did include one valuableaddition: the specification that temperature val-ues be determined at standard sea-level atmos-pheric pressure.

THE CELSIUS SCALE. With its32° freezing point and its 212° boiling point, theFahrenheit system lacks the neat orderliness of adecimal or base-10 scale. Thus when Franceadopted the metric system in 1799, it chose as itstemperature scale not the Fahrenheit but the

Celsius scale. The latter was created in 1742 bySwedish astronomer Anders Celsius (1701-1744).

Like Fahrenheit, Celsius chose the freezingand boiling points of water as his two referencepoints, but he determined to set them 100, ratherthan 180, degrees apart. The Celsius scale issometimes called the centigrade scale, because itis divided into 100 degrees, cent being a Latinroot meaning “hundred.” Interestingly, Celsiusplanned to equate 0° with the boiling point, and100° with the freezing point; only in 1750 did fel-low Swedish physicist Martin Strömer change theorientation of the Celsius scale. In accordancewith the innovation offered by Reaumur, Cel-sius’s scale was based not simply on the boilingand freezing points of water, but specifically onthose points at normal sea-level atmosphericpressure.

In SI, a scientific system of measurementthat incorporates units from the metric systemalong with additional standards used only by sci-entists, the Celsius scale has been redefined interms of the triple point of water. (Triple point isthe temperature and pressure at which a sub-stance is at once a solid, liquid, and vapor.)According to the SI definition, the triple point ofwater—which occurs at a pressure considerablybelow normal atmospheric pressure—is exactly0.01°C.

THE KELVIN SCALE. French physi-cist and chemist J. A. C. Charles (1746-1823),who is credited with the gas law that bears hisname (see below), discovered that at 0°C, the vol-ume of gas at constant pressure drops by 1/273for every Celsius degree drop in temperature.This suggested that the gas would simply disap-pear if cooled to -273°C, which of course madeno sense.

The man who solved the quandary raised byCharles’s discovery was William Thompson, LordKelvin (1824-1907), who, in 1848, put forwardthe suggestion that it was the motion of mole-cules, and not volume, that would become zero at–273°C. He went on to establish what came to beknown as the Kelvin scale. Sometimes known asthe absolute temperature scale, the Kelvin scale isbased not on the freezing point of water, but onabsolute zero—the temperature at which molec-ular motion comes to a virtual stop. This is–273.15°C (–459.67°F), which, in the Kelvinscale, is designated as 0K. (Kelvin measures donot use the term or symbol for “degree.”)



Temperatureand Heat

Though scientists normally use metric units,they prefer the Kelvin scale to Celsius because theabsolute temperature scale is directly related toaverage molecular translational energy, based onthe relative motion of molecules. Thus if theKelvin temperature of an object is doubled, thismeans its average molecular translational energyhas doubled as well. The same cannot be said ifthe temperature were doubled from, say, 10°C to20°C, or from 40°C to 80°F, since neither the Celsius nor the Fahrenheit scale is based onabsolute zero.

CONVERSIONS BETWEEN SCALES.

The Kelvin scale is closely related to the Celsiusscale, in that a difference of one degree measuresthe same amount of temperature in both. There-fore, Celsius temperatures can be converted toKelvins by adding 273.15. Conversion betweenCelsius and Fahrenheit figures, on the otherhand, is a bit trickier.

To convert a temperature from Celsius toFahrenheit, multiply by 9/5 and add 32. It isimportant to perform the steps in that order,because reversing them will produce a wrong fig-ure. Thus, 100°C multiplied by 9/5 or 1.8 equals180, which, when added to 32 equals 212°F.Obviously, this is correct, since 100°C and 212°Feach represent the boiling point of water. But ifone adds 32 to 100°, then multiplies it by 9/5, theresult is 237.6°F—an incorrect answer.

For converting Fahrenheit temperatures toCelsius, there are also two steps involving multi-plication and subtraction, but the order isreversed. Here, the subtraction step is performedbefore the multiplication step: thus 32 is sub-tracted from the Fahrenheit temperature, thenthe result is multiplied by 5/9. Beginning with212°F, when 32 is subtracted, this equals 180.Multiplied by 5/9, the result is 100°C—the cor-rect answer.

One reason the conversion formulae usesimple fractions instead of decimal fractions(what most people simply call “decimals”) is that5/9 is a repeating decimal fraction (0.55555....)Furthermore, the symmetry of 5/9 and 9/5 makesmemorization easy. One way to remember theformula is that Fahrenheit is multiplied by a frac-tion—since 5/9 is a real fraction, whereas 9/5 isactually a mixed number, or a whole numberplus a fraction.

Modern Thermometers

MERCURY THERMOMETERS.

For a thermometer, it is important that the glasstube be kept sealed; changes in atmospheric pres-sure contribute to inaccurate readings, becausethey influence the movement of the thermomet-ric medium. It is also important to have a reliablethermometric medium, and, for this reason,water—so useful in many other contexts—wasquickly discarded as an option.

Water has a number of unusual properties: itdoes not expand uniformly with a rise in tem-perature, or contract uniformly with a loweredtemperature. Rather, it reaches its maximumdensity at 39.2°F (4°C), and is less dense bothabove and below that temperature. Thereforealcohol, which responds in a much more uni-form fashion to changes in temperature, soontook the place of water, and is still used in manythermometers today. But for the reasons men-tioned earlier, mercury is generally consideredpreferable to alcohol as a thermometric medium.

In a typical mercury thermometer, mercuryis placed in a long, narrow sealed tube called acapillary. The capillary is inscribed with figuresfor a calibrated scale, usually in such a way as toallow easy conversions between Fahrenheit andCelsius. A thermometer is calibrated by measur-ing the difference in height between mercury atthe freezing point of water, and mercury at theboiling point of water. The interval betweenthese two points is then divided into equal incre-ments—180, as we have seen, for the Fahrenheitscale, and 100 for the Celsius scale.

VOLUME GAS THERMOME-

TERS. Whereas most liquids and solidsexpand at an irregular rate, gases tend to follow afairly regular pattern of expansion in response toincreases in temperature. The predictable behav-ior of gases in these situations has led to thedevelopment of the volume gas thermometer, ahighly reliable instrument against which otherthermometers—including those containing mer-cury—are often calibrated.

In a volume gas thermometer, an emptycontainer is attached to a glass tube containingmercury. As gas is released into the empty con-tainer; this causes the column of mercury tomove upward. The difference between the earlierposition of the mercury and its position after theintroduction of the gas shows the differencebetween normal atmospheric pressure and the



Temperatureand Heat

pressure of the gas in the container. It is then pos-sible to use the changes in the volume of the gasas a measure of temperature.

ELECTRIC THERMOMETERS.

All matter displays a certain resistance to electriccurrent, a resistance that changes with tempera-ture; because of this, it is possible to obtain tem-perature measurements using an electric ther-mometer. A resistance thermometer is equippedwith a fine wire wrapped around an insulator:when a change in temperature occurs, the resist-ance in the wire changes as well. This allowsmuch quicker temperature readings than thoseoffered by a thermometer containing a tradition-al thermometric medium.

Resistance thermometers are highly reliable,but expensive, and primarily are used for veryprecise measurements. More practical for every-day use is a thermistor, which also uses the prin-ciple of electric resistance, but is much simplerand less expensive. Thermistors are used for pro-viding measurements of the internal temperatureof food, for instance, and for measuring humanbody temperature.

Another electric temperature-measurementdevice is a thermocouple. When wires of two dif-ferent materials are connected, this creates asmall level of voltage that varies as a function oftemperature. A typical thermocouple uses twojunctions: a reference junction, kept at some con-stant temperature, and a measurement junction.The measurement junction is applied to the itemwhose temperature is to be measured, and anytemperature difference between it and the refer-ence junction registers as a voltage change, meas-ured with a meter connected to the system.

OTHER TYPES OF THERMOME-

TER. A pyrometer also uses electromagneticproperties, but of a very different kind. Ratherthan responding to changes in current or voltage,the pyrometer is gauged to respond to visible andinfrared radiation. As with the thermocouple, apyrometer has both a reference element and ameasurement element, which compares lightreadings between the reference filament and theobject whose temperature is being measured.

Still other thermometers, such as those in anoven that register the oven’s internal tempera-ture, are based on the expansion of metals withheat. In fact, there are a wide variety of ther-mometers, each suited to a specific purpose. Apyrometer, for instance, is good for measuring

the temperature of a object with which the ther-mometer itself is not in physical contact.

Measuring Heat

The measurement of temperature by degrees inthe Fahrenheit or Celsius scales is a part of dailylife, but measurements of heat are not as familiarto the average person. Because heat is a form ofenergy, and energy is the ability to perform work,heat is therefore measured by the same units aswork. The principal SI unit of work or energy isthe joule (J). A joule is equal to 1 newton-meter(N • m)—in other words, the amount of energyrequired to accelerate a mass of 1 kilogram at therate of 1 meter per second squared across a dis-tance of 1 meter.

The joule’s equivalent in the English systemis the foot-pound: 1 foot-pound is equal to 1.356J, and 1 joule is equal to 0.7376 ft • lbs. In theBritish system, Btu, or British thermal unit, isanother measure of energy, though it is primari-ly used for machines. Due to the cumbersomenature of the English system, contrasted with theconvenience of the decimal units in the SI system, these English units of measure are notused by chemists or other scientists for heatmeasurement.

Specific Heat Capacity

Specific heat capacity (sometimes called specificheat) is the amount of heat that must be addedto, or removed from, a unit of mass for a givensubstance to change its temperature by 1°C. Typ-ically, specific heat capacity is measured in unitsof J/g • °C (joules per gram-degree Celsius).

The specific heat capacity of water is meas-ured by the calorie, which, along with the joule, isan important SI measure of heat. Often anotherunit, the kilocalorie—which, as its name sug-gests—is 1,000 calories—is used. This is one ofthe few confusing aspects of SI, which is muchsimpler than the English system. The dietaryCalorie (capital C) with which most people arefamiliar is not the same as a calorie (lowercasec)—rather, a dietary Calorie is the same as a kilo-calorie.

COMPARING SPECIFIC HEAT

CAPACITIES. The higher the specific heatcapacity, the more resistant the substance is tochanges in temperature. Many metals, in fact,have a low specific heat capacity, making themeasy to heat up and cool down. This contributes



Temperatureand Heat

to the tendency of metals to expand when heat-ed, and thus affects their malleability. On theother hand, water has a high specific heat capac-ity, as discussed below; indeed, if it did not, lifeon Earth would hardly be possible.

One of the many unique properties of wateris its very high specific heat capacity, which iseasily derived from the value of a kilocalorie: it is4.184, the same number of joules required toequal a calorie. Few substances even come closeto this figure. At the low end of the spectrum arelead, gold, and mercury, with specific heat capac-ities of 0.13, 0.13, and 0.14 respectively. Alu-minum has a specific heat capacity of 0.89, andethyl alcohol of 2.43. The value for concrete, oneof the highest for any substance other than water,is 2.9.

As high as the specific heat capacity of con-crete is, that of water is more than 40% higher.On the other hand, water in its vapor state(steam) has a much lower specific heat capaci-ty—2.01. The same is true for solid water, or ice,with a specific heat capacity of 2.03. Nonetheless,water in its most familiar form has an astound-ingly high specific heat capacity, and this has sev-eral effects in the real world.

EFFECTS OF WATER’S HIGH

SPECIFIC HEAT CAPACITY. Forinstance, water is much slower to freeze in thewinter than most substances. Furthermore, dueto other unusual aspects of water—primarily thefact that it actually becomes less dense as asolid—the top of a lake or other body of waterfreezes first. Because ice is a poor medium for theconduction of heat (a consequence of its specificheat capacity), the ice at the top forms a layer thatprotects the still-liquid water below it from los-ing heat. As a result, the water below the ice layerdoes not freeze, and the animal and plant life inthe lake is preserved.

Conversely, when the weather is hot, water isslow to experience a rise in temperature. For thisreason, a lake or swimming pool makes a goodplace to cool off on a sizzling summer day. Giventhe high specific heat capacity of water, com-bined with the fact that much of Earth’s surface iscomposed of water, the planet is far less suscepti-ble than other bodies in the Solar System to vari-ations in temperature.

The same is true of another significant natu-ral feature, one made mostly of water: the humanbody. A healthy human temperature is 98.6°F

(37°C), and, even in cases of extremely highfever, an adult’s temperature rarely climbs bymore than 5°F (2.7°C). The specific heat capacityof the human body, though it is of course lowerthan that of water itself (since it is not entirelymade of water), is nonetheless quite high: 3.47.

Calorimetry

The measurement of heat gain or loss as a resultof physical or chemical change is called calorime-try (pronounced kal-or-IM-uh-tree). Like theword “calorie,” the term is derived from a Latinroot word meaning “heat.” The foundations ofcalorimetry go back to the mid-nineteenth cen-tury, but the field owes much to the work of sci-entists about 75 years prior to that time.

In 1780, French chemist Antoine Lavoisier(1743-1794) and French astronomer and mathe-matician Pierre Simon Laplace (1749-1827) hadused a rudimentary ice calorimeter for measur-ing heat in the formations of compounds.Around the same time, Scottish chemist JosephBlack (1728-1799) became the first scientist tomake a clear distinction between heat and tem-perature.

By the mid-1800s, a number of thinkers hadcome to the realization that—contrary to pre-vailing theories of the day—heat was a form ofenergy, not a type of material substance. (Thebelief that heat was a material substance, called“phlogiston,” and that phlogiston was the part ofa substance that burned in combustion, had orig-inated in the seventeenth century. Lavoisier wasthe first scientist to successfully challenge thephlogiston theory.) Among these were Ameri-can-British physicist Benjamin Thompson,Count Rumford (1753-1814) and Englishchemist James Joule (1818-1889)—for whom, ofcourse, the joule is named.

Calorimetry as a scientific field of studyactually had its beginnings with the work ofFrench chemist Pierre-Eugene Marcelin Berth-elot (1827-1907). During the mid-1860s, Berth-elot became intrigued with the idea of measuringheat, and, by 1880, he had constructed the firstreal calorimeter.

CALORIMETERS. Essential tocalorimetry is the calorimeter, which can be anydevice for accurately measuring the temperatureof a substance before and after a change occurs. Acalorimeter can be as simple as a styrofoam cup.Its quality as an insulator, which makes styro-



Temperatureand Heat

foam ideal both for holding in the warmth ofcoffee and protecting the human hand fromscalding, also makes styrofoam an excellentmaterial for calorimetric testing. With a styro-foam calorimeter, the temperature of the sub-stance inside the cup is measured, a reaction isallowed to take place, and afterward, the temper-ature is measured a second time.

The most common type of calorimeter usedis the bomb calorimeter, designed to measure theheat of combustion. Typically, a bomb calorime-ter consists of a large container filled with water,into which is placed a smaller container, the com-bustion crucible. The crucible is made of metal,with thick walls into which is cut an opening toallow the introduction of oxygen. In addition, thecombustion crucible is designed to be connectedto a source of electricity.

In conducting a calorimetric test using abomb calorimeter, the substance or object to bestudied is placed inside the combustion crucibleand ignited. The resulting reaction usually occursso quickly that it resembles the explosion of abomb—hence the name “bomb calorimeter.”Once the “bomb” goes off, the resulting transferof heat creates a temperature change in the water,which can be readily gauged with a thermometer.

To study heat changes at temperatures high-er than the boiling point of water, physicists usesubstances with higher boiling points. For exper-iments involving extremely large temperatureranges, an aneroid (without liquid) calorimetermay be used. In this case, the lining of the com-bustion crucible must be of a metal, such as cop-per, with a high coefficient or factor of thermalconductivity—that is, the ability to conduct heatfrom molecule to molecule.

Temperature in Chemistry

THE GAS LAWS. A collection ofstatements regarding the behavior of gases, thegas laws are so important to chemistry that a sep-arate essay is devoted to them elsewhere. Severalof the gas laws relate temperature to pressure andvolume for gases. Indeed, gases respond tochanges in temperature with dramatic changes involume; hence the term “volume,” when used inreference to a gas, is meaningless unless pressureand temperature are specified as well.

Among the gas laws, Boyle’s law holds that inconditions of constant temperature, an inverserelationship exists between the volume and pres-

sure of a gas: the greater the pressure, the less thevolume, and vice versa. Even more relevant to thesubject of thermal expansion is Charles’s law,which states that when pressure is kept constant,there is a direct relationship between volume andabsolute temperature.

CHEMICAL EQUILIBRIUM AND

CHANGES IN TEMPERATURE. Justas two systems that exchange no heat are said tobe in a state of thermal equilibrium, chemicalequilibrium describes a dynamic state in whichthe concentration of reactants and productsremains constant. Though the concentrations ofreactants and products do not change, note thatchemical equilibrium is a dynamic state—inother words, there is still considerable molecularactivity, but no net change.

Calculations involving chemical equilibriummake use of a figure called the equilibrium con-stant (K). According to Le Châtelier’s principle,named after French chemist Henri Le Châtelier(1850-1936), whenever a stress or change isimposed on a chemical system in equilibrium,the system will adjust the amounts of the varioussubstances in such a way as to reduce the impactof that stress. An example of a stress is a changein temperature, which changes the equilibriumequation by shifting K (itself dependant on tem-perature).

Using Le Châtelier’s law, it is possible todetermine whether K will change in the directionof the forward or reverse reaction. In an exother-mic reaction (a reaction that produces heat), Kwill shift to the left, or in the direction of the for-ward reaction. On the other hand, in anendothermic reaction (a reaction that absorbsheat), K will shift to the right, or in the directionof the reverse reaction.

TEMPERATURE AND REAC-

TION RATES. Another important functionof temperature in chemical processes is its func-tion of speeding up chemical reactions. Anincrease in the concentration of reacting mole-cules, naturally, leads to a sped-up reaction,because there are simply more molecules collid-ing with one another. But it is also possible tospeed up the reaction without changing the con-centration.

By definition, wherever a temperatureincrease is involved, there is always an increase inaverage molecular translational energy. Whentemperatures are high, more molecules are col-



Temperatureand Heat


ABSOLUTE ZERO: The temperature,

defined as 0K on the Kelvin scale, at which

the motion of molecules in a solid virtual-

ly ceases. The third law of thermodynamics

establishes the impossibility of actually

reaching absolute zero.

CALORIE: A measure of specific heat

capacity in the SI or metric system, equal

to the heat that must be added to or

removed from 1 gram of water to change

its temperature by 1°C. The dietary Calorie

(capital C), with which most people are

familiar, is the same as the kilocalorie.

CALORIMETRY: The measurement of

heat gain or loss as a result of physical or

chemical change.

CELSIUS SCALE: The metric scale of

temperature, sometimes known as the

centigrade scale, created in 1742 by

Swedish astronomer Anders Celsius (1701-

1744). The Celsius scale establishes the

freezing and boiling points of water at 0°

and 100° respectively. To convert a temper-

ature from the Celsius to the Fahrenheit

scale, multiply by 9/5 and add 32. Though

the worldwide scientific community uses

the metric or SI system for most measure-

ments, scientists prefer the related Kelvin

scale of absolute temperature.

CONSERVATION OF ENERGY: A

law of physics which holds that within a

system isolated from all other outside fac-

tors, the total amount of energy remains

the same, though transformations of ener-

gy from one form to another take place.

The first law of thermodynamics is the

same as the conservation of energy.

ENERGY: The ability to accomplish

work—that is, the exertion of force over a

given distance to displace or move an

object.

ENTROPY: The tendency of natural

systems toward breakdown, and specifical-

ly the tendency for the energy in a system

to be dissipated. Entropy is closely related

to the second law of thermodynamics.

FAHRENHEIT SCALE: The oldest of

the temperature scales still in use, created

in 1714 by German physicist Daniel

Fahrenheit (1686-1736). The Fahrenheit

scale establishes the freezing and boiling

points of water at 32° and 212° respective-

ly. To convert a temperature from the

Fahrenheit to the Celsius scale, subtract 32

and multiply by 5/9.

FIRST LAW OF THERMODYNAMICS:

A law which states the amount of energy in

a system remains constant, and therefore it

is impossible to perform work that results

in an energy output greater than the ener-

gy input. This is the same as the conserva-

tion of energy.

HEAT: Internal thermal energy that

flows from one body of matter to another.

JOULE: The principal unit of energy—

and thus of heat—in the SI or metric sys-

tem, corresponding to 1 newton-meter

(N • m). A joule (J) is equal to 0.7376 foot-

pounds in the English system.

KELVIN SCALE: Established by

William Thompson, Lord Kelvin (1824-

1907), the Kelvin scale measures tempera-

ture in relation to absolute zero, or 0K.

(Units in the Kelvin system, known as

Kelvins, do not include the word or symbol

KEY TERMS


Temperatureand Heat


for degree.) The Kelvin scale, which is the

system usually favored by scientists, is

directly related to the Celsius scale; hence

Celsius temperatures can be converted to

Kelvins by adding 273.15.

KILOCALORIE: A measure of specific

heat capacity in the SI or metric system,

equal to the heat that must be added to or

removed from 1 kilogram of water to

change its temperature by 1°C. As its name

suggests, a kilocalorie is 1,000 calories. The

dietary Calorie (capital C) with which

most people are familiar is the same as the

kilocalorie.

KINETIC ENERGY: The energy that

an object possesses by virtue of its motion.

MOLECULAR TRANSLATIONAL EN-

ERGY: The kinetic energy in a system

produced by the movement of molecules

in relation to one another. Thermal energy

is a manifestation of molecular transla-

tional energy.

SECOND LAW OF THERMODYNAM-

ICS: A law of thermodynamics which

states that no system can simply take heat

from a source and perform an equivalent

amount of work. This is a result of the fact

that the natural flow of heat is always from

a high-temperature reservoir to a low-tem-

perature reservoir. In the course of such a

transfer, some of the heat will always be

lost—an example of entropy. The second

law is sometimes referred to as “the law of

entropy.”

SPECIFIC HEAT CAPACITY: The

amount of heat that must be added to, or

removed from, a unit of mass of a given

substance to change its temperature by

1°C. It is typically measured in J/g • °C

(joules per gram-degree Celsius). A calorie

is the specific heat capacity of 1 gram of

water.

SYSTEM: In chemistry and other sci-

ences, the term “system” usually refers to

any set of interactions isolated from the

rest of the universe. Anything outside of

the system, including all factors and forces

irrelevant to a discussion of that system, is

known as the environment.

THERMAL ENERGY: Heat energy

resulting from internal kinetic energy.

THERMAL EQUILIBRIUM: A situa-

tion in which two systems have the same

temperature. As a result, there is no

exchange of heat between them.

THERMODYNAMICS: The study of

the relationships between heat, work, and

energy.

THERMOMETER: A device that gauges

temperature by measuring a temperature-

dependent property, such as the expansion

of a liquid in a sealed tube, or resistance to

electric current.

THERMOMETRIC MEDIUM: A sub-

stance whose physical properties change

with temperature. A mercury or alcohol

thermometer measures such changes.

THIRD LAW OF THERMODYNAMICS:

A law of thermodynamics stating that at

the temperature of absolute zero, entropy

also approaches zero. Zero entropy contra-

dicts the second law of thermodynamics,

meaning that absolute zero is therefore

impossible to reach.

KEY TERMS CONT INUED


Temperatureand Heat

liding, and the collisions that occur are moreenergetic. The likelihood is therefore increasedthat any particular collision will result in theenergy necessary to break chemical bonds, andthus bring about the rearrangements in mole-cules needed for a reaction.

WHERE TO LEARN MORE

“About Temperature” (Web site). <http://www.unidata.ucar.edu/staff/blynds/tmp.html> (April 18, 2001).

“Basic Chemical Thermodynamics” (Web site).<http://www.sunderland.ac.uk/~hs0bcl/td1.htm>(May 8, 2001).

“Chemical Thermodynamics” (Web site).<http://www.shodor.org/UNChem/advanced/thermo/> (May 8, 2001).

Ebbing, Darrell D.; R. A. D. Wentworth; and James P.

Birk. Introductory Chemistry. Boston: Houghton

Mifflin, 1995.

Gardner, Robert. Science Projects About Methods of Mea-

suring. Berkeley Heights, NJ: Enslow Publishers,

2000.

MegaConverter 2 (Web site). <http://www.

megaconverter.com> (May 7, 2001).

Royston, Angela. Hot and Cold. Chicago: Heinemann

Library, 2001.

Santrey, Laurence. Heat. Illustrated by Lloyd Birming-

ham. Mahwah, N.J.: Troll Associates, 1985.

Suplee, Curt. Everyday Science Explained. Washington,

D.C.: National Geographic Society, 1996.

Zumdahl, Steven S. Introductory Chemistry: A Founda-

tion, 4th ed. Boston: Houghton Mifflin, 2000.




MASS , DENS I T Y, AND VOLUMEMass, Density, and Volume

CONCEPTAmong the physical properties studied bychemists and other scientists, mass is one of themost fundamental. All matter, by definition, hasmass. Mass, in turn, plays a role in two propertiesimportant to the study of chemistry: density andvolume. All of these—mass, density, and vol-ume—are simple concepts, yet in order to workin chemistry or any of the other hard sciences, itis essential to understand these types of measure-ment. Measuring density, for instance, aids indetermining the composition of a given sub-stance, while volume is a necessary component tousing the gas laws.

HOW IT WORKS

Fundamental Properties in Relation to Volume and Density

Most qualities of the world studied by scientistscan be measured in terms of one or more of fourproperties: length, mass, time, and electriccharge. The volume of a cube, for instance, is aunit of length cubed—that is, length multipliedby “width,” which is then multiplied by “height.”

Width and height are not, for the purposesof science, distinct from length: they are simplyversions of it, distinguished by their orientationin space. Length provides one dimension, whilewidth provides a second perpendicular to thethird. Height, perpendicular both to length andwidth, makes the third spatial dimension—yet allof these are merely expressions of length differ-entiated according to direction.

VOLUME AND DENSITY DE-

FINED. Volume, then, is measured in terms oflength, and can be defined as the amount ofthree-dimensional space an object occupies. Vol-ume is usually expressed in cubic units oflength—for example, the milliliter (mL), alsoknown as the cubic centimeter (cc), is equal to6.10237 • 10–2 in3. As its name implies, there are1,000 milliliters in a liter.

Density is the ratio of mass to volume—or,to put its definition in terms of fundamentalproperties, of mass divided by cubed length.Density can also be viewed as the amount ofmatter within a given area. In the SI system, den-sity is typically expressed as grams per cubic cen-timeter (g/cm3), equivalent to 62.42197 poundsper cubic foot in the English system.

Mass

MASS DEFINED. Though length iseasy enough to comprehend, mass is moreinvolved. In his second law of motion, Sir IsaacNewton (1642-1727) defined mass as the ratio offorce to gravity. This, of course, is a statementthat belongs to the realm of physics; for achemist, it is more useful—and also accurate—todefine mass as the quantity of matter that anobject contains.

Matter, in turn, can be defined as physicalsubstance that occupies space; is composed ofatoms (or in the case of subatomic particles, ispart of an atom); is convertible into energy—andhas mass. The form or state of matter itself is notimportant: on Earth it is primarily observed as asolid, liquid, or gas, but it can also be found (par-ticularly in other parts of the universe) in afourth state, plasma.


Mass, Density, and Volume


The more matter an object contains, the

more mass. To refer again to the laws of motion,

in his first law, Newton identified inertia: the ten-

dency of objects in motion to remain in motion,

or of objects at rest to remain at rest, in a con-

stant velocity unless they are acted upon by some

outside force. Mass is a measure of inertia, mean-

ing that the more mass something contains, the

more difficult it is to put it into motion, or to

stop it from moving.

MASS VS. WEIGHT. Most people

who are not scientifically trained tend to think

that mass and weight are the same thing, but this

is like saying that apples and apple pies are the

same. Of course, an apple is an ingredient in an

apple pie, but the pie contains something else—

actually, a number of other things, such as flour

and sugar. In this analogy, mass is equivalent to

the apple, and weight the pie, while the accelera-

tion due to gravity is the “something else” in

weight.

It is understandable why people confuse

mass with weight, since most weight scales pro-

vide measurements in both pounds and kilo-

grams. However, the pound (lb) is a unit of

weight in the English system, whereas a kilogram

ASTRONAUTS NEIL ARMSTRONG AND BUZZ ALDRIN, THE FIRST MEN TO WALK ON THE MOON, WEIGHED LESS ON

THE MOON THAN ON EARTH. THE REASON IS BECAUSE WEIGHT DIFFERS AS A RESPONSE TO THE GRAVITATIONAL PULL

OF THE PLANET, MOON, OR OTHER BODY ON WHICH IT IS MEASURED. THUS, A PERSON WEIGHS LESS ON THE MOON,BECAUSE THE MOON POSSESSES LESS MASS THAN EARTH AND EXERTS LESS GRAVITATIONAL FORCE. (NASA/Roger Ress-

meyer/Corbis. Reproduced by permission.)


Mass, Density,

and Volume


(kg) is a unit of mass in the metric and SI sys-tems. Though the two are relatively convertibleon Earth (1 lb = 0.4536 kg; 1 kg = 2.21 lb), theyare actually quite different.

Weight is a measure of force, which New-ton’s second law of motion defined as the prod-uct of mass multiplied by acceleration. The accel-eration component of weight is a result of Earth’sgravitational pull, and is equal to 32 ft (9.8 m)per second squared. Thus a person’s weight variesaccording to gravity, and would be different ifmeasured on the Moon; mass, on the other hand,is the same throughout the universe. Given itsinvariable value, scientists typically speak interms of mass rather than weight.

Weight differs as a response to the gravita-tional pull of the planet, moon, or other body onwhich it is measured. Hence a person weighs lesson the Moon, because the Moon possesses lessmass than Earth, and, thus, exerts less gravita-tional force. Therefore, it would be easier on theMoon to lift a person from the ground, but itwould be no easier to move that person from aresting position, or to stop him or her frommoving. This is because the person’s mass, andhence his or her resistance to inertia, has notchanged.


Atomic Mass Units

Chemists do not always deal in large units ofmass, such as the mass of a human body—which,of course, is measured in kilograms. Instead,the chemist’s work is often concerned with measurements of mass for the smallest types of matter: molecules, atoms, and other ele-mentary particles. To measure these even interms of grams (0.001 kg) is absurd: a singleatom of carbon, for instance, has a mass of1.99 • 10–23 g. In other words, a gram is about50,000,000,000,000,000,000,000 times largerthan a carbon atom—hardly a usable com-parison.

Instead, chemists use an atom mass unit(abbreviated amu), which is equal to 1.66 • 10–24

g. Even so, is hard to imagine determining themass of single atoms on a regular basis, sochemists make use of figures for the averageatomic mass of a particular element. The averageatomic mass of carbon, for instance, is 12.01amu. As is the case with any average, this meansthat some atoms—different isotopes of carbon—may weigh more or less, but the figure of 12.01

TO DETERMINE WHETHER A PIECE OF GOLD IS GENUINE OR FAKE, ONE MUST MEASURE THE DENSITY OF THE SUB-STANCE. HERE, A MAN EVALUATES GOLD PIECES IN HONG KONG. (Christophe Loviny/Corbis. Reproduced by permission.)




amu is still reliable. Some other average atomicmass figures for different elements are as follows:

• Hydrogen (H): 1.008 amu

• Helium (He): 4.003 amu

• Lithium (Li): 6.941 amu

• Nitrogen (N): 14.01 amu

• Oxygen (O): 16.00

• Aluminum (Al): 26.98

• Chlorine (Cl): 35.46 amu

• Gold (Au): 197.0 amu

• Hassium (Hs): [265 amu]

The figure for hassium, with an atomicnumber of 108, is given in brackets because thisnumber is the mass for the longest-lived isotope.The average value of mass for the molecules in agiven compound can also be rendered in terms ofatomic mass units: water (H2O) molecules, forinstance, have an average mass of 18.0153 amu.Molecules of magnesium oxide (MgO), whichcan be extracted from sea water and used in mak-ing ceramics, have an average mass much higherthan for water: 40.304 amu.

These values are obtained simply by addingthose of the atoms included in the molecule:since water has two hydrogen atoms and oneoxygen, the average molecular mass is obtainedby multiplying the average atomic mass ofhydrogen by two, and adding it to the average

atomic mass of oxygen. In the case of magnesiumoxide, the oxygen is bonded to just one otheratom—but magnesium, with an average atomicmass of 24.304, weighs much more than hydrogen.

Molar Mass

It is often important for a chemist to know exact-ly how many atoms are in a given sample, partic-ularly in the case of a chemical reaction betweentwo or more samples. Obviously, it is impossibleto count atoms or other elementary particles, butthere is a way to determine whether two items—regardless of the elements or compoundsinvolved—have the same number of elementaryparticles. This method makes use of the figuresfor average atomic mass that have been estab-lished for each element.

If the average atomic mass of the substanceis 5 amu, then there should be a very large num-ber of atoms (if it is an element) or molecules (ifit is a compound) of that substance having a totalmass of 5 grams (g). Similarly, if the averageatomic mass of the substance is 7.5 amu, thenthere should be a very large number of atoms ormolecules of that substance having a total massof 7.5 g. What is needed, clearly, is a very largenumber by which elementary particles must be

Earth(to same scale)

ALTHOUGH SATURN IS MUCH LARGER THAN EARTH, IT IS MUCH LESS DENSE.


Mass, Density,

and Volume

multiplied in order to yield a mass whose valuein grams is equal to the value, in amu, of its aver-age atomic mass. This is known as Avogadro’snumber.

AVOGADRO’S NUMBER. The firstscientist to recognize a meaningful distinctionbetween atoms and molecules was Italian physi-cist Amedeo Avogadro (1776-1856). Avogadromaintained that gases consisted of particles—which he called molecules—that in turn consist-ed of one or more smaller particles. He furtherreasoned that one liter of any gas must con-tain the same number of particles as a liter ofanother gas.

In order to discuss the behavior of mole-cules, it was necessary to set a large quantity as abasic unit, since molecules themselves are verysmall. This led to the establishment of what isknown as Avogadro’s number, equal to 6.022137� 1023 (more than 600 billion trillion.)

The magnitude of Avogadro’s number isalmost inconceivable. The same number ofgrains of sand would cover the entire surface ofEarth at a depth of several feet. The same num-ber of seconds, for instance, is about 800,000times as long as the age of the universe (20 billionyears). Avogadro’s number—named after theman who introduced the concept of the mole-cule, but only calculated years after his death—serves a very useful purpose in computationsinvolving molecules.

THE MOLE. To compare two substancescontaining the same number of atoms or mole-cules, scientists use the mole, the SI fundamentalunit for “amount of substance.” A mole (abbrevi-ated mol) is, generally speaking, Avogadro’snumber of atoms or molecules; however, in themore precise SI definition, a mole is equal to thenumber of carbon atoms in 12.01 g (0.03 lb) ofcarbon. Note that, as stated earlier, carbon has anaverage atomic mass of 12.01 amu. This is nocoincidence, of course: multiplication of theaverage atomic mass by Avogadro’s numberyields a figure in grams equal to the value of theaverage atomic mass in amu.

The term “mole” can be used in the sameway we use the word “dozen.” Just as “a dozen”can refer to twelve cakes or twelve chickens, so“mole” always describes the same number ofmolecules. Just as one liter of water, or one literof mercury, has a certain mass, a mole of anygiven substance has its own particular mass,

expressed in grams. A mole of helium, forinstance, has a mass of 4.003 g (0.01 lb), whereasa mole of iron is 55.85 g (0.12 lb) These figuresrepresent the molar mass for each: that is, themass of 1 mol of a given substance.

Once again, the value of molar mass ingrams is the same as that of the average atomicmass in amu. Also, it should be clear that, giventhe fact that helium weighs much less than air—the reason why helium-filled balloons float—aquantity of helium with a mass of 4.003 g mustbe a great deal of helium. And indeed, as indicat-ed earlier, the quantity of atoms or molecules ina mole is sufficiently great to make a sample thatis large, but still usable for the purposes of studyor comparison.

Measuring Volume

Mass, because of its fundamental nature, issometimes hard to comprehend, and densityrequires an explanation in terms of mass and vol-ume. Volume, on the other hand, appears to bequite straightforward—and it is, when one isdescribing a solid of regular shape. In other situ-ations, however, volume measurement is morecomplicated.

As noted earlier, the volume of a cube can beobtained simply by multiplying length by widthby height. There are other means for measuringthe volume of other straight-sided objects, suchas a pyramid. Still other formulae, which makeuse of the constant π (roughly equal to 3.14) arenecessary for measuring the volume of a cylinder,a sphere, or a cone.

For an object that is irregular in shape, how-ever, one may have to employ calculus—but themost basic method is simply to immerse theobject in water. This procedure involves measur-ing the volume of the water before and afterimmersion, and calculating the difference. Ofcourse, the object being measured cannot bewater-soluble; if it is, its volume must be meas-ured in a non-water-based liquid such as alcohol.

LIQUID AND GAS VOLUME.

Measuring liquid volumes is even easier than forsolids, given the fact that liquids have no definiteshape, and will simply take the shape of the con-tainer in which they are placed. Gases are similarto liquids in the sense that they expand to fit theircontainer; however, measurement of gas volumeis a more involved process than that used tomeasure either liquids or solids, because gases are




highly responsive to changes in temperature andpressure.

If the temperature of water is raised from itsfreezing point to its boiling point—from 32°F(0°C) to 212°F (100°C)—its volume will increaseby only 2%. If its pressure is doubled from 1 atm(defined as normal air pressure at sea level) to 2atm, volume will decrease by only 0.01%. Yet ifair were heated from 32° to 212°F, its volumewould increase by 37%; if its pressure were dou-bled from 1 atm to 2, its volume would decreaseby 50%.

Not only do gases respond dramatically tochanges in temperature and pressure, but also,gas molecules tend to be non-attractive towardone another—that is, they tend not to sticktogether. Hence, the concept of “volume” in rela-tion to a gas is essentially meaningless unless itstemperature and pressure are known.

Comparing Densities

In the discussion of molar mass above, heliumand iron were compared, and we saw that themass of a mole of iron was about 14 times asgreat as that of a mole of helium. This may seemlike a fairly small factor of difference betweenthem: after all, helium floats on air, whereas iron(unless it is arranged in just the right way, forinstance, in a tanker) sinks to the bottom of theocean. But be careful: the comparison of molarmass is only an expression of the mass of a heli-um atom as compared to the mass of an ironatom. It makes no reference to density, which isthe ratio of mass to volume.

Expressed in terms of the ratio of mass tovolume, the difference between helium and ironbecomes much more pronounced. Suppose, onthe one hand, one had a gallon jug filled withiron. How many gallons of helium does it take toequal the mass of the iron? Fourteen? Try again:it takes more than 43,000 gallons of helium toequal the mass of the iron in one gallon jug!Clearly, what this shows is that the density of ironis much, much greater than that of helium.

This, of course, is hardly a surprising revela-tion; still, it is sometimes easy to get confused bycomparisons of mass as opposed to comparisonsof density. One might even get tricked by the oldelementary-school brain-teaser that goes some-thing like this: “Which is heavier, a ton of feath-ers or a ton of cannonballs?” Of course neither isheavier, but the trick element in the question

relates to the fact that it takes a much greater vol-ume of feathers (measured in cubic feet, forinstance) than of cannonballs to equal a ton.

One of the interesting things about density,as distinguished from mass and volume, is that ithas nothing to do with the amount of material. Akilogram of iron differs from 10 kg of iron bothin mass and volume, but the density of both sam-ples is the same. Indeed, as discussed below, theknown densities of various materials make itpossible to determine whether a sample of thatmaterial is genuine.

COMPARING DENSITIES. Asnoted several times, the densities of numerousmaterials are known quantities, and can be easilycompared. Some examples of density, allexpressed in terms of grams per cubic centime-ter, are listed below. These figures are measuredat a temperature of 68°F (20°C), and for hydro-gen and oxygen, the value was obtained at nor-mal atmospheric pressure (1 atm):

Comparisons of Densities for Various Sub-stances:

• Oxygen: 0.00133 g/cm3

• Hydrogen: 0.000084 g/cm3

• Ethyl alcohol: 0.79 g/cm3

• Ice: 0.920 g/cm3

• Water: 1.00 g/cm3

• Concrete: 2.3 g/cm3

• Iron: 7.87 g/cm3

• Lead: 11.34 g/cm3

• Gold: 19.32 g/cm3

Specific Gravity

IS IT REALLY GOLD? Note thatpure water (as opposed to sea water, which is 3%more dense) has a density of 1.0 g per cubic cen-timeter. Water is thus a useful standard for meas-uring the specific gravity of other substances, orthe ratio between the density of that substanceand the density of water. Since the specific gravi-ty of water is 1.00—also the density of water ing/cm3—the specific gravity of any substance (anumber, rather than a number combined with aunit of measure) is the same as the value of itsown density in g/cm3.

Comparison of densities make it possible todetermine whether a piece of jewelry alleged tobe solid gold is really genuine. To determine theanswer, one must drop the sample in a beaker ofwater with graduated units of measure clearly



Mass, Density,

and Volume


ATOMIC MASS UNIT: An SI unit

(abbreviated amu), equal to 1.66 • 10-24 g,

for measuring the mass of atoms.

AVERAGE ATOMIC MASS: A figure

used by chemists to specify the mass—in

atomic mass units—of the average atom in

a large sample. The average atomic mass of

carbon, for instance, is 12.01 amu. If a sub-

stance is a compound, the average atomic

mass of all atoms in a molecule of that sub-

stance must be added together to yield the

average molecular mass of that substance.

AVOGADRO’S NUMBER: A figure,

named after Italian physicist Amedeo Avo-

gadro (1776-1856), equal to 6.022137 �

�023. Avogadro’s number indicates the

number of atoms, molecules, or other ele-

mentary particles in a mole.

DENSITY: The ratio of mass to vol-

ume—in other words, the amount of mat-

ter within a given area. In the SI system,

density is typically expressed as grams per

cubic centimeter (g/cm3), equal to

62.42197 pounds per cubic foot in the Eng-

lish system.

MASS: The amount of matter an object

contains.

MATTER: Physical substance that occu-

pies space, has mass, is composed of atoms

(or in the case of subatomic particles, is

part of an atom), and is convertible to

energy.

MILLILITER: One of the most com-

monly used units of volume in the SI sys-

tem of measures. The milliliter (abbreviat-

ed mL), also known as a cubic centimeter

(cc), is equal to 6.10237 • 10-2 cubic inches

in the English system. As the name implies,

there are 1,000 milliliters in a liter.

MOLAR MASS: The mass, in grams, of

1 mole of a given substance. The value in

grams of molar mass is always equal to the

value, in atomic mass units, of the average

atomic mass of that substance: thus, car-

bon has a molar mass of 12.01 g, and an

average atomic mass of 12.01 amu.

MOLE: The SI fundamental unit for

“amount of substance.” A mole is, general-

ly speaking, Avogadro’s number of atoms,

molecules, or other elementary particles;

however, in the more precise SI definition,

a mole is equal to the number of carbon

atoms in 12.01 g of carbon.

SPECIFIC GRAVITY: The density of

an object or substance relative to the densi-

ty of water; or more generally, the ratio

between the densities of two objects or

substances. Since the specific gravity of

water is 1.00—also the density of water in

g/cm3—the specific gravity of any sub-

stance is the same as the value of its own

density in g/cm3. Specific gravity is simply

a number, without any unit of measure.

VOLUME: The amount of three-

dimensional space an object occupies. Vol-

ume is usually expressed in cubic units of

length—for instance, the milliliter.

WEIGHT: The product of mass multi-

plied by the acceleration due to gravity (32

ft or 9.8 m/sec2). A pound is a unit of

weight, whereas a kilogram is a unit of

mass.

KEY TERMS



marked. Suppose the item has a mass of 10 g. Thedensity of gold is 19 g/cm3, and because densityis equal to mass divided by volume, the volumeof water displaced should be equal to the massdivided by the density. The latter figure is equalto 10 g divided by 19 g/cm3, or 0.53 ml. Supposethat instead, the item displaced 0.88 ml of water.Clearly it is not gold, but what is it?

Given the figures for mass and volume, itsdensity is equal to 11.34 g/cm3—which happensto be the density of lead. If, on the other hand,the amount of water displaced were somewherebetween the values for pure gold and pure lead,one could calculate what portion of the item wasgold and which lead. It is possible, of course, thatit could contain some other metal, but given thehigh specific gravity of lead, and the fact that its density is relatively close to that of gold,lead is a favorite gold substitute among jewelrycounterfeiters.

SPECIFIC GRAVITY AND THE

DENSITIES OF PLANETS. Mostrocks near the surface of Earth have a specificgravity somewhere between 2 and 3, while thespecific gravity of the planet itself is about 5.How do scientists know that the density of Earthis around 5 g/cm3? The computation is fairly sim-ple, given the fact that the mass and volume ofthe planet are known. And given the fact thatmost of what lies close to Earth’s surface—seawater, soil, rocks—has a specific gravity wellbelow 5, it is clear that Earth’s interior must con-tain high-density materials, such as nickel oriron. In the same way, calculations regarding thedensity of other objects in the Solar System pro-vide a clue as to their interior composition.

This brings the discussion back around to atopic raised much earlier in this essay, whencomparing the weight of a person on Earth ver-sus that person’s weight on the Moon. It so hap-pens that the Moon is smaller than Earth, butthat is not the reason it exerts less gravitationalpull: as noted earlier, the gravitational force a

planet, moon, or other body exerts is related toits mass, not its size.

It so happens, too, that Jupiter is much larg-er than Earth, and that it exerts a gravitationalpull much greater than that of Earth. This isbecause it has a mass many times as great asEarth’s. But what about Saturn, the second-largest planet in the Solar System? In size it isonly about 17% smaller than Jupiter, and bothare much, much larger than Earth. Yet a personwould weigh much less on Saturn than onJupiter, because Saturn has a mass much smallerthan Jupiter’s. Given the close relation in sizebetween the two planets, it is clear that Saturnhas a much lower density than Jupiter, or in facteven Earth: the great ringed planet has a specificgravity of less than 1.

WHERE TO LEARN MORE

Chahrour, Janet. Flash! Bang! Pop! Fizz!: Exciting Sciencefor Curious Minds. Illustrated by Ann HumphreyWilliams. Hauppauge, NY: Barron’s, 2000.

“Density and Specific Gravity” (Web site). <http://www.tpub.com/fluid/ch1e.htm> (March 27, 2001).

“Density, Volume, and Cola” (Web site). <http://student.biology.arizona.edu/sciconn/density/density_coke.html> (March 27, 2001).

Ebbing, Darrell D.; R. A. D. Wentworth; and James P.Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.

“The Mass Volume Density Challenge” (Web site).<http://science-math-technology.com/mass_volume_density.html> (March 27, 2001).

“MegaConverter 2” (Web site). <http://www.megaconverter.com> (May 7, 2001).

“Metric Density and Specific Gravity” (Web site).<http://www.essex1.com/people/speer/density.html>(March 27, 2001).

Robson, Pam. Clocks, Scales and Measurements. NewYork: Gloucester Press, 1993.

“Volume, Mass, and Density” (Web site).<http://www.nyu.edu/pages/mathmol/modules/water/density_intro.html> (March 27, 2001).

Zumdahl, Steven S. Introductory Chemistry: A Founda-tion, 4th ed. Boston: Houghton Mifflin, 2000.



31

SC IENCE OF EVERYDAY TH INGSr e a l - l i f e c h em i s t r y

MAT TERMAT TER

PROPERT IES OF MATTER

GASES

set_vol1_sec2 10/5/01 2:42 PM Page 31


PROPERT I ES OF MAT TERProperties of Matter

CONCEPT

Matter is physical substance that occupies space,has mass, is composed of atoms—or, in the caseof subatomic particles, is part of an atom—and isconvertible to energy. On Earth, matter appearsin three clearly defined forms—solid, liquid, andgas—whose varying structural characteristics area function of the speeds at which its moleculesmove in relation to one another. A single sub-stance may exist in any of the three phases: liquidwater, for instance, can be heated to becomesteam, a vapor; or, when sufficient heat isremoved from it, it becomes ice, a solid. These aremerely physical changes, which do not affect thebasic composition of the substance itself: it is stillwater. Matter, however, can and does undergochemical changes, which (as with the variousstates or phases of matter) are an outcome ofactivity at the atomic and molecular level.

HOW IT WORKS

Matter and Energy

One of the characteristics of matter noted in itsdefinition above is that it is convertible to energy.We rarely witness this conversion, though asAlbert Einstein (1879-1955) showed with hisTheory of Relativity, it occurs in a massive way atspeeds approaching that of light.

Einstein’s famous formula, E = mc2, meansthat every item possesses a quantity of energyequal to its mass multiplied by the squared speedof light. Given the fact that light travels at186,000 mi (299,339 km) per second, the quanti-

ties of energy available from even a tiny objecttraveling at that speed are enormous indeed. Thisis the basis for both nuclear power and nuclearweaponry, each of which uses some of the small-est particles in the known universe to produceresults that are both amazing and terrifying.

Even in everyday life, it is still possible toobserve the conversion of mass to energy, if onlyon a very small scale. When a fire burns—that is,when wood experiences combustion in the pres-ence of oxygen, and undergoes chemicalchanges—a tiny fraction of its mass is convertedto energy. Likewise, when a stick of dynamiteexplodes, it too experiences chemical changesand the release of energy. The actual amount ofenergy released is, again, very small: for a stick ofdynamite weighing 2.2 lb (1 kg), the portion ofits mass that “disappears” is be equal to 6 partsout of 100 billion.

Actually, none of the matter in the fire or thedynamite blast disappears: it simply changesforms. Most of it becomes other types of mat-ter—perhaps new compounds, and certainly newmixtures of compounds. A very small part, as wehave seen, becomes energy. One of the most fun-damental principles of the universe is the conser-vation of energy, which holds that within a sys-tem isolated from all other outside factors, thetotal amount of energy remains the same, thoughtransformations of energy from one form toanother take place. In this situation, some of theenergy remains latent, or “in reserve” as matter,while other components of the energy arereleased; yet the total amount of energy remainsthe same.



Physical and ChemicalChanges

In discussing matter—as, for instance, in thecontext of matter transforming into energy—onemay speak in physical or chemical terms, or both.Generally speaking, physicists study physicalproperties and changes, while chemists are con-cerned with chemical processes and changes.

A physicist views matter in terms of its mass,temperature, mechanical properties (for exam-ple, elasticity); electrical conductivity; and otherstructural characteristics. The chemical makeupof matter, on the other hand, is of little concernto a physicist. For instance, in analyzing a fire oran explosion, the physicist is not concerned withthe interactions of combustible or explosivematerials and oxygen. The physicist’s interest,rather, is in questions such as the amount ofheat in the fire, the properties of the sound waves emitted in the explosion of the dynamite,and so on.

The changes between different states orphases of matter, as they are discussed below, arephysical changes. If water boils and vaporizes assteam, it is still water; likewise if it freezes tobecome solid ice, nothing has changed withregard to the basic chemical structure of the H2Omolecules that make up water. But if water reactswith another substance to form a new com-pound, it has undergone chemical change. Like-wise, if water molecules experience electrolysis, aprocess in which electric current is used todecompose H2O into molecules of H2 and O2,this is also a chemical change.

Similarly, a change from matter to energy,while it is also a physical change, typicallyinvolves some chemical or nuclear process toserve as “midwife” to that change. Yet physicaland chemical changes have at least one thing incommon: they can be explained in terms ofbehavior at the atomic or molecular level. This istrue of many physical processes—and of allchemical ones.

Atoms

In his highly readable Six Easy Pieces—a workthat includes considerable discussion of chem-istry as well as physics—the great Americanphysicist Richard Feynman (1918-1988) asked,“If, in some cataclysm, all of scientific knowledgewere to be destroyed, and only one sentencepassed on to the next generations of creatures,

what statement would contain the most informa-tion in the fewest words?”

The answer he gave was this: “I believe it isthe atomic hypothesis (or the atomic fact, orwhatever you wish to call it) that all things aremade of atoms—little articles that move aroundin perpetual motion, attracting each other whenthey are a little distance apart, but repelling uponbeing squeezed into one another. In that sen-tence, you will see, there is an enormous amountof information about the world, if just a littleimagination and thinking are applied.”

Indeed, what Feynman called the “atomichypothesis” is one of the most important keys tounderstanding both physical and chemicalchanges. The behavior of particles at the atomiclevel has a defining role in the shape of the worldstudied by the sciences, and an awareness of thisbehavior makes it easier to understand physicalprocesses, such as changes of state between solid,liquid, and gas; chemical processes, such as theformation of new compounds; and otherprocesses, such as the conversion of matter toenergy, which involve both physical and chemicalchanges. Only when one comprehends the atom-ic structure of matter is it possible to move on tothe chemical elements that are the most basicmaterials of chemistry.

STRUCTURE OF THE ATOM. AsFeynman went on to note, atoms are so tiny thatif an apple were magnified to the size of Earth,the atoms in it would each be about the size of aregular apple. Clearly, atoms and other atomicparticles are far too small to be glimpsed even bythe most highly powered optical microscope. Yetphysicists and other scientists are able to studythe behavior of atoms, and by doing so, they areable to form a picture of what occurs at theatomic level.

An atom is the fundamental particle in achemical element. The atom is not, however, thesmallest particle in the universe: atoms are com-posed of subatomic particles, including protons,neutrons, and electrons. These are distinguishedfrom one another in terms of electric charge: aswith the north and south poles of magnets, pos-itive and negative charges attract one another,but like charges repel. (In fact, magnetism is sim-ply a manifestation of a larger electromagneticforce that encompasses both electricity and magnetism.)

Propertiesof Matter


Propertiesof Matter

Clustered at the center, or nucleus, of theatom are protons, which are positively charged,and neutrons, which exert no charge. Spinningaround the nucleus are electrons, which exert anegative charge. The vast majority of the atom’smass is made up by the protons and neutrons,which have approximately the same mass; that ofthe electron is much smaller. If an electron had amass of 1—not a unit, but simply a figure usedfor comparison—the mass of the proton wouldbe 1,836, and of the neutron 1,839.

Atoms of the same element always have thesame number of protons, and since this figure isunique for a given element, each element isassigned an atomic number equal to the numberof protons in its nucleus. Two atoms may havethe same number of protons, and thus be of thesame element, yet differ in their number of neu-trons. Such atoms are called isotopes.

The number of electrons is usually the sameas the number of protons, and thus atoms have aneutral charge. In certain situations, however, theatom may lose or gain one or more electrons andacquire a net charge, becoming an ion. But elec-tric charge, like energy, is conserved, and the elec-trons are not “lost” when an atom becomes anion: they simply go elsewhere.

It is useful, though far from precise, to com-pare the interior of an atom to a planet spinningvery quickly around a sun. If the nucleus wereour own Sun, then the electrons spinning at theedge of the atom would be on an orbit some-where beyond Mars: in other words, the ratiobetween the size of the nucleus and the furthestedge of the atom is like that between the Sun’sdiameter and an orbital path about 80 millionmiles beyond Mars.

One of many differences between an atomand a solar system, however, is the fact that theelectrons are spinning around the nucleus at arelative rate of motion much, much greater thanany planet is revolving around the Sun. Further-more, what holds the atom together is not gravi-tational force, as in the Solar System, but electro-magnetic force. A final and critical difference isthe fact that electrons move in much more com-plex orbital patterns than the elliptical paths thatplanets make in their movement around the Sun.

Molecules

Though an atom is the fundamental unit of mat-ter, most of the substances people encounter in


the world are not pure elements such as oxygenor iron. They are compounds in which atoms ofmore than one element join—usually in mole-cules. All molecules are composed of more thanone atom, but not necessarily of more than oneelement: oxygen, for instance, generally appearsin the form of molecules in which two oxygenatoms are bonded. Because of this, pure oxygen isrepresented by the chemical symbol O2, asopposed to the symbol for the element oxygen,which is simply O.

One of the most well-known molecularforms in the world is water, or H2O, composed oftwo hydrogen atoms and one oxygen atom. Thearrangement is extremely precise and nevervaries: scientists know, for instance, that the twohydrogen atoms join the oxygen atom (which ismuch larger than the hydrogen atoms) at anangle of 105°3’. Since the oxygen atom is muchlarger than the two hydrogens, its shape can becompared to a basketball with two softballsattached.

Other molecules are much more complexthan those of water, and some are much, muchmore complex, a fact reflected in the sometimes

A COAL GASIFICATION PLANT. COAL GASIFICATION

MAKES IT POSSIBLE TO BURN “CLEAN” COAL. (Roger Ress-

meyer/Corbis. Reproduced by permission.)


Propertiesof Matter

lengthy names and complicated symbolic repre-sentations required to identify their chemicalcomponents. On the other hand, not all materi-als are made up of molecules: salt, for instance, isan ionic solid, as discussed below.

Quantifying Atoms and Molecules

The nucleus of an atom is about 10–13 cm indiameter, and the diameter of the entire atom isabout 10–8 cm—about 0.0000003937 in. Obvi-ously, special units are required for describingthe size of atoms, and usually measurements areprovided in terms of the angstrom, equal to 10–10

m, or 10–8 cm. To put this on some sort of imag-inable scale, there are 10 million angstroms in amillimeter.

Measuring the spatial dimensions of anatom, however, is not as important as measuringits mass—and naturally, the mass of an atom isalso almost inconceivably small. For instance, ittakes about 5.0 • 1023 carbon atoms to equal justone gram of mass. Again, the numbers boggle themind, but the following may put this into per-spective. We have already established just howtiny an angstrom is; now consider the following.If 5.0 • 1023 angstrom lengths were laid end toend, they would stretch for a total of about107,765 round trips from Earth to the Sun!

It is obvious, then, that an entirely differentunit should be used for measuring the mass of anatom, and for this purpose, chemists and otherscientists use an atom mass unit (abbreviatedamu). The latter is equal to 1.66 • 10–24 g. Even so,scientists can hardly be expected to be constantlymeasuring the mass of individual atoms; rather,they rely on figures determined for the averageatomic mass of a particular element.

Average atomic mass figures range from1.008 amu for hydrogen to over 250 amu for ele-ments of very high atomic numbers. Figures foraverage atomic mass can be used to determinethe average mass of a molecule as well, simply bycombining the average atomic mass figures foreach atom the molecule contains. A water mole-cule, for instance, has an average mass equal tothe average atomic mass of hydrogen multipliedby two, and added to the average atomic mass ofoxygen.

AVOGADRO’S NUMBER AND

THE MOLE. Just as using average atomicmass is much more efficient than measuring the

mass of individual atoms or molecules, scientistsneed a useful means for comparing atoms ormolecules of different substances—and for doingso in such a way that they know they are analyz-ing equal numbers of particles. This cannot bedone in terms of mass, because the number ofatoms in each sample would vary: a gram ofhydrogen, for instance, would contain about 12times as many atoms as a gram of carbon, whichhas an average atomic mass of 12.01 amu. Whatis needed, instead, is a way to designate a certainnumber of atoms or molecules, such that accu-rate comparisons are possible.

In order to do this, scientists make use of afigure known as Avogadro’s number. Named afterItalian physicist Amedeo Avogadro (1776-1856),it is equal to 6.022137 � 1023 Earlier, we estab-lished the almost inconceivable scale representedby the figure 5.0 • 1023; here we are confrontedwith a number 20% larger. But Avogadro’s num-ber, which is equal to 6,022,137 followed by 17zeroes, is more than simply a mind-bogglingseries of digits.

In general terms, Avogadro’s number desig-nates the quantity of molecules (and sometimesatoms, if the substance in question is an elementthat, unlike oxygen, appears as single atoms) in amole (abbreviated mol). A mole is the SI unit for“amount of substance,” and is defined preciselyas the number of carbon atoms in 12.01 g of car-bon. It is here that the value of Avogadro’s num-ber becomes clear: as noted, carbon has an aver-age atomic mass of 12.01 amu, and multiplica-tion of the average atomic mass by Avogadro’snumber yields a figure in grams equal to thevalue of the average atomic mass in atomic massunits.

MOLAR MASS AND DENSITY.

By comparison, a mole of helium has a molarmass of 4.003 g (0.01 lb) The molar mass of iron(that is, the mass of 1 mole of iron) is 55.85 g(0.12 lb) Note that there is not a huge ratio ofdifference between the molar mass of iron andthat of helium: iron has a molar mass about 14times greater. This, of course, seems very smallin light of the observable differences betweeniron and helium: after all, who ever heard of aballoon filled with iron, or a skyscraper withhelium girders?

The very striking differences between ironand helium, clearly, must come from somethingother than the molar mass differential between



Propertiesof Matter

them. Of course, a mole of iron contains thesame number of atoms as a mole of helium, butthis says nothing about the relative density ofthe two substances. In terms of volume—that is, the amount of space that something occu-pies—the difference is much more striking: thevolume of a mole of helium is about 43,000 timesas large as that of a mole of iron.

What this tells us is that the densities of ironand helium—the amount of mass per unit ofvolume—are very different. This difference indensity is discussed in the essay on Mass, Densi-ty, and Volume; here the focus is on a larger judg-ment that can be formed by comparing the twodensities. Helium, of course, is almost always inthe form of a gas: to change it to a solid requiresa temperature near absolute zero. And iron is asolid, meaning that it only turns into a liquid atextraordinarily high temperatures. These differ-ences in overall structure can, in turn, be attrib-uted to the relative motion, attraction, and ener-gy of the molecules in each.

Molecular Attraction and Motion

At the molecular level, every item of matter in theworld is in motion, and the rate of that motion isa function of the attraction between molecules.Furthermore, the rate at which molecules movein relation to one another determines phase ofmatter—that is, whether a particular item can bedescribed as solid, liquid, or gas. The movementof molecules generates kinetic energy, or theenergy of movement, which is manifested asthermal energy—what people call “heat” in ordi-nary language. (The difference between thermalenergy and heat is explained in the essay on Tem-perature and Heat.) In fact, thermal energy is theresult of molecules’ motion relative to one anoth-er: the faster they move, the greater the kineticenergy, and the greater the “heat.”

When the molecules in a material moveslowly—merely vibrating in place—they exert astrong attraction toward one another, and thematerial is called a solid. Molecules of liquid, bycontrast, move at moderate speeds and exert amoderate attraction. A material substance whosemolecules move at high speeds, and thereforeexert little or no attraction, is known as a gas. Inshort, the weaker the attraction, the greater therate of relative motion—and the greater theamount of thermal energy the object contains.


Types of Solids

Particles of solids resist attempts to compressthem, or push them together, and because oftheir close proximity, solid particles are fixed inan orderly and definite pattern. As a result, a solidusually has a definite volume and shape.

A crystalline solid is a type of solid in whichthe constituent parts are arranged in a simple,definite geometric pattern that is repeated in alldirections. But not all crystalline solids are thesame. Table salt is an example of an ionic solid: aform of crystalline solid that contains ions. Whenmixed with a solvent such as water, ions from thesalt move freely throughout the solution, makingit possible to conduct an electric current.

Regular table sugar (sucrose) is a molecularsolid, or one in which the molecules have a neu-tral electric charge—that is, there are no ionspresent. Therefore, a solution of water and sugarwould not conduct electricity. Finally, there arecrystalline solids known as atomic solids, inwhich atoms of one element bond to one anoth-er. Examples include diamonds (made of purecarbon), silicon, and all metals.

Other solids are said to be amorphous,meaning that they possess no definite shape.Amorphous solids—an example of which isclay—either possess very tiny crystals, or consistof several varieties of crystal mixed randomly.Still other solids, among them glass, do not con-tain crystals.

Freezing and Melting

VIBRATIONS AND FREEZING.

Because of their slow movement in relation toone another, solid particles exert strong attrac-tions; yet as slowly as they move, solid particlesdo move—as is the case with all forms of matterat the atomic level. Whereas the particles in a liquid or gas move fast enough to be in rela-tive motion with regard to one another, how-ever, solid particles merely vibrate from a fixedposition.

As noted earlier, the motion and attractionof particles in matter has a direct effect on ther-mal energy, and thus on heat and temperature.The cooler the solid, the slower and weaker thevibrations, and the closer the particles are to one



Propertiesof Matter

another. Thus, most types of matter contractwhen freezing, and their density increases.Absolute zero, or 0K on the Kelvin scale of tem-perature—equal to –459.67°F (–273°C)—is thepoint at which vibration virtually ceases.

Note that the vibration virtually stops, butdoes not totally stop. In fact, as established in the third law of thermodynamics, absolute zero is impossible to achieve: thus, the relative mo-tion of molecules never ceases. The lowest temperature actually achieved, at a Finnishnuclear laboratory in 1993, is 2.8 • 10–10 K, or0.00000000028K—still above absolute zero.

UNUSUAL CHARACTERISTICS

OF SOLID AND LIQUID WATER.

The behavior of water when frozen is interestingand exceptional. Above 39.2°F (4°C) water, likemost substances, expands when heated. In otherwords, the molecules begin moving further apart as expected, because—in this temperaturerange, at least—water behaves like other sub-stances, becoming “less solid” as the temperatureincreases.

Between 32°F (0°C) and 39.2°F (4°C), how-ever, water actually contracts. In this temperaturerange, it is very “cold” (that is, it has relatively lit-tle heat), but it is not frozen. The density of waterreaches its maximum—in other words, watermolecules are as closely packed as they can be—at 39.2°F; below that point, the density starts todecrease again. This is highly unusual: in mostsubstances, the density continues to increase withlowered temperatures, whereas water is actuallymost dense slightly above the freezing point.

Below the freezing point, then, waterexpands, and therefore when water in pipesfreezes, it may increase in volume to the pointwhere it bursts the pipe. This is also the reasonwhy ice floats on water: its weight is less than thatof the water it has displaced, and thus it is buoy-ant. Additionally, the buoyant qualities of iceatop very cold water helps explain the behaviorof lake water in winter; although the top of a lakemay freeze, the entire lake rarely freezes solid—even in the coldest of inhabited regions.

Instead of freezing from the bottom up, as itwould if ice were less buoyant than the water, thelake freezes from the top down—an importantthing to remember when ice-fishing! Further-more, water in general (and ice in particular) is apoor conductor of heat, and thus little of the heatfrom the water below it escapes. Therefore, the

lake does not freeze completely—only a layer atthe top—and this helps preserve animal andplant life in the body of water.

MELTING. When heated, particles beginto vibrate more and more, and therefore movefurther apart. If a solid is heated enough, it losesits rigid structure and becomes a liquid. The tem-perature at which a solid turns into a liquid iscalled the melting point, and melting points aredifferent for different substances. The meltingpoint of a substance, incidentally, is the same asits freezing point: the difference is a matter oforientation—that is, whether the process is oneof a solid melting to become a liquid, or of a liq-uid freezing to become a solid.

The energy required to melt 1 mole of a solidsubstance is called the molar heat of fusion. Itcan be calculated by the formula Q = smδT,where Q is energy, s is specific heat capacity, m ismass, and δT means change in temperature. (Inthe symbolic language often employed by scien-tists, the Greek letter δ, or delta, stands for“change in.”) Specific heat capacity is measuredin units of J/g • °C (joules per gram-degree Cel-sius), and energy in joules or kilojoules (kJ)—that is, 1,000 joules.

In melting, all the thermal energy in a solidis used in breaking up the arrangement of crys-tals, called a lattice. This is why water meltedfrom ice does not feel any warmer than the icedid: the thermal energy has been expended, andthere is none left over for heating the water. Onceall the ice is melted, however, the absorbed ener-gy from the particles—now moving at muchgreater speeds than when the ice was in a solidstate—causes the temperature to rise.

For the most part, solids composed of parti-cles with a higher average atomic mass requiremore energy—and hence higher temperatures—to induce the vibrations necessary for melting.Helium, with an average atomic mass of 4.003amu, melts or freezes at an incredibly low tem-perature: –457.6°F (–272°C), or close to absolutezero. Water, for which, as noted earlier, the aver-age atomic mass is the sum of the masses for itstwo hydrogen atoms and one oxygen atom, hasan average molecular mass of 18.016 amu. Icemelts (or water freezes) at much higher tempera-tures than helium: 32°F (0°C). Copper, with anaverage atomic mass of 63.55 amu, melts atmuch, much higher temperatures than water:1,985°F (1,085°C).



Propertiesof Matter

Liquids

The particles of a liquid, as compared to those ofa solid, have more energy, more motion, and—generally speaking—less attraction to one anoth-er. The attraction, however, is still fairly strong:thus, liquid particles are in close enough pro-ximity that the liquid resists attempts at com-pression.

On the other hand, their arrangement isloose enough that the particles tend to movearound one another rather than simply vibrate inplace the way solid particles do. A liquid is there-fore not definite in shape. Due to the fact that theparticles in a liquid are farther apart than thoseof a solid, liquids tend to be less dense thansolids. The liquid phase of a substance thus tendsto be larger in volume than its equivalent in solidform. Again, however, water is exceptional in thisregard: liquid water actually takes up less spacethan an equal mass of frozen water.

Boiling

When a liquid experiences an increase in temper-ature, its particles take on energy and begin tomove faster and faster. They collide with oneanother, and at some point the particles nearestthe surface of the liquid acquire enough energyto break away from their neighbors. It is at thispoint that the liquid becomes a gas or vapor.

As heating continues, particles throughoutthe liquid begin to gain energy and move faster,but they do not immediately transform into gas.The reason is that the pressure of the liquid,combined with the pressure of the atmosphereabove the liquid, tends to keep particles in place.Those particles below the surface, therefore,remain where they are until they acquire enoughenergy to rise to the surface.

The heated particle moves upward, leavingbehind it a hollow space—a bubble. A bubble isnot an empty space: it contains smaller trappedparticles, but its small mass, relative to that of theliquid it disperses, makes it buoyant. Therefore, abubble floats to the top, releasing its trapped par-ticles as gas or vapor. At that point, the liquid issaid to be boiling.

THE EFFECT OF ATMOSPHER-

IC PRESSURE. The particles thus have toovercome atmospheric pressure as they rise,which means that the boiling point for any liquiddepends in part on the pressure of the surround-

ing air. Normal atmospheric pressure (1 atm) isequal to 14 lb/in2 (1.013 x 105 Pa), and is meas-ured at sea level. The greater the altitude, the lessthe air pressure, because molecules of air—sinceair is a gas, and therefore its particles are fast-moving and non-attractive—respond less toEarth’s gravitational pull. This is why airplanesrequire pressurized cabins to maintain an ade-quate oxygen supply; but even at altitudes muchlower than the flight path of an airplane, differ-ences in air pressure are noticeable.

It is for this reason that cooking instructionsoften vary with altitude. Atop Mt. Everest, Earth’shighest peak at about 29,000 ft (8,839 m) abovesea level, the pressure is approximately one-thirdof normal atmospheric pressure. Water boils at amuch lower temperature on Everest than it doeselsewhere: 158°F (70°C), as opposed to 212°F(100°C) at sea level. Of course, no one lives onthe top of Mt. Everest—but people do live inDenver, Colorado, where the altitude is 5,577 ft(1,700 m) and the boiling point of water is 203°F(95°C).

Given the lower boiling point, one mightassume that food would cook faster in Denverthan in New York, Los Angeles, or in any cityclose to sea level. In fact, the opposite is true:because heated particles escape the water somuch faster at high altitudes, they do not havetime to acquire the energy needed to raise thetemperature of the water. It is for this reason thata recipe may include a statement such as “at alti-tudes above XX feet, add XX minutes to cookingtime.”

If lowered atmospheric pressure means alowered boiling point, what happens in outerspace, where there is no atmospheric pressure?Liquids boil at very, very low temperatures. Thisis one of the reasons why astronauts have to wearpressurized suits: if they did not, their bloodwould boil—even though space itself is incredi-bly cold.

LIQUID TO GAS AND BACK

AGAIN. Note that the process of changing aliquid to a gas is similar to that which occurswhen a solid changes to a liquid: particles gainheat and therefore energy, begin to move faster,break free from one another, and pass a certainthreshold into a new phase of matter. And just asthe freezing and melting point for a given sub-stance are the same temperature—the only dif-ference being one of orientation—the boiling



Propertiesof Matter

point of a liquid transforming into a gas is thesame as the condensation point for a gas turninginto a liquid.

The behavior of water in boiling and con-densation makes possible distillation, one of theprincipal methods for purifying sea water in var-ious parts of the world. First the water is boiled,then it is allowed to cool and condense, thusforming water again. In the process, the waterseparates from the salt, leaving it behind in theform of brine. A similar separation takes placewhen salt water freezes: because salt, like mostcrystalline solids, has a much lower freezingpoint than water, very little of it remains joinedto the water in ice. Instead, the salt takes the formof a briny slush.

Gases

A liquid that is vaporized, or any substance thatexists normally as a gas, is quite different in phys-ical terms from a solid or a liquid. This is illus-trated by the much higher energy component inthe molar heat of vaporization, or the amount ofenergy required to turn 1 mole of a liquid into a gas.

Consider, for instance, what happens towater when it experiences phase changes. Assum-ing that heat is added at a uniform rate, when icereaches its melting point, there is only a relative-ly small period of time when the H2O is com-posed of both ice and liquid. But when the liquidreaches its boiling point, the water is presentboth as a liquid and a vapor for a much longerperiod of time. In fact, it takes almost seven timesas much energy to turn liquid water into puresteam than it does to turn ice into purely liquidwater. Thus, the molar heat of fusion for water is6.02 kJ/mol, while the molar heat of vaporizationis 40.6 kJ/mol.

Although liquid particles exert a moderateattraction toward one another, particles in a gas(particularly a substance that normally exists as agas at ordinary temperatures on Earth) exert lit-tle to no attraction. They are thus free to move,and to move quickly. The overall shape andarrangement of gas is therefore random andindefinite—and, more importantly, the motionof gas particles provides much greater kineticenergy than is present in any other major form ofmatter on Earth.

The constant, fast, and random motion ofgas particles means that they are regularly collid-

ing and thereby transferring kinetic energy backand forth without any net loss of energy. Thesecollisions also have the overall effect of produc-ing uniform pressure in a gas. At the same time,the characteristics and behavior of gas particlesindicate that they tend not to remain in an opencontainer. Therefore, in order to have any pres-sure on a gas—other than normal atmosphericpressure—it is necessary to keep it in a closedcontainer.

The Phase Diagram

The vaporization of water is an example of achange of phase—the transition from one phaseof matter to another. The properties of any sub-stance, and the points at which it changes phase,are plotted on what is known as a phase diagram.The phase diagram typically shows temperaturealong the x-axis, and pressure along the y-axis.

For simple substances, such as water andcarbon dioxide (CO2), the solid form of the sub-stance appears at a relatively low temperatureand at pressures anywhere from zero upward.The line between solids and liquids, indicatingthe temperature at which a solid becomes a liq-uid at any pressure above a certain level, is calledthe fusion curve. Though it appears to be a moreor less vertical line, it is indeed curved, indicatingthat at high pressures, a solid well below the nor-mal freezing point may be melted to create a liquid.

Liquids occupy the area of the phase dia-gram corresponding to relatively high tempera-tures and high pressures. Gases or vapors, on theother hand, can exist at very low temperatures,but only if the pressure is also low. Above themelting point for the substance, gases exist athigher pressures and higher temperatures. Thus,the line between liquids and gases often looksalmost like a 45° angle. But it is not a straightline, as its name, the vaporization curve, implies.The curve of vaporization demonstrates that atrelatively high temperatures and high pressures, asubstance is more likely to be a gas than a liquid.

THE CRITICAL POINT. There areseveral other interesting phenomena mapped ona phase diagram. One is the critical point, foundat a place of very high temperature and pressurealong the vaporization curve. At the criticalpoint, high temperatures prevent a liquid fromremaining a liquid, no matter how high the pressure.



Propertiesof Matter

At the same time, the pressure causes gasbeyond that point to become increasingly moredense, but due to the high temperatures, it doesnot condense into a liquid. Beyond the criticalpoint, the substance cannot exist in anythingother than the gaseous state. The temperaturecomponent of the critical point for water is705.2°F (374°C)—at 218 atm, or 218 times ordi-nary atmospheric pressure. For helium, however,critical temperature is just a few degrees aboveabsolute zero. This is, in part, why helium israrely seen in forms other than a gas.

THE SUBLIMATION CURVE.

Another interesting phenomenon is the sublima-tion curve, or the line between solid and gas. Atcertain very low temperatures and pressures, asubstance may experience sublimation, meaningthat a gas turns into a solid, or a solid into a gas,without passing through a liquid stage.

A well-known example of sublimationoccurs when “dry ice,” made of carbon dioxide,vaporizes at temperatures above (–78.5°C). Car-bon dioxide is exceptional, however, in that itexperiences sublimation at relatively high pres-sures that occur in everyday life: for most sub-stances, the sublimation point transpires at sucha low pressure point that it is seldom witnessedoutside of a laboratory.

THE TRIPLE POINT. The phenom-enon known as the triple point shows how anordinary substance such as water or carbon diox-ide can actually be a liquid, solid, and vapor—allat once. Most people associate water as a gas orvapor (that is, steam) with very high tempera-tures. Yet, at a level far below normal atmospher-ic pressure, water can be a vapor at temperaturesas low as –4°F (–20 °C). (All of the pressure val-ues in the discussion of water at or near the triplepoint are far below atmospheric norms: the pres-sure at which water turns into a vapor at –4°F, forinstance, is about 0.001 atm.)

Just as water can exist as a vapor at low tem-peratures and low pressures, it is also possible forwater at temperatures below freezing to remainliquid. Under enough pressure, ice melts and isthereby transformed from a solid to a liquid, attemperatures below its normal freezing point.On the other hand, if the pressure of ice fallsbelow a very low threshold, it will sublimate.

The phase diagram of water shows a linebetween the solid and liquid states that is almost,but not quite, exactly perpendicular to the x-axis.But in fact, it is a true fusion curve: it slopesslightly upward to the left, indicating that solidice turns into water with an increase of pressure.Below a certain level of pressure is the vaporiza-


pres

sure

(y-a

xis)

temperature(x-axis)

triple point

(solid)(liquid)

(vapor)

A PHASE DIAGRAM FOR WATER.


Propertiesof Matter

tion curve, and where the fusion curve intersectsthe vaporization curve, there is a place called thetriple point. Just below freezing, in conditionsequivalent to about 0.007 atm, water is a solid,liquid, and vapor all at once.

Other States of Matter

PLASMA. Principal among states ofmatter other than solid, liquid, and gas is plasma,which is similar to gas. (The term “plasma,” whenreferring to the state of matter, has nothing to dowith the word as it is often used, in reference toblood plasma.) As with gas, plasma particles col-lide at high speeds—but in plasma the speeds areeven greater, and the kinetic energy levels evenhigher.

The speed and energy of these collisions isdirectly related to the underlying property thatdistinguishes plasma from gas. So violent are thecollisions between plasma particles that electronsare knocked away from their atoms. As a result,plasma does not have the atomic structure typi-cal of a gas; rather, it is composed of positive ionsand electrons. Plasma particles are thus electri-cally charged, and therefore greatly influenced byelectric and magnetic fields.

Formed at very high temperatures, plasma isfound in stars. The reaction between plasma andatomic particles in the upper atmosphere isresponsible for the aurora borealis, or “northernlights.” Though found on Earth only in verysmall quantities, plasma—ubiquitous in otherparts of the universe—may be the most plentifulof all the states of matter.

QUASI-STATES. Among the quasi-states of matter discussed by scientists are severalterms describing the structure in which particlesare joined, rather than the attraction and relativemovement of those particles. Thus “crystalline,”“amorphous,” and “glassy” are all terms todescribe what may be individual states of matter;so too is “colloidal.”

A colloid is a structure intermediate in sizebetween a molecule and a visible particle, and ithas a tendency to be dispersed in another medi-um—the way smoke, for instance, is dispersed inair. Brownian motion describes the behavior ofmost colloidal particles. When one sees dustfloating in a ray of sunshine through a window,the light reflects off colloids in the dust, whichare driven back and forth by motion in the airotherwise imperceptible to the human senses.

DARK MATTER. The number of statesor phases of matter is clearly not fixed, and it isquite possible that more will be discovered inouter space, if not on Earth. One intriguing can-didate is called dark matter, so described becauseit neither reflects nor emits light, and is thereforeinvisible. In fact, luminous or visible matter mayvery well make up only a small fraction of themass in the universe, with the rest being taken upby dark matter.

If dark matter is invisible, how doastronomers and physicists know it exists? Byanalyzing the gravitational force exerted on visi-ble objects in such cases where there appears tobe no visible object to account for that force. Anexample is the center of our galaxy, the MilkyWay. It appears to be nothing more than a dark“halo,” but in order to cause the entire galaxy torevolve around it—in the same way that planetsrevolve around the Sun, though on a vastly larg-er scale—it must contain a staggering quantity ofinvisible mass.

The Bose-Einstein Condensate

Physicists at the Joint Institute of LaboratoryAstrophysics in Boulder, Colorado, in 1995revealed a highly interesting aspect of atomicbehavior at temperatures approaching absolutezero. Some 70 years before, Einstein had predict-ed that, at extremely low temperatures, atomswould fuse to form one large “superatom.” Thishypothesized structure was dubbed the Bose-Einstein Condensate (BEC) after Einstein andSatyendranath Bose (1894-1974), an Indianphysicist whose statistical methods contributedto the development of quantum theory.

Cooling about 2,000 atoms of the elementrubidium to a temperature just 170 billionths ofa degree Celsius above absolute zero, the physi-cists succeeded in creating an atom 100 microm-eters across—still incredibly small, but vast incomparison to an ordinary atom. The super-atom, which lasted for about 15 seconds, cooleddown all the way to just 20 billionths of a degreeabove absolute zero. The Colorado physicistswon the Nobel Prize in physics in 1997 for theirwork.

In 1999, researchers in a lab at Harvard University also created a superatom of BEC,and used it to slow light to just 38 MPH (61.2km/h)—about 0.02% of its ordinary speed.



Propertiesof Matter

Dubbed a “new” form of matter, the BEC maylead to a greater understanding of quantummechanics, and may aid in the design of smaller,more powerful computer chips.

Some Unusual Phase Transitions

At places throughout this essay, references havebeen made variously to “phases” and “states” ofmatter. This is not intended to confuse, butrather to emphasize a particular point. Solids,liquids, and gases are referred to as “phases”because many (though far from all) substanceson Earth regularly move from one phase toanother.

There is absolutely nothing incorrect inreferring to “states of matter.” But “phases ofmatter” is used in the present context as a meansof emphasizing the fact that substances, at theappropriate temperature and pressure, can besolid, liquid, or gas. The phases of matter, in fact,can be likened to the phases of a person’s life:infancy, babyhood, childhood, adolescence,adulthood, old age. The transition between thesestages is indefinite, yet it is easy enough to saywhen a person is at a certain stage.

LIQUID CRYSTALS. A liquid crystalis a substance that, over a specific range of tem-perature, displays properties both of a liquid anda solid. Below this temperature range, it isunquestionably a solid, and above this range it isjust as certainly a liquid. In between, however,liquid crystals exhibit a strange solid-liquidbehavior: like a liquid, their particles flow, butlike a solid, their molecules maintain specificcrystalline arrangements.

The cholesteric class of liquid crystals is sonamed because the spiral patterns of lightthrough the crystal are similar to those whichappear in cholesterols. Depending on the physi-cal properties of a cholesteric liquid crystal, onlycertain colors may be reflected. The response ofliquid crystals to light makes them useful in liq-uid crystal displays (LCDs) found on laptopcomputer screens, camcorder views, and in otherapplications.

LIQUEFACTION OF GASES.

One interesting and useful application of phasechange is the liquefaction of gases, or the changeof gas into liquid by the reduction in its molecu-

lar energy levels. Liquefied natural gas (LNG)and liquefied petroleum gas (LPG), the latter amixture of by-products obtained from petrole-um and natural gas, are among the examples ofliquefied gas in daily use. In both cases, the vol-ume of the liquefied gas is far less than it wouldbe if the gas were in a vaporized state, thusenabling ease and economy of transport.

Liquefied gases are used as heating fuel formotor homes, boats, and homes or cabins inremote areas. Other applications of liquefiedgases include liquefied oxygen and hydrogen inrocket engines; liquefied oxygen and petroleumused in welding; and a combination of liquefiedoxygen and nitrogen used in aqualung devices.The properties of liquefied gases figure heavily inthe science of producing and studying low-tem-perature environments. In addition, liquefiedhelium is used in studying the behavior of matterat temperatures close to absolute zero.

COAL GASIFICATION. Coal gasifi-cation, as one might discern from the name, isthe conversion of coal to gas. Developed beforeWorld War II, it fell out of favor after the war, dueto the lower cost of oil and natural gas. However,increasingly stringent environmental regulationsimposed by the federal government on industryduring the 1970s, combined with a growing con-cern for the environment on the part of the pop-ulace as a whole, led to a resurgence of interest incoal gasification.

Though widely used as a fuel in powerplants, coal, when burned by ordinary means,generates enormous air pollution. Coal gasifica-tion, on the other hand, makes it possible to burn“clean” coal. Gasification involves a number ofchemical reactions, some exothermic or heat-releasing, and some endothermic or heat-absorb-ing. At one point, carbon monoxide is released inan exothermic reaction, then mixed with hydro-gen released from the coal to create a secondexothermic reaction. The energy discharged inthese first two reactions is used to initiate a third,endothermic, reaction.

The finished product of coal gasification is amixture containing carbon monoxide, methane,hydrogen, and other substances, and this—ratherthan ordinary coal—is burned as a fuel. Thecomposition of the gases varies according to theprocess used. Products range from coal synthesisgas and medium-Btu gas (both composed of car-



Propertiesof Matter


ABSOLUTE ZERO: The temperature,

defined as 0K on the Kelvin scale, at which

the motion of molecules in a solid virtual-

ly ceases. Absolute zero is equal to

–459.67°F (–273.15°C).

ATOM: The smallest particle of an ele-

ment that retains the chemical and physical

properties of the element. An atom can

exist either alone or in combination with

other atoms in a molecule. Atoms are made

up of protons, neutrons, and electrons. An

atom that loses or gains one or more elec-

trons, and thus has a net charge, is an ion.

Atoms that have the same number of pro-

tons—that is, are of the same element—

but differ in number of neutrons, are

known as isotopes.


(abbreviated amu), equal to 1.66 • 10–24 g,


ATOMIC NUMBER: The number of

protons in the nucleus of an atom. Since

this number is different for each element,

elements are listed on the periodic table in

order of atomic number.

ATOMIC SOLID: A form of crystalline

solid in which atoms of one element bond

to one another. Examples include dia-

monds (made of pure carbon), silicon, and

all metals.




a large sample. If a substance is a com-

pound, the average atomic mass of all

atoms in a molecule of that substance must

be added together to yield the average

molecular mass of that substance.



gadro (1776-1856), equal to 6.022137 x

1023. Avogadro’s number indicates the

number of atoms, molecules, or other ele-

mentary particles in a mole.

CHANGE OF PHASE: The transition

from one phase of matter to another.

CRITICAL POINT: A coordinate, plot-

ted on a phase diagram, above which a sub-

stance cannot exist in anything other than

the gaseous state. Located at a position of

very high temperature and pressure, the

critical point marks the termination of the

vaporization curve.

COMPOUND: A substance made up of

atoms of more than one element. These

atoms are usually, but not always, joined in

molecules.

CRYSTALLINE SOLID: A type of

solid in which the constituent parts have a

simple and definite geometric arrange-

ment that is repeated in all directions.

Types of crystalline solids include atomic

solids, ionic solids, and molecular solids.

ELEMENT: A substance made up of

only one kind of atom.

ELECTRON: A negatively charged par-

ticle in an atom. Electrons, which spin

around the protons and neutrons that

make up the atom’s nucleus, constitute a

very small portion of the atom’s mass. In

most atoms, the number of electrons and

protons is the same, thus canceling out one

another. When an atom loses one or more

electrons, however—thus becoming an

ion—it acquires a net electric charge.

KEY TERMS


Propertiesof Matter


ENERGY: The ability to accomplish

work—that is, the exertion of force over a

given distance to displace or move an

object.

FUSION CURVE: The boundary be-

tween solid and liquid for any given sub-

stance as plotted on a phase diagram.

GAS: A phase of matter in which mole-

cules move at high speeds, and therefore

exert little or no attraction toward one

another.

ION: An atom or atoms that has lost or

gained one or more electrons, and thus has

a net electric charge.

IONIC SOLID: A form of crystalline

solid that contains ions. When mixed with

a solvent such as water, ions from table

salt—an example of an ionic solid—move

freely throughout the solution, making it

possible to conduct an electric current.

ISOTOPES: Atoms that have an equal

number of protons, and hence are of the

same element, but differ in their number of

neutrons.


an object possesses by virtue of its

movement.

LIQUID: A phase of matter in which

molecules move at moderate speeds, and

therefore exert moderate attractions

toward one another.

MATTER: Physical substance that occu-

pies space, has mass, is composed of atoms

(or in the case of subatomic particles, is

part of an atom), and is convertible to

energy.

MOLAR HEAT OF FUSION: The

amount of energy required to melt 1 mole

of a solid substance.

MOLAR HEAT OF VAPORIZATION:

The amount of energy required to turn 1

mole of a liquid into a gas. Because gases

possess much more energy than liquids or

solids, molar heat of vaporization is usual-

ly much higher than molar heat of fusion.





atomic mass of that substance: thus car-










MOLECULAR SOLID: A form of crys-

talline solid in which the molecules have a

neutral electric charge, meaning that there

are no ions present as there are in an ionic

solid. Table sugar (sucrose) is an example

of a molecular solid. When mixed with a

solvent such as water, the resulting solution

does not conduct electricity.

MOLECULE: A group of atoms, usual-

ly of more than one element, joined in a

structure.

NEUTRON: A subatomic particle that

has no electric charge. Neutrons are found

at the nucleus of an atom, alongside

protons.



Propertiesof Matter


bon monoxide and hydrogen, though combinedin different forms) to substitute natural gas,which consists primarily of methane.

Not only does coal gasification produce aclean-burning product, but it does so without thehigh costs associated with flue-gas desulfuriza-tion systems. The latter, often called “scrubbers,”were originally recommended by the federal gov-ernment to industry, but companies discoveredthat coal gasification could produce the sameresults for much less money. In addition, thewaste products from coal gasification can be used

for other purposes. At the Cool Water Integrated

Gasification Combined Cycle Plant, established

in Barstow, California, in 1984, sulfur obtained

from the reduction of sulfur dioxide is sold off

for about $100 a ton.

The Chemical Dimension toChanges of Phase

Throughout much of this essay, we have dis-

cussed changes of phase primarily in physical

terms; yet clearly these changes play a significant

role in chemistry. Furthermore, coal gasification

NUCLEUS: The center of an atom, a

region where protons and neutrons are

located, and around which electrons spin.

PHASE DIAGRAM: A chart, plotted

for any particular substance, identifying

the particular phase of matter for that sub-

stance at a given temperature and pressure

level. A phase diagram usually shows tem-

perature along the x-axis, and pressure

along the y-axis.

PHASES OF MATTER: The various

forms of material substance (matter),

which are defined primarily in terms of the

behavior exhibited by their atomic or

molecular structures. On Earth, three prin-

cipal phases of matter exist, namely solid,

liquid, and gas. Other forms of matter

include plasma.

PLASMA: One of the phases of matter,

closely related to gas. Plasma is found pri-

marily in stars. Containing neither atoms

nor molecules, plasma is made up of elec-

trons and positive ions.

PROTON: A positively charged particle

in an atom. Protons and neutrons, which

together form the nucleus around which

electrons spin, have approximately the

same mass—a mass that is many times

greater than that of an electron. The num-

ber of protons in the nucleus of an atom is

the atomic number of an element.

SOLID: A phase of matter in which

molecules move slowly relative to one

another, and therefore exert strong attrac-

tions toward one another.

SUBLIMATION CURVE: The bound-

ary between solid and gas for any given

substance, as plotted on a phase diagram.








THERMAL ENERGY: “Heat” energy, a

form of kinetic energy produced by the rel-

ative motion of atoms or molecules. The

greater the movement of these particles,

the greater the thermal energy.

VAPORIZATION CURVE: The bound-

ary between liquid and gas for any given

substance as plotted on a phase diagram.



Propertiesof Matter

serves to illustrate the impact chemical processescan have on changes of state.

Much earlier, figures were given for the melt-ing points of copper, water, and helium, andthese were compared with the average atomicmass of each. Those figures, again, are:

Average Atomic Mass and Melting Points of aSample Gas, Liquid, and Solid

• Helium: 4.003 amu; –457.6°F (–210°C)

• Water: 40.304 amu 32°F (0°C)

• Copper: 63.55 amu; 1,985°F (1,085°C).

Something seems a bit strange about thosecomparisons: specifically, the differences in melt-ing point appear to be much more dramatic thanthe differences in average atomic mass. Clearly,another factor is at work—a factor that relates tothe difference in the attractions between mole-cules in each. Although the differences betweensolids, liquids, and gases are generally physical,the one described here—a difference betweensubstances—is clearly chemical in nature.

To discuss this in the detail it deserves wouldrequire a lengthy digression on the chemicaldimensions of intermolecular attraction.Nonetheless, it is possible here to offer at least acursory answer to the question raised by thesestriking differences in response to temperature.

Dipoles, Electron Seas, and London Dispersion

Water molecules are polar, meaning that one areaof a water molecule is positively charged, whileanother area has a negative charge. Thus the pos-itive side of one molecule is drawn to the nega-tive side of another, and vice versa, which giveswater a much stronger intermolecular bondthan, for instance, oil, in which the positive andnegative charges are evenly distributed through-out the molecule.

Yet the intermolecular attraction betweenthe dipoles (as they are called) in water is notnearly as strong as the bond that holds together ametal. Particles in copper or other metals “float”in a tightly packed “sea” of highly mobile elec-trons, which provide a bond that is powerful, yetlacking in a firm directional orientation. Thusmetals are both strong and highly malleable (thatis, they can be hammered very flat withoutbreaking.)

Water, of course, appears most often as a liq-uid, and copper as a solid, precisely because waterhas a very high boiling point (the point at whichit becomes a vapor) and copper has a very highmelting point. But consider helium, which hasthe lowest freezing point of any element: justabove absolute zero. Even then, a pressure equalto 25 times that of normal atmospheric pressureis required to push it past the freezing point.

Helium and other Group 8 or Group 18 ele-ments, as well as non-polar molecules such asoils, are bonded by what is called London disper-sion forces. The latter, as its name suggests, tendsto keep molecules dispersed, and induces instan-taneous dipoles when most of the electrons hap-pen to be on one side of an atom. Of course, thishappens only for an infinitesimal fraction oftime, but it serves to create a weak attraction.Only at very low temperatures do London dis-persion forces become strong enough to result inthe formation of a solid.

WHERE TO LEARN MORE

Biel, Timothy L. Atom: Building Blocks of Matter. SanDiego, CA: Lucent Books, 1990.

Feynman, Richard. Six Easy Pieces: Essentials of PhysicsExplained by Its Most Brilliant Teacher. New intro-duction by Paul Davies. Cambridge, MA: PerseusBooks, 1995.

“High School Chemistry Table of Contents—Solids andLiquids” Homeworkhelp.com (Web site). <http://www.homeworkhelp.com/homeworkhelp/freemember/text/chem/hig h/topic09.htm> (April 10, 2001).

“Matter: Solids, Liquids, Gases.” Studyweb (Web site).<http://www.studyweb.com/links/4880.html> (April10, 2001).

“The Molecular Circus” (Web site). <http://www.cpo.com/Weblabs/circus.htm> (April 10, 2001).

Paul, Richard. A Handbook to the Universe: Explorationsof Matter, Energy, Space, and Time for Beginning Scientific Thinkers. Chicago: Chicago Review Press,1993.

“Phases of Matter” (Web site). <http://pc65.frontier.osrhe.edu/hs/science/pphase.htm> (April 10, 2001).

Royston, Angela. Solids, Liquids, and Gasses. Chicago:Heinemann Library, 2001.

Wheeler, Jill C. The Stuff Life’s Made Of: A Book AboutMatter. Minneapolis, MN: Abdo & Daughters Pub-lishing, 1996.





GASESGases

CONCEPTThe number of elements that appear ordinarilyin the form of a gas is relatively small: oxygen,hydrogen, fluorine, and chlorine in the halogen“family”; and a handful of others, most notablythe noble gases in Group 8 of the periodic table.Yet many substances can exist in the form of agas, depending on the relative attraction andmotion of molecules in that substance. A simpleexample, of course, is water, or H2O, which,though it appears as a liquid at room tempera-ture, begins to vaporize and turn into steam at212°F (100°C). In general, gases respond moredramatically to changes in pressure and temper-ature than do most other types of matter, andthis allows scientists to predict gas behaviorsunder certain conditions. These predictions canexplain mundane occurrences, such as the factthat an open can of soda will soon lose its fizz,but they also apply to more dramatic, life-and-death situations.

HOW IT WORKS

Molecular Motion and Phas-es of Matter

On Earth, three principal phases or states of mat-ter exist: solid, liquid, and gas. The differencesbetween these three are, on the surface at least,easily perceivable. Clearly water is a liquid, just asice is a solid and steam a vapor or gas. Yet theways in which various substances convertbetween phases are often complex, as are theinterrelations between these phases. Ultimately,understanding of the phases depends on an

awareness of what takes place at the molecularlevel.

All molecules are in motion, and the rate ofthat motion determines the attraction betweenthem. The movement of molecules generateskinetic energy, or the energy of movement, whichis manifested as thermal energy. In everyday lan-guage, thermal energy is what people mean whenthey say “heat”; but in scientific terms, heat has adifferent definition.

The force that attracts atoms to atoms, ormolecules to molecules, is not the same as gravi-tational force, which holds the Moon in orbitaround Earth, Earth in orbit around the Sun, andso on. By contrast, the force of interatomic andintermolecular attraction is electromagnetic. Justas the north pole of a magnet is attracted to thesouth pole of another magnet and repelled bythat other magnet’s north pole, so positive elec-tric charges are attracted to negative charges, andnegatives to positives. (In fact, electricity andmagnetism are both manifestations of an electro-magnetic interaction.)

The electromagnetic attractions betweenmolecules are much more complex than thisexplanation makes it seem, and they play a high-ly significant role in chemical bonding. In simpleterms, however, one can say that the greater therate of motion for the molecules in relation toone another, the less the attraction between mol-ecules. In addition, the kinetic energy, and hencethe thermal energy, is greater in a substancewhose molecules are relatively free to move.

When the molecules in a material moveslowly in relation to one another, they exert astrong attraction, and the material is called asolid. Molecules of liquid, by contrast, move at


Gases


moderate speeds and exert a moderate attrac-tion. A material substance whose moleculesmove at high speeds, and therefore exert little orno attraction, is known as a gas.

Comparison of Gases toOther Phases of Matter

WATER AND AIR COMPARED.

Gases respond to changes in pressure and tem-perature in a manner remarkably different fromthat of solids or liquids. Consider the behavior ofliquid water as compared with air—a combina-tion of oxygen (O2), nitrogen (N2), and othergases—in response to experiments involvingchanges in pressure and temperature.

In the first experiment, both samples aresubjected to an increase in pressure from 1 atm(that is, normal atmospheric pressure at sealevel) to 2 atm. In the second, both experience anincrease in temperature from 32°F (0°C) to 212°F(100°C). The differences in the responses ofwater and air are striking.

A sample of water that experiences anincrease in pressure from 1 to 2 atm will decreasein volume by less than 0.01%, while a tempera-ture increase from the freezing point to the boil-ing point will result in only a 2% increase in vol-ume. For air, however, an equivalent pressure

increase will decrease the volume by a whopping

50%, and an equivalent temperature increase

results in a volume increase of 37%.

Air and other gases, by definition, have a

boiling point below room temperature. If they

did not boil and thus become gas well below

ordinary temperatures, they would not be

described as substances that are in the gaseous

state in most circumstances. The boiling point of

water, of course, is higher than room tempera-

ture, and that of solids is much higher.

THE ARRANGEMENT OF PAR-

TICLES. Solids possess a definite volume and

a definite shape, and are relatively noncompress-

ible: for instance, if one applies extreme pressure

to a steel plate, it will bend, but not much. Liq-

uids have a definite volume, but no definite

shape, and tend to be noncompressible. Gases, on

the other hand, possess no definite volume or

shape, and are highly compressible.

At the molecular level, particles of solids

tend to be definite in their arrangement and close

in proximity—indeed, part of what makes a solid

“solid,” in the everyday meaning of that term, is

the fact that its constituent parts are basically

immovable. Liquid molecules, too, are close in

proximity, though random in arrangement. Gas

A COMPARISON OF GAS, LIQUID, AND SOLID STATES.


Gases units of pressure represent some ratio of force tosurface area.

UNITS OF PRESSURE. The prin-cipal SI unit of pressure is called a pascal (Pa), or1 N/m2. It is named for French mathematicianand physicist Blaise Pascal (1623-1662), who iscredited with Pascal’s principle. The latter holdsthat the external pressure applied on a fluid—which, in the physical sciences, can mean either agas or a liquid—is transmitted uniformlythroughout the entire body of that fluid.

A newton (N), the SI unit of force, is equal tothe force required to accelerate 1 kg of mass at arate of 1 m/sec2. Thus a Pascal is the pressure of 1newton over a surface area of 1 m2. In the Englishor British system, pressure is measured in termsof pounds per square inch, abbreviated aslbs./in2. This is equal to 6.89 • 103 Pa, or 6,890 Pa.

Another important measure of pressure isthe atmosphere (atm), which is the average pres-sure exerted by air at sea level. In English units,this is equal to 14.7 lb/in2, and in SI units, to1.013 • 105 Pa.

There are two other specialized units ofpressure measurement in the SI system: the bar,equal to 105 Pa, and the torr, equal to 133 Pa.Meteorologists, scientists who study weather pat-terns, use the millibar (mb), which, as its nameimplies, is equal to 0.001 bars. At sea level, atmos-pheric pressure is approximately 1,013 mb.

The torr, also known as the millimeter ofmercury (mm Hg), is the amount of pressurerequired to raise a column of mercury (chemicalsymbol Hg) by 1 mm. It is named for Italianphysicist Evangelista Torricelli (1608-1647), whoinvented the barometer, an instrument for meas-uring atmospheric pressure.

THE BAROMETER. The barometerconstructed by Torricelli in 1643 consisted of along glass tube filled with mercury. The tube wasopen at one end, and turned upside down into adish containing more mercury: the open end wassubmerged in mercury, while the closed end atthe top constituted a vacuum—that is, an areadevoid of matter, including air.

The pressure of the surrounding air pusheddown on the surface of the mercury in the bowl,while the vacuum at the top of the tube providedan area of virtually no pressure into which themercury could rise. Thus the height to which the


molecules are random in arrangement, but tendto be more widely spaced than liquid molecules.

Pressure

There are a number of statements, collectivelyknown as the “gas laws,” that describe and predictthe behavior of gases in response to changes intemperature, pressure, and volume. Temperatureand volume are discussed elsewhere in this book.However, the subject of pressure requires someattention before we can continue with a discus-sion of the gas laws.

When a force is applied perpendicular to asurface area, it exerts pressure on that surface.Hence the formula for pressure is p = F/A, wherep is pressure, F force, and A the area over whichthe force is applied. The greater the force, and thesmaller the area of application, the greater thepressure; conversely, an increase in area—evenwithout a reduction in force—reduces the overallpressure.

Pressure is measured by a number of units inthe English and SI systems. Because p = F/A, all

THE DIVERS PICTURED HERE HAVE ASCENDED FROM A

SUNKEN SHIP AND HAVE STOPPED AT THE 10-FT (3-METER) DECOMPRESSION LEVEL TO AVOID GETTING

DECOMPRESSION SICKNESS, BETTER KNOWN AS THE

“BENDS.” (Jonathan Blair/Corbis. Reproduced by permission)


Gases


mercury rose in the glass tube represented

normal air pressure (that is, 1 atm.) Torricelli

discovered that at standard atmospheric pres-

sure, the column of mercury rose to 760 mm

(29.92 in).

The value of 1 atm was thus established as

equal to the pressure exerted on a column of

mercury 760 mm high at a temperature of 0°C

(32°F). In time, Torricelli’s invention became a

fixture both of scientific laboratories and of

households. Since changes in atmospheric pres-

sure have an effect on weather patterns, many

home indoor-outdoor thermometers today also

include a barometer.


Introduction to the Gas Laws

English chemist Robert Boyle (1627-1691), whomade a number of important contributions tochemistry—including his definition and identifi-cation of elements—seems to have been influ-enced by Torricelli. If so, this is an interestingexample of ideas passing from one great thinkerto another: Torricelli, a student of Galileo Galilei(1564-1642), was no doubt influenced byGalileo’s thermoscope.

Like Torricelli, Boyle conducted tests involv-ing the introduction of mercury to a tube closed

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THE MAJOR COMPONENTS OF AN INTERNAL COMBUSTION ENGINE (TOP), AND THE FOUR STROKES OF ITS COMBUS-TION SEQUENCE (BOTTOM).


Gases at the other end. The tube Boyle used was shapedlike the letter J, and it was so long that he had touse the multi-story foyer of his house as a labo-ratory. At the tip of the curved bottom was anarea of trapped gas, and into the top of the tube,Boyle introduced increasing quantities of mercu-ry. He found that the greater the volume of mer-cury, the greater the pressure on the gas, and theless the volume of gas at the end of the tube. As aresult, he formulated the gas law associated withhis name.

The gas laws are not a set of government reg-ulations concerning use of heating fuel; rather,they are a series of statements concerning thebehavior of gases in response to changes in tem-perature, pressure, and volume. These werederived, beginning with Boyle’s law, during theseventeenth, eighteenth, and nineteenth cen-turies by scientists whose work is commemorat-ed through the association of their names withthe laws they discovered. In addition to Boyle,these men include fellow English chemists JohnDalton (1766-1844) and William Henry (1774-1836); French physicists and chemists J. A. C.Charles (1746-1823) and Joseph Gay-Lussac(1778-1850); and Italian physicist Amedeo Avo-gadro (1776-1856).

There is a close relationship between Boyle’s,Charles’s, and Gay-Lussac’s laws. All of these treatone of three parameters—temperature, pressure,or volume—as fixed quantities in order toexplain the relationship between the other twovariables. Avogadro’s law treats two of theparameters as fixed, thereby establishing a rela-tionship between volume and the number ofmolecules in a gas. The ideal gas law sums upthese four laws, and the kinetic theory of gasesconstitutes an attempt to predict the behavior ofgases based on these laws. Finally, Dalton’s andHenry’s laws both relate to partial pressure ofgases.

Boyle’s, Charles’s, and Gay-Lussac’s Laws

BOYLE’S AND CHARLES’S

LAWS. Boyle’s law holds that in isothermalconditions (that is, a situation in which tempera-ture is kept constant), an inverse relationshipexists between the volume and pressure of a gas.(An inverse relationship is a situation involvingtwo variables, in which one of the two increasesin direct proportion to the decrease in the other.)

In this case, the greater the pressure, the less thevolume and vice versa. Therefore, the product ofthe volume multiplied by the pressure remainsconstant in all circumstances.

Charles’s law also yields a constant, but inthis case the temperature and volume are allowedto vary under isobarometric conditions—that is,a situation in which the pressure remains thesame. As gas heats up, its volume increases, andwhen it cools down, its volume reduces accord-ingly. Hence, Charles established that the ratio oftemperature to volume is constant.

ABSOLUTE TEMPERATURE. Inabout 1787, Charles made an interesting discov-ery: that at 0°C (32°F), the volume of gas at con-stant pressure drops by 1/273 for every Celsiusdegree drop in temperature. This seemed to sug-gest that the gas would simply disappear if cooledto -273°C (-459.4°F), which, of course, made nosense. In any case, the gas would most likelybecome first a liquid, and then a solid, longbefore it reached that temperature.

The man who solved the quandary raised byCharles’s discovery was born a year after Charlesdied. He was William Thomson, Lord Kelvin(1824-1907); in 1848, he put forward the sug-gestion that it was molecular translational energy—the energy generated by molecules inmotion—and not volume, that would becomezero at -273°C. He went on to establish whatcame to be known as the Kelvin scale of absolutetemperature.

Sometimes known as the absolute tempera-ture scale, the Kelvin scale is based not on thefreezing point of water, but on absolute zero—the temperature at which molecular motioncomes to a virtual stop. This is -273.15°C (-459.67°F). In the Kelvin scale, which uses nei-ther the term nor the symbol for “degree,”absolute zero is designated as 0K.

Scientists prefer the Kelvin scale to the Cel-sius, and certainly to the Fahrenheit, scales. If theKelvin temperature of an object is doubled, itsaverage molecular translational energy has dou-bled as well. The same cannot be said if the tem-perature were doubled from, say, 10°C to 20°C, orfrom 40°F to 80°F, since neither the Celsius northe Fahrenheit scale is based on absolute zero.

GAY-LUSSAC’S LAW. From Boyle’sand Charles’s law, a pattern should be emerging:both treat one parameter (temperature inBoyle’s, pressure in Charles’s) as unvarying, while



Gasestwo other factors are treated as variables. Both, inturn, yield relationships between the two vari-ables: in Boyle’s law, pressure and volume areinversely related, whereas in Charles’s law, tem-perature and volume are directly related.

In Gay-Lussac’s law, a third parameter, vol-ume, is treated as a constant, and the result is aconstant ratio between the variables of pressureand temperature. According to Gay-Lussac’s law,the pressure of a gas is directly related to itsabsolute temperature.

Avogadro’s Law

Gay-Lussac also discovered that the ratio inwhich gases combine to form compounds can beexpressed in whole numbers: for instance, wateris composed of one part oxygen and two partshydrogen. In the language of modern chemistry,this is expressed as a relationship between mole-cules and atoms: one molecule of water containsone oxygen atom and two hydrogen atoms.

In the early nineteenth century, however, sci-entists had yet to recognize a meaningful distinc-tion between atoms and molecules, and Avo-gadro was the first to achieve an understandingof the difference. Intrigued by the whole-numberrelationship discovered by Gay-Lussac, Avogadroreasoned that one liter of any gas must containthe same number of particles as a liter of anoth-er gas. He further maintained that gas consists ofparticles—which he called molecules—that inturn consist of one or more smaller particles.

In order to discuss the behavior of mole-cules, Avogadro suggested the use of a largequantity as a basic unit, since molecules them-selves are very small. Avogadro himself did notcalculate the number of molecules that should beused for these comparisons, but when that num-ber was later calculated, it received the name“Avogadro’s number” in honor of the man whointroduced the idea of the molecule. Equal to6.022137 • 1023, Avogadro’s number designatesthe quantity of atoms or molecules (dependingon whether the substance in question is an ele-ment or a compound) in a mole.

Today the mole (abbreviated mol), the SIunit for “amount of substance,” is defined pre-cisely as the number of carbon atoms in 12.01 gof carbon. The term “mole” can be used in thesame way we use the word “dozen.” Just as “adozen” can refer to twelve cakes or twelve chick-ens, so “mole” always describes the same number

of molecules. The ratio of mass between a moleof Element A and Element B, or Compound Aand Compound B, is the same as the ratiobetween the mass of Atom A and Atom B, orMolecule A and Molecule B. Avogadro’s lawdescribes the connection between gas volumeand number of moles. According to Avogadro’slaw, if the volume of gas is increased underisothermal and isobarometric conditions, thenumber of moles also increases. The ratiobetween volume and number of moles is there-fore a constant.

The Ideal Gas Law

Once again, it is easy to see how Avogadro’s lawcan be related to the laws discussed earlier. Likethe other three, this one involves the parametersof temperature, pressure, and volume, but it alsointroduces a fourth—quantity of molecules (thatis, number of moles). In fact, all the laws so fardescribed are brought together in what is knownas the ideal gas law, sometimes called the com-bined gas law.

The ideal gas law can be stated as a formula,pV = nRT, where p stands for pressure, V for vol-ume, n for number of moles, and T for tempera-ture. R is known as the universal gas constant, afigure equal to 0.0821 atm • liter/mole • K. (Likemost figures in chemistry, this one is bestexpressed in metric rather than English units.)

Given the equation pV = nRT and the factthat R is a constant, it is possible to find the valueof any one variable—pressure, volume, numberof moles, or temperature—as long as one knowsthe value of the other three. The ideal gas law alsomakes it possible to discern certain relationships:thus, if a gas is in a relatively cool state, the prod-uct of its pressure and volume is proportionatelylow; and if heated, its pressure and volume prod-uct increases correspondingly.

The Kinetic Theory of Gases

From the preceding gas laws, a set of proposi-tions known collectively as the kinetic theory ofgases has been derived. Collectively, these putforth the proposition that a gas consists ofnumerous molecules, relatively far apart in space,which interact by colliding. These collisions areresponsible for the production of thermal energy,because when the velocity of the moleculesincreases—as it does after collision—the temper-ature increases as well.



Gases There are five basic postulates to the kinetictheory of gases:

• 1. Gases consist of tiny molecular or atomicparticles.

• 2. The proportion between the size of theseparticles and the distances between them isso small that the individual particles can beassumed to have negligible volume.

• 3. These particles experience continual ran-dom motion. When placed in a container,their collisions with the walls of the con-tainer constitute the pressure exerted by thegas.

• 4. The particles neither attract nor repel oneanother.

• 5. The average kinetic energy of the particlesin a gas is directly related to absolute tem-perature.

These observations may appear to resemblestatements made earlier concerning the differ-ences between gases, liquids, and solids in termsof molecular behavior. If so, that is no accident:the kinetic theory constitutes a generally accept-ed explanation for the reasons why gases behaveas they do. Kinetic theories do not work as wellfor explaining the behaviors of solids and liquids;nonetheless, they do go a long way toward iden-tifying the molecular properties inherent in thevarious phases of matter.

Laws of Partial Pressure

In addition to all the gas laws so far discussed,two laws address the subject of partial pressure.When two or more gases are present in a con-tainer, partial pressure is the pressure that one ofthem exerts if it alone is in the container.

Dalton’s law of partial pressure states thatthe total pressure of a gas is equal to the sum ofits partial pressures. As noted earlier, air is com-posed mostly of nitrogen and oxygen. Along withthese are small components, carbon dioxide, andgases collectively known as the rare or noblegases: argon, helium, krypton, neon, radon, andxenon. Hence, the total pressure of a given quan-tity of air is equal to the sum of the pressuresexerted by each of these gases.

Henry’s law states that the amount of gasdissolved in a liquid is directly proportional tothe partial pressure of the gas above the surfaceof the solution. This applies only to gases such asoxygen and hydrogen that do not react chemical-ly to liquids. On the other hand, hydrochloric

acid will ionize when introduced to water: one ormore of its electrons will be removed, and itsatoms will convert to ions, which are either posi-tive or negative in charge.

Applications of Dalton’s andHenry’s Laws

PARTIAL PRESSURE: A MAT-

TER OF LIFE AND POSSIBLE

DEATH FOR SCUBA DIVERS. Thegas laws are not just a series of abstract state-ments. Certainly, they do concern the behavior ofideal as opposed to real gases. Like all scientificmodels, they remove from the equation all out-side factors, and treat specific properties in isola-tion. Yet, the behaviors of the ideal gasesdescribed in the gas laws provide a key to under-standing the activities of real gases in the realworld. For instance, the concept of partial pressure helps scuba divers avoid a possibly fatalsickness.

Imagine what would happen if a substancewere to bubble out of one’s blood like carbondioxide bubbling out of a soda can, as describedbelow. This is exactly what can happen to anundersea diver who returns to the surface tooquickly: nitrogen rises up within the body, pro-ducing decompression sickness—known collo-quially as “the bends.” This condition may mani-fest as itching and other skin problems, jointpain, choking, blindness, seizures, unconscious-ness, permanent neurological defects such asparaplegia, and possibly even death.

If a scuba diver descending to a depth of 150ft (45.72 m) or more were to use ordinary air inhis or her tanks, the results would be disastrous.The high pressure exerted by the water at suchdepths creates a high pressure on the air in thetank, meaning a high partial pressure on thenitrogen component in the air. The result wouldbe a high concentration of nitrogen in the blood,and hence the bends.

Instead, divers use a mixture of helium andoxygen. Helium gas does not dissolve well inblood, and thus it is safer for a diver to inhale thisoxygen-helium mixture. At the same time, theoxygen exerts the same pressure that it wouldnormally—in other words, it operates in accor-dance with Dalton’s observations concerningpartial pressure.

OPENING A SODA CAN. Inside acan or bottle of carbonated soda is carbon diox-



Gaseside gas (CO2), most of which is dissolved in thedrink itself. But some of it is in the space (some-times referred to as “head space”) that makes upthe difference between the volume of the softdrink and the volume of the container.

At the bottling plant, the soda manufactureradds high-pressure carbon dioxide (CO2) to thehead space in order to ensure that more CO2 willbe absorbed into the soda itself. This is in accor-dance with Henry’s law: the amount of gas (inthis case CO2) dissolved in the liquid (soda) isdirectly proportional to the partial pressure ofthe gas above the surface of the solution—that is,the CO2 in the head space. The higher the pres-sure of the CO2 in the head space, the greater theamount of CO2 in the drink itself; and the greaterthe CO2 in the drink, the greater the “fizz” of thesoda.

Once the container is opened, the pressurein the head space drops dramatically. Once again,Henry’s law indicates that this drop in pressurewill be reflected by a corresponding drop in theamount of CO2 dissolved in the soda. Over aperiod of time, the soda will release that gas, andeventually, it will go “flat.”

FIRE EXTINGUISHERS. A fireextinguisher consists of a long cylinder with anoperating lever at the top. Inside the cylinder is atube of carbon dioxide surrounded by a quantityof water, which creates pressure around the CO2

tube. A siphon tube runs vertically along thelength of the extinguisher, with one opening inthe water near the bottom. The other end opensin a chamber containing a spring mechanismattached to a release valve in the CO2 tube.

The water and the CO2 do not fill the entirecylinder: as with the soda can, there is “headspace,” an area filled with air. When the operatinglever is depressed, it activates the spring mecha-nism, which pierces the release valve at the top ofthe CO2 tube. When the valve opens, the CO2

spills out in the “head space,” exerting pressureon the water. This high-pressure mixture ofwater and carbon dioxide goes rushing out of thesiphon tube, which was opened when the releasevalve was depressed. All of this happens, ofcourse, in a fraction of a second—plenty of timeto put out the fire.

AEROSOL CANS. Aerosol cans aresimilar in structure to fire extinguishers, thoughwith one important difference. As with the fire

extinguisher, an aerosol can includes a nozzlethat depresses a spring mechanism, which in turnallows fluid to escape through a tube. But insteadof a gas cartridge surrounded by water, most of the can’s interior is made up of the product(for instance, deodorant), mixed with a liquidpropellant.

The “head space” of the aerosol can is filledwith highly pressurized propellant in gas form,and, in accordance with Henry’s law, a corre-sponding proportion of this propellant is dis-solved in the product itself. When the nozzle isdepressed, the pressure of the propellant forcesthe product out through the nozzle.

A propellant, as its name implies, propels theproduct itself through the spray nozzle when thenozzle is depressed. In the past, chlorofluorocar-bons (CFCs)—manufactured compounds con-taining carbon, chlorine, and fluorine atoms—were the most widely used form of propellant.Concerns over the harmful effects of CFCs on theenvironment, however, has led to the develop-ment of alternative propellants, most notablyhydrochlorofluorocarbons (HCFCs), CFC-likecompounds that also contain hydrogen atoms.

Applications of Boyle’s,Charles’s, and Gay-Lussac’sLaws

WHEN THE TEMPERATURE

CHANGES. A number of interesting resultsoccur when gases experience a change in temper-ature, some of them unfortunate and somepotentially lethal. In these instances, it is possibleto see the gas laws—particularly Boyle’s andCharles’s—at work.

There are numerous examples of the disas-trous effects that result from an increase in thetemperature of combustible gases, including nat-ural gas and petroleum-based products. In addi-tion, the pressure on the gases in aerosol cansmakes the cans highly explosive—so much sothat discarded cans at a city dump may explodeon a hot summer day. Yet, there are otherinstances when heating a gas can produce posi-tive effects.

A hot-air balloon, for instance, floatsbecause the air inside it is not as dense than theair outside. According to Charles’s law, heating agas will increase its volume, and since gas mole-cules exert little attraction toward one another,



Gases they tend to “spread out” even further with anincrease of volume. This, in turn, creates a signif-icant difference in density between the air in theballoon and the air outside, and as a result, theballoon floats.

Although heating a gas can be beneficial,cooling a gas is not always a wise idea. If someonewere to put a bag of potato chips into a freezer,thinking this would preserve their flavor, hewould be in for a disappointment. Much of whatmaintains the flavor of the chips is the pressur-ization of the bag, which ensures a consistentinternal environment so that preservative chem-icals, added during the manufacture of the chips,can keep them fresh. Placing the bag in the freez-er causes a reduction in pressure, as per Gay-Lus-sac’s law, and the bag ends up a limp version of itsformer self.

Propane tanks and tires offer an example ofthe pitfalls that may occur by either allowing agas to heat up or cool down by too much.Because most propane tanks are made accordingto strict regulations, they are generally safe, but itis not entirely inconceivable that the extremeheat of a summer day could cause a defectivetank to burst. An increase in temperature leads toan increase in pressure, in accordance with Gay-Lussac’s law, and could lead to an explosion.

Because of the connection between heat andpressure, propane trucks on the highways duringthe summer are subjected to weight tests toensure that they are not carrying too much gas.On the other hand, a drastic reduction in tem-perature could result in a loss in gas pressure. If apropane tank from Florida were transported bytruck during the winter to northern Canada, thepressure is dramatically reduced by the time itreaches its destination.

THE INTERNAL-COMBUSTION

ENGINE. In operating a car, we experiencetwo applications of the gas laws. One of these iswhat makes the car run: the combustion of gasesin the engine, which illustrates the interrelationof volume, pressure, and temperature expressedin the laws attributed to Boyle, Charles, and Gay-Lussac. The other is, fortunately, a less frequentphenomenon—but it can and does save lives.This is the operation of an airbag, whichdepends, in part, on the behaviors explained inCharles’s law.

When the driver of a modern, fuel-injectionautomobile pushes down on the accelerator, this

activates a throttle valve that sprays droplets ofgasoline mixed with air into the engine. The mix-ture goes into the cylinder, where the pistonmoves up, compressing the gas and air. While themixture is still at a high pressure, the electricspark plug produces a flash that ignites the gaso-line-air mixture. The heat from this controlledexplosion increases the volume of air, whichforces the piston down into the cylinder. Thisopens an outlet valve, causing the piston to riseand release exhaust gases.

As the piston moves back down again, aninlet valve opens, bringing another burst of gaso-line-air mixture into the chamber. The piston,whose downward stroke closed the inlet valve,now shoots back up, compressing the gas and airto repeat the cycle. The reactions of the gasolineand air to changes in pressure, temperature, andvolume are what move the piston, which turns acrankshaft that causes the wheels to rotate.

THE AIRBAG. So much for moving—what about stopping? Most modern cars areequipped with an airbag, which reacts to suddenimpact by inflating. This protects the driver andfront-seat passenger, who, even if they are wear-ing seatbelts, may otherwise be thrown againstthe steering wheel or dashboard.

In order to perform its function properly,the airbag must deploy within 40 milliseconds(0.04 seconds) of impact. Not only that, but it hasto begin deflating before the body hits it. If a per-son’s body, moving forward at speeds typical inan automobile accident, were to smash against afully inflated airbag, it would feel like hittingconcrete—with all the expected results.

The airbag’s sensor contains a steel ballattached to a permanent magnet or a stiff spring.The spring or magnet holds the ball in placethrough minor mishaps when an airbag is notwarranted—for instance, if a car were simply tobe “tapped” by another in a parking lot. But in acase of sudden deceleration, the magnet orspring releases the ball, sending it down a smoothbore. The ball flips a switch, turning on an elec-trical circuit. This in turn ignites a pellet of sodi-um azide, which fills the bag with nitrogen gas.

At this point, the highly pressurized nitrogengas molecules begin escaping through vents.Thus, as the driver’s or rider’s body hits theairbag, the deflation of the bag is moving it in thesame direction that the body is moving—only



Gases


ABSOLUTE TEMPERATURE: Tem-

perature in relation to absolute zero

(-273.15°C or -459.67°F), as measured on

the Kelvin scale. The Kelvin and Celsius

scales are directly related; hence, Celsius

temperatures can be converted to Kelvins

(for which neither the word nor the sym-

bol for “degree” are used) by adding

273.15.

ATMOSPHERE: A measure of pres-

sure, abbreviated “atm” and equal to the

average pressure exerted by air at sea level.

In English units, this is equal to 14.7 lb/in2,

and in SI units, to 101,300 pascals.

AVOGADRO’S LAW: A statement,

derived by the Italian physicist Amedeo

Avogadro (1776-1856), which holds that as

the volume of gas increases under isother-

mal and isobarometric conditions, the

number of molecules (expressed in terms

of mole number), increases as well. Thus,

the ratio of volume to mole number is a

constant.

BAROMETER: An instrument for

measuring atmospheric pressure.

BOYLE’S LAW: A statement, derived

by English chemist Robert Boyle (1627-

1691), which holds that for gases in

isothermal conditions, an inverse relation-

ship exists between the volume and pres-

sure of a gas. This means that the greater

the pressure, the less the volume and vice

versa, and therefore the product of pres-

sure multiplied by volume yields a constant

figure.

CHARLES’S LAW: A statement,

derived by French physicist and chemist

J. A. C. Charles (1746-1823), which holds

that for gases in isobarometric conditions,

the ratio between the volume and temper-

ature of a gas is constant. This means that

the greater the temperature, the greater the

volume, and vice versa.

DALTON’S LAW OF PARTIAL PRES-

SURE: A statement, derived by the

English chemist John Dalton (1766-1844),

which holds that the total pressure of a gas

is equal to the sum of its partial pres-

sures—that is, the pressure exerted by each

component of the gas mixture.

GAS: A phase of matter in which mole-

cules move at high speeds, and therefore

exert little or no attraction toward one

another.

GAS LAWS: A series of statements

concerning the behavior of gases in

response to changes in temperature, pres-

sure, and volume. The gas laws, developed

by scientists during the seventeenth, eigh-

teenth, and nineteenth centuries, include

Avogadro’s law, Boyle’s law, Charles’s law,

Dalton’s law of partial pressures, Gay-Lus-

sac’s law, and Henry’s law. These are

summed up in the ideal gas law. The kinet-

ic theory of gases is based on observations

garnered from these laws.

GAY-LUSSAC’S LAW: A statement,

derived by French physicist and chemist

Joseph Gay-Lussac (1778-1850), which

holds that the pressure of a gas is directly

related to its absolute temperature. Hence,

the ratio of pressure to absolute tempera-

ture is a constant.

HENRY’S LAW: A statement, derived

KEY TERMS


Gases


much, much more slowly. Two seconds afterimpact, which is an eternity in terms of theprocesses involved, the pressure inside the baghas returned to 1 atm.

The chemistry of the airbag is particularlyinteresting. The bag releases inert, or non-reac-tive, nitrogen gas, which poses no hazard tohuman life; yet one of the chemical ingredients in

by English chemist William Henry (1774-

1836), which holds that the amount of gas

dissolved in a liquid is directly proportion-

al to the partial pressure of the gas above

the solution. This holds true only for gases,

such as hydrogen and oxygen, which do

not react chemically to liquids.

IDEAL GAS LAW: A proposition, also

known as the combined gas law, which

draws on all the gas laws. The ideal gas law

can be expressed as the formula pV = nRT,

where p stands for pressure, V for volume,

n for number of moles, and T for tempera-

ture. R is known as the universal gas con-

stant, a figure equal to 0.0821 atm •

liter/mole • K.

ISOTHERMAL: Referring to a situa-

tion in which temperature is kept constant.

ISOBAROMETRIC: Referring to a sit-

uation in which pressure is kept constant.


an object possesses by virtue of its motion.

KINETIC THEORY OF GASES: A set

of propositions describing a gas as consist-

ing of numerous molecules, relatively far

apart in space, which interact by colliding.

These collisions are responsible for the

production of thermal energy, because

when the velocity of the molecules increas-

es—as it does after collision—the temper-

ature increases as well.

MILLIMETER OF MERCURY: Anoth-

er name for the torr, abbreviated mm Hg.


“amount of substance.” The quantity of

molecules or atoms in a mole is, generally

speaking, the same as Avogadro’s number:

6.022137 • 1023. However, in the more pre-

cise SI definition, a mole is equal to the

number of carbon atoms in 12.01 g of car-

bon.

PARTIAL PRESSURE: When two or

more gases are present in a container, par-

tial pressure is the pressure that one of

them exerts if it alone is in the container.

Dalton’s law of partial pressure and

Henry’s law relate to the partial pressure of

gases.

PASCAL: The principle SI or metric

unit of pressure, abbreviated “Pa” and

equal to 1 N/m2.

PRESSURE: The ratio of force to sur-

face area, when force is applied in a direc-

tion perpendicular to that surface.

THERMAL ENERGY: A form of kinet-

ic energy—commonly called “heat”—that

is produced by the movement of atomic or

molecular particles. The greater the

motion of these particles relative to one

another, the greater the thermal energy.

TORR: An SI unit, also known as the

millimeter of mercury, that represents the

pressure required to raise a column of mer-

cury 1 mm. The torr, equal to 133 Pascals,

is named for the Italian physicist Evange-

lista Torricelli (1608-1647), who invented

the barometer.



Gasesthe airbag is so lethal that some environmentalistgroups have begun to raise concerns over itspresence in airbags. This is sodium azide (NaN3),one of three compounds—along with potassiumnitrate (KNO3) and silicon dioxide (SiO2)—pres-ent in an airbag prior to inflation.

The sodium azide and potassium nitratereact to one another, producing a burst of hotnitrogen gas in two back-to-back reactions. Inthe fractions of a second during which thisoccurs, the airbag becomes like a solid-rocketbooster, experiencing a relatively slow detonationknown as “deflagration.”

The first reaction releases nitrogen gas,which fills the bag, while the second reactionleaves behind the by-products potassium oxide(K2O) and sodium oxide (Na2O). These combinewith the silicon dioxide to produce a safe, stablecompound known as alkaline silicate. The latter,similar to the sand used for making glass, is allthat remains in the airbag after the nitrogen gashas escaped.

WHERE TO LEARN MORE

“Atmospheric Pressure: The Force Exerted by the Weight ofAir” (Web site). <http://kids.earth.nasa.gov/archive/air_pressure/> (April 7, 2001).

“Chemical Sciences Structure: Structure of Matter: Natureof Gases” University of Alberta Chemistry Department(Web site). <http://www.chem.ualberta.ca/~plam-beck/che/struct/s0307.htm> (May 12, 2001).

“Chemistry Units: Gas Laws.”<http://bio.bio.rpi.edu/MS99/ausemaW/chem/gases.html> (February 21, 2001).

“Homework: Science: Chemistry: Gases” Channelone.com(Web site).<http://www.channelone.com/fasttrack/science/chemistry/gases.html> (May 12, 2001).

“Kinetic Theory of Gases: A Brief Review” University ofVirginia Department of Physics (Web site).<http://www.phys.virginia.edu/classes/252/kinetic_theory.html> (April 15, 2001).

Laws of Gases. New York: Arno Press, 1981.

Macaulay, David. The New Way Things Work. Boston:Houghton Mifflin, 1998.

Mebane, Robert C. and Thomas R. Rybolt. Air and OtherGases. Illustrations by Anni Matsick. New York:Twenty-First Century Books, 1995.

“Tutorials—6.” Chemistrycoach.com (Web site).<http://www.chemistrycoach.com/tutorials-6.htm>(February 21, 2001).




61

SCIENCE OF EVERYDAY THINGSreal-l ife chemistry

ATOMS AND MOLECULES

ATOMS AND MOLECULES

ATOMS

ATOMIC MASS

ELECTRONS

ISOTOPES

IONS AND IONIZATION

MOLECULES


63SCIENCE OF EVERYDAY THINGS VOLUME 1: REAL-LIFE CHEMISTRY

ATOMSAtoms

CONCEPTOur world is made up of atoms, yet the atomicmodel of the universe is nonetheless considered a“theory.” When scientists know beyond all rea-sonable doubt that a particular principle is thecase, then it is dubbed a law. Laws address the factthat certain things happen, as well as how theyhappen. A theory, on the other hand, attempts toexplain why things happen. By definition, an ideathat is dubbed a theory has yet to be fully proven,and such is the case with the atomic theory ofmatter. After all, the atom cannot be seen, evenwith electron microscopes—yet its behavior canbe studied in terms of its effects. Atomic theoryexplains a great deal about the universe, includ-ing the relationship between chemical elements,and therefore (as with Darwin’s theory concern-ing biological evolution), it is generally acceptedas fact. The particulars of this theory, includingthe means by which it evolved over the centuries,are as dramatic as any detective story. Nonethe-less, much still remains to be explained about theatom—particularly with regard to the smallestitems it contains.

HOW IT WORKS

Why Study Atoms?

Many accounts of the atom begin with a historyof the growth in scientists’ understanding of itsstructure, but here we will take the oppositeapproach, first discussing the atom in terms ofwhat physicists and chemists today understand.Only then will we examine the many challengesscientists faced in developing the current atomicmodel: false starts, wrong theories, right roadsnot taken, incomplete models. In addition, we

will explore the many insights added along theway as, piece by piece, the evidence concerningatomic behavior began to accumulate.

People who are not scientifically trainedtend to associate studies of the atom withphysics, not chemistry. While it is true that physi-cists study atomic structure, and that much ofwhat scientists know today about atoms comesfrom the work of physicists, atomic studies areeven more integral to chemistry than to physics.At heart, chemistry is about the interaction ofdifferent atomic and molecular structures: theirproperties, their reactions, and the ways in whichthey bond.

WHAT THE ATOM MEANS TO

CHEMISTRY. Just as a writer in Englishworks with the 26 letters of the alphabet, achemist works with the 100-plus known ele-ments, the fundamental and indivisible sub-stances of all matter. And what differentiates theelements, ultimately, from one another is nottheir color or texture, or even the phase of mat-ter—solid, gas, or liquid—in which they are nor-mally found. Rather, the defining characteristicof an element is the atom that forms its basicstructure.

The number of protons in an atom is thecritical factor in differentiating between ele-ments, while the number of neutrons alongsidethe protons in the nucleus serves to distinguishone isotope from another. However, as importantas elements and even isotopes are to the work ofa chemist, the components of the atom’s nucleushave little direct bearing on the atomic activitythat brings about chemical reactions and chemi-cal bonding. All the chemical “work” of an atomis done by particles vastly smaller in mass than


Atoms element. In fact, these two definitions do notform a closed loop, as they would if it were stat-ed that an element is something made up ofatoms. Every item of matter that exists, except forthe subatomic particles discussed in this essay, ismade up of atoms. An element, on the otherhand, is—as stated in its definition—made up ofonly one kind of atom. “Kind of atom” in thiscontext refers to the number of protons in itsnucleus.

Protons are one of three basic subatomicparticles, the other two being electrons and neu-trons. As we shall see, there appear to be particleseven smaller than these, but before approachingthese “sub-subatomic” particles, it is necessary toaddress the three most significant components ofan atom. These are distinguished from oneanother in terms of electric charge: protons arepositively charged, electrons are negative incharge, and neutrons have no electrical charge.As with the north and south poles of magnets,positive and negative charges attract one another,whereas like charges repel. Atoms have no netcharge, meaning that the protons and electronscancel out one another.

EVOLVING MODELS OF THE

ATOM. Scientists originally thought of an atomas a sort of closed sphere with a relatively hardshell, rather like a ball bearing. Nor did they ini-tially understand that atoms themselves are divis-ible, consisting of the parts named above. Even asawareness of these three parts emerged in the lastyears of the nineteenth century and the first partof the twentieth, it was not at all clear how theyfit together.

At one point, scientists believed that elec-trons floated in a cloud of positive charges. Thiswas before the discovery of the nucleus, wherethe protons and neutrons reside at the heart ofthe atom. It then became clear that electronswere moving around the nucleus, but how? For atime, a planetary model seemed appropriate: inother words, electrons revolved around thenucleus much as planets orbit the Sun. Eventual-ly, however—as is often the case with scientificdiscovery—this model became unworkable, andhad to be replaced by another.

The model of electron behavior acceptedtoday depicts the electrons as forming a cloudaround the nucleus—almost exactly the oppositeof what physicists believed a century ago. The useof the term “cloud” may perhaps be a bit mis-

64 SCIENCE OF EVERYDAY THINGSVOLUME 1: REAL-LIFE CHEMISTRY

either the protons or neutrons—fast-moving lit-tle bundles of energy called electrons.

Moving rapidly through the space betweenthe nucleus and the edge of the atom, electronssometimes become dislodged, causing the atomto become a positively charged ion. Conversely,sometimes an atom takes on one or more elec-trons, thus acquiring a negative charge. Ions arecritical to the formation of some kinds of chem-ical bonds, but the chemical role of the electronis not limited to ionic bonds.

In fact, what defines an atom’s ability tobond with another atom, and therefore to form amolecule, is the specific configuration of its elec-trons. Furthermore, chemical reactions are theresult of changes in the arrangement of electrons,not of any activity involving protons or neutrons.So important are electrons to the interactionsstudied in chemistry that a separate essay hasbeen devoted to them.

What an Atom Is

BASIC ATOMIC STRUCTURE. Thedefinitions of atoms and elements seems, at firstglance, almost circular: an element is a substancemade up of only one kind of atom, and an atomis the smallest particle of an element that retainsall the chemical and physical properties of the

JOHN DALTON.


Atomsleading, implying as it does something that sim-ply hovers. In fact, the electron, under normalcircumstances, is constantly moving. The pathsof its movement around the nucleus are nothinglike that of a planet’s orbit, except inasmuch asboth models describe a relatively small objectmoving around a relatively large one.

The furthest edges of the electron’s move-ment define the outer perimeters of the atom.Rather than being a hard-shelled little nugget ofmatter, an atom—to restate the metaphor men-tioned above—is a cloud of electrons surround-ing a nucleus. Its perimeters are thus not sharplydelineated, just as there is no distinct barrierbetween Earth’s atmosphere and space itself. Justas the air gets thinner the higher one goes, so it iswith an atom: the further a point is from thenucleus, the less the likelihood that an electronwill pass that point on a given orbital path.

Nucleons

MASS NUMBER AND ATOMIC

NUMBER. The term nucleon is used generi-cally to describe the relatively heavy particles thatmake up an atomic nucleus. Just as “sport” canrefer to football, basketball, or baseball, or anyother item in a similar class, such as soccer ortennis, “nucleon” refers to protons and neutrons.The sum of protons and neutrons is sometimescalled the nucleon number, although a morecommonly used term is mass number.

Though the electron is the agent of chemicalreactions and bonding, it is the number of pro-tons in the nucleus that defines an atom as to itselement. Atoms of the same element always havethe same number of protons, and since this fig-ure is unique for a given element, each element isassigned an atomic number equal to the numberof protons in its nucleus. The atoms are listed inorder of atomic number on the periodic table ofelements.

ATOMIC MASS AND ISOTOPES.

A proton has a mass of 1.673 • 10–24 g, which isvery close to the established figure for measuringatomic mass, the atomic mass unit. At one time,the basic unit of atomic mass was equal to themass of one hydrogen atom, but hydrogen is soreactive—that is, it tends to combine readily withother atoms to form a molecule, and hence acompound—that it is difficult to isolate. Instead,the atomic mass unit is today defined as 1/12 of


the mass of a carbon-12 atom. That figure isexactly 1.66053873 • 10–24 grams.

The mention of carbon-12, a substancefound in all living things, brings up the subject ofisotopes. The “12” in carbon-12 refers to its massnumber, or the sum of protons and neutrons.Two atoms may be of the same element, and thushave the same number of protons, yet differ intheir number of neutrons—which means a dif-ference both in mass number and atomic mass.Such differing atoms of the same element arecalled isotopes. Isotopes are often designated bysymbols showing mass number to the upper leftof the chemical symbol or element symbol—forinstance, 12C for carbon-12.

ELECTRIC CHARGE. Protons have apositive electric charge of 1, designated either as1+ or +1. Neutrons, on the other hand, have noelectric charge. It appears that the 1+ charge of aproton and the 0 charge of a neutron are theproducts of electric charges on the part of evensmaller particles called quarks. A quark mayeither have a positive electric charge of less than1+, in which case it is called an “up quark”; or anegative charge of less than 1–, in which case it iscalled a “down quark.”

Research indicates that a proton containstwo up quarks, each with a charge of 2/3+, and

NIELS BOHR.


Atoms


one down quark with a charge of 1/3–. Thisresults in a net charge of 1+. On the other hand,a neutron is believed to hold one up quark witha charge of 2/3+, and two down quarks withcharges of 1/3– each. Thus, in the neutron, the upand down quarks cancel out one another, and thenet charge is zero.

A neutron has about the same mass as a pro-ton, but other than its role in forming isotopes,the neutron’s function is not exactly clear. Per-haps, it has been speculated, it binds protons—which, due to their positive charges, tend to repelone another—together at the nucleus.

Electrons

An electron is much smaller than a proton orneutron, and has much less mass; in fact, its massis equal to 1/1836 that of a proton, and 1/1839that of a neutron. Yet the area occupied by elec-trons—the region through which they move—constitutes most of the atom’s volume. If thenucleus of an atom were the size of a BB (which,in fact, is billions of times larger than a nucleus),the furthest edge of the atom would be equiva-lent to the highest ring of seats around an indoorsports arena. Imagine the electrons as incrediblyfast-moving insects buzzing constantly throughthe arena, passing by the BB but then flitting to

the edges or points in between, and you havesomething approaching an image of the atom’sinterior.

How fast does an electron move? Speedsvary depending on a number of factors, but it canmove nearly as fast as light: 186,000 mi (299,339km) per second. On the other hand, for an itemof matter near absolute zero in temperature, thevelocity of the electron is much, much less. In anycase, given the fact that an electron has enoughnegative charge to cancel out that of the proton,it must be highly energized. After all, this wouldbe like an electric generator weighing 1 lb havingas much power as a generator that weighed 1 ton.

According to what modern scientists knowor hypothesize concerning the inner structure ofthe atom, electrons are not made up of quarks;rather, they are part of a class of particles calledleptons. It appears that leptons, along withquarks and what are called exchange particles,constitute the elementary particles of atoms—particles on a much more fundamental level thanthat of the proton and neutron.

Electrons are perhaps the most intriguingparts of an atom. Their mass is tiny, even inatomic terms, yet they possess enough charge tocounteract a “huge” proton. They are capable, incertain situations, of moving from one atom to

THE EVOLUTION OF ATOMIC THEORY.


Atomsanother, thus creating ions, and depending ontheir highly complex configuration and ability torearrange their configuration, they facilitate orprevent chemical reactions.

REAL-LIFEAPPLICATIONS

Ancient Greek Theories of Matter

The first of the Greek philosophers, and the firstindividual in Western history who deserves to becalled a scientist, was Thales (c. 625-c. 547 B.C.) ofMiletus. (Miletus is in Greek Asia Minor, nowpart of Turkey.) Among his many achievementswere the correct prediction of a solar eclipse, andone of the first-ever observations of electricity,when he noted the electrification of amber byfriction.

But perhaps the greatest of Thales’s legacieswas his statement that “Everything is water.” Thisrepresented the first attempt to characterize thenature of all physical reality. It set off a debateconcerning the fundamental nature of matterthat consumed Greek philosophers for two cen-turies. Later, philosophers attempted to charac-terize matter in terms of fire or air. In time, how-ever, there emerged a school of thought con-cerned not with identifying matter as one partic-ular thing or another, but with recognizing astructural consistency in all of matter. Amongthese were the philosophers Leucippus (c. 480-c.420 B.C.) and his student Democritus (c. 460-370B.C.)

DEMOCRITUS’S “ATOMS”. Leu-cippus and Democritus proposed a new andhighly advanced model for the tiniest point ofphysical space. Democritus, who actually articu-lated these ideas (far less is known about Leucip-pus) began with a “thought experiment,” imagin-ing what would happen if an item of matter weresubdivided down to its smallest piece. This tini-est fragment, representing an item of matter thatcould not be cut into smaller pieces, he called bya Greek term meaning “no cut”: atomos.

Democritus was not necessarily describingmatter in a concrete, scientific way: his “atoms”were idealized philosophical constructs ratherthan purely physical units. Yet, he came amazing-ly close, and indeed much closer than any thinkerfor the next 22 centuries, to identifying the fun-

damental structure of physical reality. Why did ittake so long for scientists to come back around tothe atomic model? The principal culprit, whoadvanced an erroneous theory of matter, alsohappened to be one of the greatest thinkers of alltime: Aristotle (384-322 B.C..)

ARISTOTLE’S “ELEMENTS”.

Aristotle made numerous contributions to sci-ence, including his studies in botany and zoolo-gy, as well as his explanation of the four causes, asignificant attempt to explain events by meansother than myth or superstition. In the area ofthe physical sciences, however, Aristotle’s impactwas less than beneficial. Most notably, in explain-ing why objects fall when dropped, he claimedthat the ground was their “natural” destination—a fallacy later overturned with the gravitationalmodel developed by Galileo Galilei (1564-1642)and Sir Isaac Newton (1642-1727).

The ideas Aristotle put forward concerningwhat he called “natural motion” were a productof his equally faulty theories with regard to whattoday’s scientists refer to as chemistry. In ancienttimes, chemistry, as such, did not exist. Longbefore Aristotle’s time, Egyptian embalmers andmetallurgists used chemical processes, but theydid so in a practical, applied manner, exerting lit-tle effort toward what could be described as sci-entific theory. Philosophers such as Aristotle,who were some of the first scientists, made littledistinction between physical and chemicalprocesses. Thus, whereas physics is understoodtoday as an important background for chemistry,Aristotle’s “physics” was actually an outgrowth ofhis “chemistry.”

Rejecting Democritus’s atomic model, Aris-totle put forward his own view of matter. LikeDemocritus, he believed that matter was com-posed of very small components, but these heidentified not as atoms, but as “elements”: earth,air, fire, and water. He maintained that all objectsconsisted, in varying degrees, of one or more ofthese, and based his explanation of gravity on therelative weights of each element. Water sits ontop of the earth, he explained, because it islighter, yet air floats above the water because it islighter still—and fire, lightest of all, rises highest.Furthermore, he claimed that the planets beyondEarth were made up of a “fifth element,” or quin-tessence, of which little could be known.

In fairness to Aristotle, it should be pointedout that it was not his fault that science all but



Atoms died out in the Western world during the periodfrom about A.D. 200 to about 1200. Furthermore,he did offer an accurate definition of an element,in a general sense, as “one of those simple bodiesinto which other bodies can be decomposed, andwhich itself is not capable of being divided intoothers.” As we shall see, the definition used todayis not very different from Aristotle’s. However, todefine an element scientifically, as modernchemists do, it is necessary to refer to somethingAristotle rejected: the atom. So great was hisopposition to Democritus’s atomic theory, andso enormous was Aristotle’s influence on learn-ing for more than 1,500 years following hisdeath, that scientists only began to reconsideratomic theory in the late eighteenth century.

A Maturing Concept of Elements

BOYLE’S IDEA OF ELEMENTS.

One of the first steps toward an understanding ofthe chemical elements came with the work ofEnglish physicist and chemist Robert Boyle(1627-1691). Building on the usable definition ofan element provided by Aristotle, Boyle main-tained no substance was an element if it could bebroken down into other substances. Thus, aircould be eliminated from the list of “elements,”because, clearly, it could be separated into morethan one elemental substance. (In fact, none ofthe four “elements” identified by Aristotle evenremotely qualifies as an element in modernchemistry.)

Boyle, nonetheless, still clung to aspects ofalchemy, a pseudo-science based on the transfor-mation of “base metals,” for example, the meta-morphosis of iron into gold. Though true chem-istry grew out of alchemy, the fundamentalproposition of alchemy was faulty: if one metalcan be turned into another, then that means thatmetals are not elements, which, in fact, they are.Nonetheless, Boyle’s studies led to the identifica-tion of numerous elements—that is, items thatreally are elements—in the years that followed.

LAVOISIER AND PROUST: CON-

STANT COMPOSITION. A few years afterBoyle came two French chemists who extendedscientific understanding of the elements. AntoineLavoisier (1743-1794) affirmed the definition ofan element as a simple substance that could notbe broken down into a simpler substance, and

noted that elements always react with one anoth-er in the same proportions.

Joseph-Louis Proust (1754-1826) put for-ward the law of constant composition, whichholds that a given compound always contains thesame proportions of mass between elements.Another chemist of the era had claimed that thecomposition of a compound varies in accordancewith the reactants used to produce it. Proust’s lawof constant composition made it clear that anyparticular compound will always have the samecomposition.

Early Modern Understandingof the Atom

DALTON AND AVOGADRO:

ATOMS AND MOLECULES. The workof Lavoisier and Proust influenced a critical fig-ure in the development of the atomic model:English chemist John Dalton (1766-1844). In ANew System of Chemical Philosophy (1808), Dal-ton put forward the idea that nature is composedof tiny particles, and in so doing he adoptedDemocritus’s word “atom” to describe these basicunits. This idea, which Dalton had formulatedfive years earlier, marked the starting-point ofmodern atomic theory.

Dalton recognized that the structure ofatoms in a particular element or compound isuniform, but maintained that compounds aremade up of compound atoms: in other words,water, for instance, is a compound of “wateratoms.” However, water is not an element, andthus, it was necessary to think of its atomic com-position in a different way—in terms of mole-cules rather than atoms. Dalton’s contemporaryAmedeo Avogadro (1776-1856), an Italian physi-cist, became the first scientist to clarify the dis-tinction between atoms and molecules.

The later development of the mole, whichprovided a means whereby equal numbers ofmolecules could be compared, paid tribute toAvogadro by designating the number of mole-cules in a mole as “Avogadro’s number.” Anothercontemporary, Swedish chemist Jons Berzelius(1779-1848), maintained that equal volumes ofgases at the same temperature and pressure con-tained equal numbers of atoms. Using this idea,he compared the mass of various reacting gases,and developed a system of comparing the mass ofvarious atoms in relation to the lightest one,hydrogen. Berzelius also introduced the system



Atomsof chemical symbols—H for hydrogen, O foroxygen, and so on—in use today.

BROWNIAN MOTION AND

KINETIC THEORY. Yet another figurewhose dates overlapped with those of Dalton,Avogadro, and Berzelius was Scottish botanistRobert Brown (1773-1858). In 1827, Brownnoted a phenomenon that later had an enormousimpact on the understanding of the atom. Whilestudying pollen grains under a microscope,Brown noticed that the grains underwent a curi-ous zigzagging motion in the water. The pollenassumed the shape of a colloid, a pattern thatoccurs when particles of one substance are dis-persed—but not dissolved—in another sub-stance. At first, Brown assumed that the motionhad a biological explanation—that is, it resultedfrom life processes within the pollen—but later,he discovered that even pollen from long-deadplants behaved in the same way.

Brown never understood what he was wit-nessing. Nor did a number of other scientists,who began noticing other examples of whatcame to be known as Brownian motion: the con-stant but irregular zigzagging of colloidal parti-cles, which can be seen clearly through a micro-scope. Later, however, Scottish physicist JamesClerk Maxwell (1831-1879) and others were ableto explain this phenomenon by what came to beknown as the kinetic theory of matter.

Kinetic theory is based on the idea that mol-ecules are constantly in motion: hence, the watermolecules were moving the pollen grains Brownobserved. Pollen grains are many thousands oftimes as large as water molecules, but since thereare so many molecules in even a drop of water,and their motion is so constant but apparentlyrandom, they are bound to move a pollen grainonce every few thousand collisions.

Mendeleev and the PeriodicTable

In 1869, Russian chemist Dmitri Mendeleev(1834-1907) introduced a highly useful systemfor organizing the elements, the periodic table.Mendeleev’s table is far more than just a handychart listing elements: at once simple and highlycomplex, it shows elements in order of increasingatomic mass, and groups together those exhibit-ing similar forms of chemical behavior andstructure.

Reading from right to left and top to bot-tom, the periodic table, as it is configured today,lists atoms in order of atomic number, generallyreflected by a corresponding increase in averageatomic mass. As Mendeleev observed, everyeighth element on the chart exhibits similar characteristics, and thus the chart is organized in columns representing specific groups ofelements.

The patterns Mendeleev observed were soregular that for any “hole” in his table, he pre-dicted that an element would be discovered thatwould fill that space. For instance, at one pointthere was a gap between atomic numbers 71 and73 (lutetium and tantalum, respectively).Mendeleev indicated that an atom would befound for the space, and 15 years after this pre-diction, the element germanium was isolated.

However, much of what defines an element’splace on the chart today relates to subatomic par-ticles—protons, which determine atomic num-ber, and electrons, whose configurations explaincertain chemical similarities. Mendeleev wasunaware of these particles: from the time he cre-ated his table, it was another three decades beforethe discovery of the first of these particles, theelectron. Instead, he listed the elements in anorder reflecting outward characteristics nowunderstood to be the result of the quantity anddistribution of protons and electrons.

Electromagnetism and Radiation

The contribution of Mendeleev’s contemporary,Maxwell, to the understanding of the atom wasnot limited to his kinetic theory. Building on thework of British physicist and chemist MichaelFaraday (1791-1867) and others, in 1865 he pub-lished a paper outlining a theory of a fundamen-tal interaction between electricity and magnet-ism. The electromagnetic interaction, as it laterturned out, explained something that gravita-tion, the only other form of fundamental inter-action known at the time, could not: the forcethat held together particles in an atom.

The idea of subatomic particles was still along time in coming, but the model of electro-magnetism helped make it possible. In the longrun, electromagnetism was understood toencompass a whole spectrum of energy radia-tion, including radio waves; infrared, visible, andultraviolet light; x rays; and gamma rays. But this,



Atoms too, was the product of work on the part ofnumerous individuals, among whom was Englishphysicist William Crookes (1832-1919).

In the 1870s, Crookes developed an appara-tus later termed a Crookes tube, with which hesought to analyze the “rays”—that is, radiation—emitted by metals. The tube consisted of a glassbulb, from which most of the air had beenremoved, encased between two metal plates orelectrodes, referred to as a cathode and an anode.A wire led outside the bulb to an electric source,and when electricity was applied to the elec-trodes, the cathodes emitted rays. Crookes con-cluded that the cathode rays were particles with anegative electric charge that came from the metalin the cathode plate.

RADIATION. In 1895, German physicistWilhelm Röntgen (1845-1923) noticed that pho-tographic plates held near a Crookes tubebecame fogged, and dubbed the rays that hadcaused the fogging “x rays.” A year after Röntgen’sdiscovery, French physicist Henri Becquerel(1852-1908) left some photographic plates in adrawer with a sample of uranium. Uranium hadbeen discovered more than a century before;however, there were few uses for it until Becquer-el discovered that the uranium likewise caused afogging of the photographic plates.

Thus radioactivity, a type of radiationbrought about by atoms that experience radioac-tive decay was discovered. The term was coinedby Polish-French physicist and chemist MarieCurie (1867-1934), who with her husband Pierre(1859-1906), a French physicist, was responsiblefor the discovery of several radioactive elements.

The Rise and Fall of thePlum Pudding Model

Working with a Crookes tube, English physicist J. J. Thomson (1856-1940) hypothesized that thenegatively charged particles Crookes hadobserved were being emitted by atoms, and in1897, he gave a name to these particles: electrons.The discovery of the electron raised a new ques-tion: if Thomson’s particles exerted a negativecharge, from whence did the counterbalancingpositive charge come?

An answer, of sorts, came from WilliamThomson, not related to the other Thomson and,in any case, better known by his title as LordKelvin (1824-1907). Kelvin compared the struc-ture of an atom to an English plum pudding: the

electrons were like raisins, floating in a positivelycharged “pudding”—that is, an undifferentiatedcloud of positive charges.

Kelvin’s temperature scale contributedgreatly to the understanding of molecularmotion as encompassed in the kinetic theory ofmatter. However, his model for the distributionof charges in an atom—charming as it may havebeen—was incorrect. Nonetheless, for severaldecades, the “plum pudding model,” as it came tobe known, remained the most widely accepteddepiction of the way that electric charges weredistributed in an atom. The overturning of theplum pudding model was the work of Englishphysicist Ernest Rutherford (1871-1937), a stu-dent of J. J. Thomson.

RUTHERFORD IDENTIFIES THE

NUCLEUS. Rutherford did not set out to dis-prove the plum pudding model; rather, he wasconducting tests to find materials that wouldblock radiation from reaching a photographicplate. The two materials he identified, whichwere, respectively, positive and negative in elec-tric charge, he dubbed alpha and beta particles.(An alpha particle is a helium nucleus stripped ofits electrons, such that it has a positive charge of2; beta particles are either electrons or positivelycharged subatomic particles called positrons. Thebeta particle Rutherford studied was an electronemitted during radioactive decay.)

Using a piece of thin gold foil with photo-graphic plates encircling it, Rutherford bom-barded the foil with alpha particles. Most of thealpha particles went straight through the foil—asthey should, according to the plum puddingmodel. However, a few particles were deflectedfrom their course, and some even bounced back.Rutherford later said it was as though he hadfired a gun at a piece of tissue paper, only to seethe tissue deflect the bullets. Analyzing theseresults, Rutherford concluded that there was no“pudding” of positive charges: instead, the atomhad a positively charged nucleus at its center.

The Nucleus Emerges

PROTONS AND ISOTOPES. Inaddition to defining the nucleus, Rutherford alsogave a name to the particles that imparted itspositive charge: protons. But just as the identifi-cation of the electron had raised new questionsthat, in being answered, led to the discovery ofthe proton, Rutherford’s achievement only



Atomsbrought up new anomalies concerning thebehavior of the nucleus.

Together with English chemist FrederickSoddy (1877-1956), Rutherford discovered thatwhen an atom emitted alpha or beta particles, itsatomic mass changed. Soddy had a name foratoms that displayed this type of behavior: iso-topes. Certain types of isotopes, Soddy andRutherford went on to conclude, had a tendencyto decay, moving toward stabilization, and thisdecay explained radioactivity.

CLARIFYING THE PERIODIC

TABLE. Soddy concluded that atomic mass, asmeasured by Berzelius, was actually an average ofthe mass figures for all isotopes within that ele-ment. This explained a problem withMendeleev’s periodic table, in which thereseemed to be irregularities in the increase ofatomic mass from element to element. Theanswer to these variations in mass, it turned out,related to the number of isotopes associated witha given element: the greater the number of iso-topes, the more these affected the overall meas-ure of the element’s mass.

By this point, physicists and chemists hadcome to understand that various levels of energyin matter emitted specific electromagnetic wave-lengths. Welsh physicist Henry Moseley (1887-1915) experimented with x rays, bombardingatoms of different elements with high levels ofenergy and observing the light they gave off asthey cooled. In the course of these tests, heuncovered an astounding mathematical relation-ship: the amount of energy a given element emit-ted was related to its atomic number.

Furthermore, the atomic number corre-sponded to the number of positive charges—thiswas in 1913, before Rutherford had named theproton—in the nucleus. Mendeleev had beenable to predict the discovery of new elements, butsuch predictions had remained problematic.When scientists understood the idea of atomicnumber, however, it became possible to predictthe existence of undiscovered elements withmuch greater accuracy.

NEUTRONS. Yet again, discoveries—thenucleus, protons, and the relationship betweenthese and atomic number—only created newquestions. (This, indeed, is one of the hallmarksof an active scientific theory. Rather than settlingquestions, science is about raising new ones, andthus improving the quality of the questions that

are asked.) Once Rutherford had identified theproton, and Moseley had established the numberof protons, the mystery at the heart of the atomonly grew deeper.

Scientists had found that the measured massof atoms could not be accounted for by the num-ber of protons they contained. Certainly, theelectrons had little to do with atomic mass: bythen it had been shown that the electron weighedabout 0.06% as much as a proton. Yet for all ele-ments other than protium (the first of threehydrogen isotopes), there was a discrepancybetween atomic mass and atomic number. Clear-ly, there had to be something else inside thenucleus.

In 1932, English physicist James Chadwick(1891-1974) identified that “something else.”Working with radioactive material, he found thata certain type of subatomic particle could pene-trate lead. All other known types of radiationwere stopped by the lead, and therefore, Chad-wick reasoned that this particle must be neutralin charge. In 1932, he won the Nobel Prize inPhysics for his discovery of the neutron.

The Nuclear Explosion

ISOTOPES AND RADIOACTIVI-

TY. Chadwick’s discovery clarified another mys-tery, that of the isotope, which had been raised byRutherford and Soddy several decades earlier.Obviously, the number of protons in a nucleusdid not change, but until the identification of theneutron, it had not been clear what it was thatdid change. At that point, it was understood thattwo atoms may have the same atomic number—and hence be of the same element—yet they may differ in number of neutrons, and thus be isotopes.

As the image of what an isotope was becameclearer, so too did scientists’ comprehension ofradioactivity. Radioactivity, it was discovered,was most intense where an isotope was the mostunstable—that is, in cases where an isotope hadthe greatest tendency to experience decay. Urani-um had a number of radioactive isotopes, such as235U, and these found application in the burgeon-ing realm of nuclear power—both the destruc-tive power of atomic bombs, and later the con-structive power of nuclear energy plants.

FISSION VS. FUSION. In nuclearfission, or the splitting of atoms, uranium iso-topes (or other radioactive isotopes) are bom-



Atoms barded with neutrons, splitting the uraniumnucleus in half and releasing huge amounts ofenergy. As the nucleus is halved, it emits severalextra neutrons, which spin off and split moreuranium nuclei, creating still more energy andsetting off a chain reaction. This explains thedestructive power in an atomic bomb, as well asthe constructive power—providing energy tohomes and businesses—in a nuclear power plant.Whereas the chain reaction in an atomic bombbecomes an uncontrolled explosion, in a nuclearplant the reaction is slowed and controlled.

Yet nuclear fission is not the most powerfulform of atomic reaction. As soon as scientistsrealized that it was possible to force particles outof a nucleus, they began to wonder if particlescould be forced into the nucleus. This type ofreaction, known as fusion, puts even nuclear fis-sion, with its awesome capabilities, to shame:nuclear fusion is, after all, the power of the Sun.On the surface of that great star, hydrogen atomsreach incredible temperatures, and their nucleifuse to create helium. In other words, one ele-ment actually transforms into another, releasingenormous amounts of energy in the process.

NUCLEAR ENERGY IN WAR AND

PEACE. The atomic bombs dropped by theUnited States on Japan in 1945 were fissionbombs. These were the creation of a group of sci-entists—legendary figures such as Americanphysicist J. Robert Oppenheimer (1904-1967),American mathematician John von Neumann(1903-1957), American physicist Edward Teller(1908-), and Italian physicist Enrico Fermi(1901-1954)—involved in the Manhattan Projectat Las Alamos, New Mexico.

Some of these geniuses, particularly Oppen-heimer, were ambivalent about the moral impli-cations of the enormous destructive power theycreated. However, most military historiansbelieve that far more lives—both Japanese andAmerican—would have been lost if America hadbeen forced to conduct a land invasion of Japan.As it was, the Japanese surrendered shortly afterthe cities of Hiroshima and Nagasaki suffered thedevastating effects of fission-based explosions.

By 1952, U.S. scientists had developed a“hydrogen,” or fusion bomb, thus raising thestakes greatly. This was a bomb that possessed farmore destructive capability than the onesdropped over Japan. Fortunately, the Hiroshimaand Nagasaki bombs were the only ones dropped

in wartime, and a ban on atmospheric nucleartesting has greatly reduced the chances of humanexposure to nuclear fallout of any kind. With theend of the arms race between the United Statesand the Soviet Union, the threat of nucleardestruction has receded somewhat—though itwill perhaps always be a part of human life.

Nonetheless, fear of nuclear power, spawnedas a result of the arms race, continues to cloudthe future of nuclear plants that generate elec-tricity—even though these, in fact, emit lessradioactive pollution than coal- or gas-burningpower plants. At the same time, scientists contin-ue to work on developing a process of power gen-eration by means of nuclear fusion, which, if andwhen it is achieved, will be one of the great mir-acles of science.

PARTICLE ACCELERATORS. Oneof the tools used by scientists researching nuclearfusion is the particle accelerator, which movesstreams of charged particles—protons, forinstance—faster and faster. These fast particlesare then aimed at a thin plate composed of a lightelement, such as lithium. If the proton managesto be “captured” in the nucleus of a lithium atom,the resulting nucleus is unstable, and breaks intoalpha particles.

This method of induced radioactivity isamong the most oft- used means of studyingnuclear structure and subatomic particles. In1932, the same year that Chadwick discoveredthe neutron, English physicist John D. Cockcroft(1897-1967) and Irish physicist Ernest Walton(1903-1995) built the first particle accelerator.Some particle accelerators today race the parti-cles in long straight lines or, to save space, inringed paths several miles in diameter.

Quantum Theory and Beyond

THE CONTRIBUTION OF RELA-

TIVITY. It may seem strange that in this lengthy(though, in fact, quite abbreviated!) overview ofdevelopments in understanding of the atom, nomention has been made of the figure most asso-ciated with the atom in the popular mind: Ger-man-American physicist Albert Einstein (1879-1955). The reasons for this are several. Einstein’srelativity theory addresses physical, rather thanchemical, processes, and did not directly con-tribute to enhanced understanding of atomicstructure or elements. The heart of relativity the-



Atomsory is the famous formula E = mc2, which meansthat every item of matter possesses energy pro-portional to its mass multiplied by the squaredspeed of light.

The value of mc2, of course, is an enormousamount of energy, and in order to be released insignificant quantities, an article of matter mustexperience the kinetic energy associated withvery, very high speeds—speeds close to that oflight. Obviously, the easiest thing to accelerate tosuch a speed is an atom, and hence, nuclear energy is a result of Einstein’s famous equa-tion. Nonetheless, it should be stressed thatalthough Einstein is associated with unlockingthe power of the atom, he did little to explainwhat atoms are.

However, in the course of developing his rel-ativity theory in 1905, Einstein put to rest a ques-tion about atoms and molecules that stillremained unsettled after more than a century.Einstein’s analysis of Brownian motion, com-bined with the confirmation of his results byFrench physicist Jean Baptiste Perrin (1870-1942), showed conclusively that yes, atoms andmolecules do exist. It may seem amazing that asrecently as 1905, this was still in doubt; however,this only serves to illustrate the arduous path sci-entists must tread in developing a theory thataccurately explains the world.

PLANCK’S QUANTUM THEORY.

A figure whose name deserves to be as much ahousehold word as Einstein’s—though it is not—is German physicist Max Planck (1858-1947). Itwas Planck who initiated the quantum theorythat Einstein developed further, a theory thatprevails today in the physical sciences.

At the atomic level, Planck showed, energy isemitted in tiny packets or “quanta.” Each of theseenergy packets is indivisible, and the behavior ofquanta redefine the old rules of physics handeddown from Newton and Maxwell. Thus, it isPlanck’s quantum theory, rather than Einstein’srelativity, that truly marks the watershed, or“before and after,” between classical physics andmodern physics.

Quantum theory is important not only tophysics, but to chemistry as well. It helps toexplain the energy levels of electrons, which arenot continuous, as in a spectrum, but jumpbetween certain discrete points. The quantummodel is now also applied to the overall behavior

of the electron; but before this could be fullyachieved, scientists had to develop a new under-standing of the way electrons move around thenucleus.

BOHR’S PLANETARY MODEL OF

THE ATOM. As was often the case in the his-tory of the atom, a man otherwise respected as agreat scientist put forward a theory of atomicstructure that at first seemed convincing, butultimately turned out to be inaccurate. In thiscase, it was Danish physicist Niels Bohr (1885-1962), a seminal figure in the development ofnuclear fission.

Using the observation, derived from quan-tum theory, that electrons only occupied specificenergy levels, Bohr hypothesized that electronsorbited around a nucleus in the same way thatplanets orbit the Sun. There is no reason tobelieve that Bohr formed this hypothesis for anysentimental reasons—though, of course, scien-tists are just as capable of prejudice as anyone.His work was based on his studies; nonetheless, itis easy to see how this model seemed appealing,showing as it did an order at the subatomic levelreflecting an order in the heavens.

ELECTRON CLOUDS. Many peopletoday who are not scientifically trained continueto think that an atom is structured much like theSolar System. This image is reinforced by sym-bolism, inherited from the 1950s, that represents“nuclear power” by showing a dot (the nucleus)surrounded by ovals at angles to one another,representing the orbital paths of electrons. How-ever, by the 1950s, this model of the atom hadalready been overturned.

In 1923, French physicist Louis de Broglie(1892-1987) introduced the particle-wavehypothesis, which indicated that electrons couldsometimes have the properties of waves—aneventuality not encompassed in the Bohr model.It became clear that though Bohr was correct inmaintaining that electrons occupy specific ener-gy levels, his planetary model was inadequate forexplaining the behavior of electrons.

Two years later, in 1925, German physicistWerner Heisenberg (1901-1976) introducedwhat came to be known as the HeisenbergUncertainty Principle, showing that the preciseposition and speed of an electron cannot beknown at the same time. Austrian physicistErwin Schrödinger (1887-1961) developed an



Atoms


equation for calculating how an electron with a

certain energy moves, identifying regions in an

atom where an electron possessing a certain

energy level is likely to be. Schrödinger’s equation

cannot, however, identify the location exactly.

Rather than being called orbits, which sug-

gest the orderly pattern of Bohr’s model,

Schrödinger’s regions of probability are called

orbitals. Moving within these orbitals, electrons

describe the shape of a cloud, as discussed much

earlier in this essay; as a result, the “electron

cloud” theory prevails today. This theory incor-

porates aspects of Bohr’s model, inasmuch as

electrons move from one orbital to another by

absorbing or emitting a quantum of energy.

WHERE TO LEARN MORE

“The Atom.” Thinkquest (Web site).<http://library.thinkquest.org/17940/texts/atom/atom.html> (May 18, 2001).

“Elements” (Web site).<http://home.school.net.hk/~chem/main/F5notes/atom/element.html> (May 18, 2001).

“Explore the Atom” CERN—European Organization forNuclear Research (Web site). <http://public.web.cern.ch/Public/SCIENCE/Welcome.html> (May 18, 2001).


ment that retains all the chemical and

physical properties of the element. An

atom can exist either alone or in combina-

tion with other atoms in a molecule. Atoms

are made up of protons, neutrons, and

electrons.







elements are listed on the periodic table of

elements in order of atomic number.




a large sample.

CHEMICAL SYMBOL: A one- or two-

letter abbreviation for the name of an

element.



atoms are usually joined in molecules.

ELECTRON: Negatively charged parti-

cles in an atom. Electrons, which spin



very small portion of the atom’s mass. The

number of electrons and protons is the

same, thus canceling out one another; on

the other hand, if an atom loses or gains

electrons, it becomes an ion.

ELEMENT SYMBOL: Another term for

chemical symbol.







neutrons.

MASS NUMBER: The sum of protons

and neutrons in an atom’s nucleus.

MOLECULE: A group of atoms, usually

(but not always) representing more

than one element, joined in a structure.

Compounds are typically made up of

molecules.

KEY TERMS


Atoms


Gallant, Roy A. The Ever-Changing Atom. New York:Benchmark Books, 1999.

Goldstein, Natalie. The Nature of the Atom. New York:Rosen Publishing Group, 2001.

“A Look Inside the Atom” (Web site). <http://www.aip.org/history/electron/jjhome.htm> (May 18, 2001).

“Portrait of the Atom” (Web site). <http://www.inetarena.com/~pdx4d/snelson/Portrait.html> (May18, 2001).

“A Science Odyssey: You Try It: Atom Builder.” PBS—Pub-

lic Broadcasting System (Web site). <http://www.pbs.

org/wgbh/aso/tryit/atom/> (May 18, 2001).

Spangenburg, Ray and Diane K. Moser. The History of

Science in the Nineteenth Century. New York: Facts on

File, 1994.






protons.

NUCLEON: A generic term for the

heavy particles—protons and neutrons—

that make up the nucleus of an atom.

NUCLEON NUMBER: Another term

for mass number.




PERIODIC TABLE OF ELEMENTS: A

chart that shows the elements arranged in

order of atomic number, along with chem-

ical symbol and the average atomic mass

(in atomic mass units) for that particular

element. Vertical columns within the peri-

odic table indicate groups or “families” of

elements with similar chemical character-

istics.






greater than that of an electron.

QUARK: A particle believed to be a com-

ponent of protons and neutrons. A quark

may either have a positive electric charge of

less than 1+, in which case it is called an

“up quark”; or a negative charge of less

than 1-, in which case it is called a “down

quark.”

RADIATION: In a general sense, radia-

tion can refer to anything that travels in a

stream, whether that stream be composed

of subatomic particles or electromagnetic

waves. In a more specific sense, the term

relates to the radiation from radioactive

materials, which can be harmful to human

beings.

RADIOACTIVITY: A term describing a

phenomenon whereby certain isotopes are

subject to a form of decay brought about

by the emission of high-energy particles or

radiation, such as alpha particles, beta par-

ticles, or gamma rays.

KEY TERMS CONTINUED



ATOMIC MASSAtomic Mass

CONCEPTEvery known item of matter in the universe has

some amount of mass, even if it is very small. But

what about something so insignificant in mass

that comparing it to a gram is like comparing a

millimeter to the distance between Earth and the

nearest galaxy? Obviously, special units are need-

ed for such measurements; then again, one might

ask why it is necessary to weigh atoms at all. One

answer is that everything is made of atoms. More

specifically, the work of a chemist requires the

use of accurate atomic proportions in forming

the molecules that make up a compound. The

measurement of atomic mass was thus a historic

challenge that had to be overcome, and the story

of the ways that scientists met this challenge is an

intriguing one.

HOW IT WORKS

Why Mass and Not Weight?

Some textbooks and other sources use the term

atomic weight instead of atomic mass. The first

of these is not as accurate as the second, which

explains why atomic mass was chosen as the sub-

ject of this essay. Indeed, the use of “atomic

weight” today merely reflects the fact that scien-

tists in the past used that expression and spoke of

“weighing” atoms. Though “weigh” is used as a

verb in this essay, this is only because it is less

cumbersome than “measure the mass of.” (In

addition, “atomic weight” may be mentioned

when discussing studies by scientists of the nine-

teenth century, who applied that term rather

than atomic mass.)

One might ask why such pains have beentaken to make the distinction. Mass is, after all,basically the same as weight, is it not? In fact it isnot, though people are accustomed to thinkingin those terms since most weight scales providemeasurements in both pounds and kilograms.However, the pound is a unit of weight in theEnglish system, whereas a kilogram is a unit ofmass in the metric and SI systems. Though thetwo are relatively convertible on Earth, they areactually quite different. (Of course it would makeno more sense to measure atoms in pounds orkilograms than to measure the width of a hair in light-years; but pounds and kilograms are the most familiar units of weight and massrespectively.)

Weight is a measure of force affected byEarth’s gravitational pull. Therefore a person’sweight varies according to gravity, and would bedifferent if measured on the Moon, whereas massis the same throughout the universe. Its invari-ability makes mass preferable to weight as aparameter of scientific measure.

Putting an Atom’s Size and Mass in Context

Mass does not necessarily relate to size, thoughthere is enough of a loose correlation that moreoften than not, we can say that an item of verysmall size will have very small mass. And atomsare very, very small—so much so that, until theearly twentieth century, chemists and physicistshad no accurate means of isolating them todetermine their mass.

The diameter of an atom is about 10-8 cm.This is equal to about 0.000000003937 in—or toput it another way, an inch is about as long as 250


Atomic Massmillion atoms lined up side by side. Obviously,special units are required for describing the sizeof atoms. Usually, measurements are provided interms of the angstrom, equal to 10-10 m. (In otherwords, there are 10 million angstroms in a mil-limeter.)

Measuring the spatial dimensions of anatom, however, is not nearly as important forchemists’ laboratory work as measuring its mass.The mass of an atom is almost inconceivablysmall. It takes about 5.0 • 1023 carbon atoms toequal just one gram in mass. At first, 1023 doesnot seem like such a huge number, until one considers that 106 is already a million, meaningthat 1023 is a million times a million times a mil-lion times 100,000. If 5.0 • 1023 angstromlengths—angstroms, not meters or even millime-ters—were laid end to end, they would stretchfrom Earth to the Sun and back 107,765 times!

Atomic Mass Units

It is obvious, then, that an entirely different unitis needed for measuring the mass of an atom, andfor this purpose, chemists and other scientistsuse an atom mass unit (abbreviated amu), whichis equal to 1.66 • 10-24 g.

Though the abbreviation amu is used in thisbook, atomic mass units are sometimes designat-ed simply by a u. On the other hand, they may bepresented as numbers without any unit of meas-ure—as for instance on the periodic table ofelements.

Within the context of biochemistry andmicrobiology, often the term dalton (abbreviatedDa or D) is used. This is useful for describing themass of large organic molecules, typically ren-dered in kilodaltons (kDa). The Latin prefix kilo-indicates 1,000 of something, and “kilodalton” ismuch less of a tongue-twister than “kilo-amu”.The term “dalton” honors English chemist JohnDalton (1766-1844), who, as we shall see, intro-duced the concept of the atom to science.

AVERAGE ATOMIC MASS. Since1960, when its value was standardized, the atom-ic mass unit has been officially known as the“unified atomic mass unit.” The addition of theword “unified” reflects the fact that atoms are notweighed individually—a labor that would beproblematic at the very least. In any case, to do sowould be to reinvent the wheel, as it were,


because average atomic mass figures have beenestablished for each element.

Average atomic mass figures range from1.008 amu for hydrogen, the first element listedon the periodic table of elements, to over 250amu for elements of very high atomic number.Figures for average atomic mass can be used todetermine the average mass of a molecule as well,since a molecule is just a group of atoms joinedin a structure. The mass of a molecule can bedetermined simply by adding together averageatomic mass figures for each atom the moleculecontains. A water molecule, for instance, consistsof two hydrogen atoms and one oxygen atom;therefore, its mass is equal to the average atomicmass of hydrogen multiplied by two, and addedto the average atomic mass of oxygen.

Avogadro’s Number and the Mole

Atomic mass units and average atomic mass arenot the only components necessary for obtainingaccurate mass figures where atoms are con-cerned. Obviously, as suggested several timesalready, it would be fruitless to determine themass of individual atoms or molecules. Norwould it do to measure the mass of a few hun-dred, or even a few million, of these particles.

ERNEST RUTHERFORD.


Atomic Mass becomes clear: as noted, carbon has an averageatomic mass of 12.01 amu, and multiplication ofthe average atomic mass by Avogadro’s numberyields a figure in grams equal to the value of theaverage atomic mass in atomic mass units.


Early Ideas of Atomic Mass

Dalton was not the first to put forth the idea ofthe atom: that concept, originated by the ancientGreeks, had been around for more than 2,000years. However, atomic theory had never takenhold in the world of science—or, at least, whatpassed for science prior to the seventeenth centu-ry revolution in thinking brought about byGalileo Galilei (1564-1642) and others.

Influenced by several distinguished prede-cessors, Dalton in 1803 formulated the theorythat nature is formed of tiny particles, an idea hepresented in A New System of Chemical Philoso-phy (1808). Dalton was the first to treat atoms asfully physical constructs; by contrast, ancientproponents of atomism conceived these funda-mental particles in ideal or spiritual terms. Dal-ton described atoms as hard, solid, indivisibleparticles with no inner spaces—a definition thatdid not endure, as later scientific inquiry revealedthe complexities of the atom. Yet he was correctin identifying atoms as having weight—or, as sci-entists say today, mass.

THE FIRST TABLE OF ATOMIC

WEIGHTS. The question was, how could any-one determine the weight of something as smallas an atom? A year after the publication of Dal-ton’s book, a discovery by French chemist andphysicist Joseph Gay-Lussac (1778-1850) andGerman naturalist Alexander von Humboldt(1769-1859) offered a clue. Humboldt and Gay-Lussac—famous for his gas law associating pres-sure and temperature—found that gases com-bine to form compounds in simple proportionsby volume.

For instance, as Humboldt and Gay-Lussacdiscovered, water is composed of only two ele-ments: hydrogen and oxygen, and these two com-bine in a whole-number ratio of 8:1. By separat-ing water into its components, they found thatfor every part of oxygen, there were eight parts ofhydrogen. Today we know that water molecules


After all, as we have seen, it takes about 500,000trillion million carbon atoms to equal just onegram—and a gram, after all, is still rather smallin mass compared to most objects encounteredin daily life. (There are 1,000 g in a kilogram, anda pound is equal to about 454 g.)

In addition, scientists need some means forcomparing atoms or molecules of different sub-stances in such a way that they know they areanalyzing equal numbers of particles. This can-not be done in terms of mass, because the num-ber of atoms in each sample varies: a gram ofhydrogen, for instance, contains about 12 timesas many atoms as a gram of carbon, which has anaverage atomic mass of 12.01 amu. What is need-ed, instead, is a way to designate a certain num-ber of atoms or molecules, such that accuratecomparisons are possible.

In order to do this, chemists make use of afigure known as Avogadro’s number. Named afterItalian physicist Amedeo Avogadro (1776-1856),it is equal to 6.022137 x 1023. Avogadro’s number,which is 6,022,137 followed by 17 zeroes, desig-nates the quantity of molecules in a mole (abbre-viated mol). The mole, a fundamental SI unit for“amount of substance,” is defined precisely as thenumber of carbon atoms in 12.01 g of carbon. Itis here that the value of Avogadro’s number

FREDERICK SODDY.


Atomic Massare formed by two hydrogen atoms, with an aver-age atomic mass of 1.008 amu each, and one oxy-gen atom. The ratio between the average atomicmass of oxygen (16.00 amu) and that of the twohydrogen atoms is indeed very nearly 8:1.

In the early nineteenth century, however,chemists had no concept of molecular structure,or any knowledge of the atomic masses of ele-ments. They could only go on guesswork: henceDalton, in preparing the world’s first “Table ofAtomic Weights,” had to make some assumptionsbased on Humboldt’s and Gay-Lussac’s findings.Presumably, Dalton reasoned, only one atom ofhydrogen combines with one atom of oxygen toform a “water atom.” He assigned to hydrogen aweight of 1, and according to this, calculated theweight of oxygen as 8.

AVOGADRO AND BERZELIUS

IMPROVE ON DALTON’S WORK. Theimplications of Gay-Lussac’s discovery that sub-stances combined in whole-number ratios wereastounding. (Gay-Lussac, who studied gases formuch of his career, is usually given more creditthan Humboldt, an explorer and botanist whohad his hand in many things.) On the one hand,the more scientists learned about nature, themore complex it seemed; yet here was somethingamazingly simple. Instead of combining in pro-portions of, say, 8.3907 to 1.4723, oxygen andhydrogen molecules formed a nice, clean, ratio of8 to 1. This served to illustrate the fact that, asDalton had stated, the fundamental particles ofmatter must be incredibly tiny; otherwise, itwould be impossible for every possible quantityof hydrogen and oxygen in water to have thesame ratio.

Intrigued by the work of Gay-Lussac, Avo-gadro in 1811 proposed that equal volumes ofgases have the same number of particles if meas-ured at the same temperature and pressure. Healso went on to address a problem raised by Dal-ton’s work. If atoms were indivisible, as Daltonhad indicated, how could oxygen exist both as itsown atom and also as part of a water “atom”?Water, as Avogadro correctly hypothesized, is notcomposed of atoms but of molecules, which arethemselves formed by the joining of two hydro-gen atoms with one oxygen atom.

Avogadro’s molecular theory opened theway to the clarification of atomic mass and thedevelopment of the mole, which, as we have seen,makes it possible to determine mass for large

quantities of molecules. However, his ideas didnot immediately gain acceptance. Only in 1860,four years after Avogadro’s death, did Italianchemist Stanislao Cannizzaro (1826-1910) resur-rect the concept of the molecule as a way ofaddressing disagreements among scientistsregarding the determination of atomic mass.

In the meantime, Swedish chemist JonsBerzelius (1779-1848) had adopted Dalton’smethod of comparing all “atomic weights” tothat of hydrogen. In 1828, Berzelius published atable of atomic weights, listing 54 elements alongwith their weight relative to that of hydrogen.Thus carbon, in Berzelius’s system, had a weightof 12. Unlike Dalton’s figures, Berzelius’s are veryclose to those used by scientists today. By thetime Russian chemist Dmitri IvanovitchMendeleev (1834-1907) created his periodictable in 1869, there were 63 known elements.That first table retained the system of measuringatomic mass in comparison to hydrogen.

The Discovery of SubatomicStructures

Until scientists began to discover the existence ofsubatomic structures, measurements of atomicmass could not really progress. Then in 1897,English physicist J. J. Thomson (1856-1940)identified the electron. A particle possessing neg-ative charge, the electron contributes little to anatom’s mass, but it pointed the way to the exis-tence of other particles within an atom. First ofall, there had to be a positive charge to offset thatof the electron, and secondly, the item or itemsproviding this positive charge had to account forthe majority of the atom’s mass.

Early in the twentieth century, Thomson’sstudent Ernest Rutherford (1871-1937) discov-ered that the atom has a nucleus, a center aroundwhich electrons move, and that the nucleus con-tains positively charged particles called protons.Protons have a mass 1,836 times as great as thatof an electron, and thus seemed to account forthe total atomic mass. Later, however, Rutherfordand English chemist Frederick Soddy (1877-1956) discovered that when an atom emitted cer-tain types of particles, its atomic mass changed.

ISOTOPES AND ATOMIC MASS.

Rutherford and Soddy named these atoms of dif-fering mass isotopes, though at that point—because the neutron had yet to be discovered—



Atomic Mass they did not know exactly what had caused thechange in mass. Certain types of isotopes, Soddyand Rutherford concluded, had a tendency todecay, moving (sometimes over a great period oftime) toward stabilization. Such isotopes wereradioactive.

Soddy concluded that atomic mass, as meas-ured by Berzelius, was actually an average of themass figures for all isotopes within that element.This explained a problem with Mendeleev’s peri-odic table, in which there seemed to be irregular-ities in the increase of atomic mass from elementto element. The answer to these variations inmass, it turned out, related to the number of iso-topes associated with a given element: the greaterthe number of isotopes, the more these affectedthe overall measure of the element’s mass.

A NEW DEFINITION OF ATOMIC

NUMBER. Up to this point, the term “atomicnumber” had a different, much less precise,meaning than it does today. As we have seen, theearly twentieth century periodic table listed ele-ments in order of their atomic mass in relation tohydrogen, and thus atomic number referred sim-ply to an element’s position in this ordering.Then, just a few years after Rutherford and Soddydiscovered isotopes, Welsh physicist HenryMoseley (1887-1915) determined that every ele-ment has a unique number of protons in itsnucleus.

Today, the number of protons in the nucle-us, rather than the mass of the atom, determinesthe atomic number of an element. Carbon, forinstance, has an atomic number of 6, not becausethere are five elements lighter—though this isalso true—but because it has six protons in itsnucleus. The ordering by atomic number hap-pens to correspond to the ordering by atomicmass, but atomic number provides a much moreprecise means of distinguishing elements. Forone thing, atomic number is always a whole inte-ger—1 for hydrogen, for instance, or 17 for chlo-rine, or 92 for uranium. Figures for mass, on theother hand, are almost always rendered with dec-imal fractions (for example, 1.008 for hydrogen).

NEUTRONS COMPLETE THE

PICTURE. As with many other discoveriesalong the way to uncovering the structure of theatom, Moseley’s identification of atomic numberwith the proton raised still more questions. Inparticular, if the unique number of protons iden-tified an element, what was it that made isotopes

of the same element different from one another?Hydrogen, as it turned out, indeed had a massvery nearly equal to that of one proton—thusjustifying its designation as the basic unit ofatomic mass. Were it not for the isotope knownas deuterium, which has a mass nearly twice asgreat as that of hydrogen, the element wouldhave an atomic mass of exactly 1 amu.

A discovery by English physicist JamesChadwick (1891-1974) in 1932 finally explainedwhat made an isotope an isotope. It was Chad-wick who identified the neutron, a particle withno electric charge, which resides in the nucleusalongside the protons. In deuterium, which hasone proton, one neutron, and one electron, theelectron accounts for only 0.0272% of the totalmass—a negligible figure. The proton, on theother hand, makes up 49.9392% of the mass.Until the discovery of the neutron, there hadbeen no explanation of the other 50.0336% ofthe mass in an atom with just one proton andone electron.

Average Atomic Mass Today

Thanks to Chadwick’s discovery of the neutron,it became clear why deuterium weighs almosttwice as much as ordinary hydrogen. This in turnis the reason why a large sample of hydrogen,containing as it does a few molecules of deuteri-um here and there, does not have the same aver-age atomic mass as a proton. Today scientistsknow that there are literally thousands of iso-topes—many of them stable, but many more ofthem unstable or radioactive—for the 100-pluselements on the periodic table. Each isotope, ofcourse, has a slightly different atomic mass. Thisrealization has led to clarification of atomic massfigures.

One might ask how figures of atomic massare determined. In the past, as we have seen, itwas largely a matter of guesswork, but todaychemists and physicist use a highly sophisticatedinstrument called a mass spectrometer. First,atoms are vaporized, then changed to positivelycharged ions, or cations, by “knocking off” elec-trons. The cations are then passed through amagnetic field, and this causes them to bedeflected by specific amounts, depending on thesize of the charge and its atomic mass. The parti-cles eventually wind up on a deflector plate,where the amount of deflection can be measuredand compared with the charge. Since 1 amu has



Atomic Massbeen calculated to equal approximately 931.494MeV, or mega electron-volts, very accurate fig-ures can be determined.

CALIBRATION OF THE ATOMIC

MASS UNIT. When 1 is divided by Avo-gadro’s number, the result is 1.66 • 10-24—thevalue, in grams, of 1 amu. However, in accor-dance with a 1960 agreement among members ofthe international scientific community, measure-ments of atomic mass take as their referencepoint the mass of carbon-12. Not only is the car-bon-12 isotope found in all living things, buthydrogen is a problematic standard because itbonds so readily with other elements. Accordingto the 1960 agreement, 1 amu is officially 1/12the mass of a carbon-12 atom, whose exact value(retested in 1998), is 1.6653873 • 10-24 g.

Carbon-12, sometimes represented as �162�C,

contains six protons and six neutrons. (Asexplained in the essay on Isotopes, where an iso-tope is indicated, the number to the upper left ofthe chemical symbol indicates the total numberof protons and neutrons. Sometimes this is theonly number shown; but if a number is includedon the lower left, this indicates only the numberof protons, which remains the same for each ele-ment.) The value of 1 amu thus obtained is, ineffect, an average of the mass for a proton andneutron—a usable figure, given the fact that aneutron weighs only 0.163% more than a proton.

Of all the carbon found in nature (asopposed to radioactive isotopes created in labo-ratories), 98.89% of it is carbon-12. The remain-der is mostly carbon-13, with traces of carbon-14, an unstable isotope produced in nature. Bydefinition, carbon-12 has an atomic mass ofexactly 12 amu; that of carbon-13 (about 1.11%of all carbon) is 13 amu. Thus the atomic mass ofcarbon, listed on the periodic table as 12.01 amu,is obtained by taking 98.89% of the mass of car-bon-12, combined with 1.11% of the mass ofcarbon-13.

ATOMIC MASS UNITS AND THE

PERIODIC TABLE. The periodic table as itis used today includes figures, in atomic massunits, for the average mass of each atom. As itturns out, Berzelius was not so far off in his useof hydrogen as a standard, since its mass is almostexactly 1 amu—but not quite, because (as notedabove) deuterium increases the average masssomewhat. Figures increase from there along the

periodic table, though not by a regular pattern.Sometimes the increase from one element to thenext is by just over 1 amu, and in other cases, theincrease is by more than 3 amu. This only servesto prove that atomic number, rather than atomicmass, is a more straightforward means of order-ing the elements.

Mass figures for many elements that tend toappear in the form of radioactive isotopes areusually shown in parentheses. This is particular-ly true for elements with atomic numbers above92, because samples of these elements do not stayaround long enough to be measured. Some havea half-life—the period in which half the isotopesdecay to a stable form—of just a few minutes,and for others, the half-life is but a fraction of asecond. Therefore, atomic mass figures representthe mass of the longest-lived isotope.

Uses of Atomic Mass in Chemistry

MOLAR MASS. Just as the value ofatomic mass units has been calibrated to themass of carbon-12, the mole is no longer official-ly defined in terms of Avogadro’s number,though in general its value has not changed. Byinternational scientific agreement, the moleequals the number of carbon atoms in 12.01 g ofcarbon. Note that, as stated earlier, carbon has anaverage atomic mass of 12.01 amu.

This is no coincidence, of course: multiplica-tion of the average atomic mass by Avogadro’snumber yields a figure in grams equal to thevalue of the average atomic mass in amu. A moleof helium, with an average atomic mass of 4.003,is 4.003 g. Iron, on the other hand, has an averageatomic mass of 55.85, so a mole of iron is 55.85g. These figures represent the molar mass—themass of 1 mole—for each of the elements men-tioned.

THE NEED FOR EXACT PRO-

PORTIONS. When chemists discover newsubstances in nature or create new ones in thelaboratory, the first thing they need to determineis the chemical formula—in other words, theexact quantities and proportions of elements ineach molecule. By chemical means, they separatethe compound into its constituent elements, thendetermine how much of each element is present.



Atomic Mass


Since they are using samples in relativelylarge quantities, molar mass figures for each ele-ment make it possible to determine the chemicalcomposition. To use a very simple example, sup-pose a quantity of water is separated, and theresult is 2.016 g of hydrogen and 16 g of oxygen.The latter is the molar mass of oxygen, and theformer is the molar mass of hydrogen multipliedby two. Thus we know that there are two moles ofhydrogen and one mole of oxygen, which com-bine to make one mole of water.

Of course the calculations used by chemistsworking in the research laboratories of universi-ties, government institutions, and corporationsare much, much more complex than the examplewe have given. In any case, it is critical that achemist be exact in making these determinations,

so as to know the amount of reactants needed to

produce a given amount of product, or the

amount of product that can be produced from a

given amount of reactant.

When a company produces millions or bil-

lions of a single item in a given year, a savings of

very small quantities in materials—thanks to

proper chemical measurement—can result in a

savings of billions of dollars on the bottom line.

Proper chemical measurement can also save lives.

Again, to use a very simple example, if a mole of

compounds weighs 44.01 g and is found to con-

tain two moles of oxygen and one of carbon, then

it is merely carbon dioxide—a compound essen-

tial to plant life. But if it weighs 28.01 g and has

one mole of oxygen with one mole of carbon, it

is poisonous carbon monoxide.



physical properties of that element.









ATOMIC WEIGHT: An old term for

atomic mass. Since weight varies depend-

ing on gravitational field, whereas mass is

the same throughout the universe, scien-

tists typically use the term “atomic mass”

instead.




a large sample.



gadro (1776-1856), equal to 6.022137 x


number of atoms or molecules in a mole.




DALTON: An alternate term for atomic

mass units, used in biochemistry and

microbiology for describing the mass of

large organic molecules. The dalton

(abbreviated Da or D) is named after

English chemist John Dalton (1766-1844),

who introduced the concept of the atom to

science.

ELEMENT: A substance made up of only

one kind of atom, which cannot be chemi-

cally broken into other substances.

HALF-LIFE: The length of time it takes a

substance to diminish to one-half its initial

amount.

KEY TERMS


Atomic Mass


WHERE TO LEARN MORE

“Atomic Weight” (Web site). <http://www.colorado.edu/physics/2000/periodic_table/atomic_weight.html>(May 23, 2001).

“An Experiment with ‘Atomic Mass’” (Web site).<http://www.carlton.paschools.pa.sk.ca/chemical/molemass/moles3a.htm> (May 23, 2001).

Knapp, Brian J. and David Woodroffe. The Periodic Table.Danbury, CT: Grolier Educational, 1998.

Oxlade, Chris. Elements and Compounds. Chicago:Heinemann Library, 2001.

“Periodic Table: Atomic Mass.” ChemicalElements.com

(Web site). <http://www.chemicalelements.com/

show/mass.html> (May 23, 2001).

“Relative Atomic Mass” (Web site). <http://www.

chemsoc.org/viselements/pages/mass.html> (May 23,

2001).

“What Are Atomic Number and Atomic Weight?” (Web

site). <http://tis.eh.doe.gov/ohre/roadmap/achre/

intro_9_3.html> (May 23, 2001).

ION: An atom or group of atoms that has

lost or gained one or more electrons, and

thus has a net electric charge.

ISOTOPES: Atoms of the same element

(that is, they have the same number of pro-

tons) that differ in terms of mass. Isotopes

may be either stable or unstable. The latter

type, known as radioisotopes, are radioac-

tive.

MASS: The amount of matter an object

contains.





atomic mass of that substance. Thus car-





ly speaking, Avogadro’s number of atoms

or molecules; however, in the more precise

SI definition, a mole is equal to the number

of carbon atoms in 12.01 g of carbon.

MOLECULE: A group of atoms, usually,

but not always, representing more


Compounds are typically made of up

molecules.






element.


phenomenon whereby certain isotopes

known as radioisotopes are subject to a

form of decay brought about by the emis-

sion of high-energy particles. “Decay” does

not mean that the isotope “rots”; rather, it

decays to form another isotope, until even-

tually (though this may take a long time), it

becomes stable.

KEY TERMS CONTINUED



ELECTRONSElectrons

CONCEPTNo one can see an electron. Even an electronmicroscope, used for imaging the activities ofthese subatomic particles, does not offer aglimpse of an electron as one can look at anamoeba; instead, the microscope detects the pat-terns of electron deflection. In any case, a single-cell organism is gargantuan in comparison to anelectron. Even when compared to a proton or aneutron, particles at the center of an atom, elec-trons are minuscule, being slightly more than1/2000 the size of either. Yet the electron is thekey to understanding the chemical process ofbonding, and electron configurations clarify anumber of aspects of the periodic table that may,at first, seem confusing.

HOW IT WORKS

The Electron’s Place in the Atom

The atom is discussed in detail elsewhere in thisbook; here, the particulars of atomic structurewill be presented in an abbreviated form, so thatthe discussion of electrons may proceed. At thecenter of an atom, the smallest particle of an ele-ment, there is a nucleus, which contains protonsand neutrons. Protons are positive in charge,while neutrons exert no charge.

Moving around the nucleus are elec-trons, negatively charged particles whose mass is very small in comparison to the proton or neutron: 1/1836 and 1/1839 the mass of the proton and neutron respectively.The mass of an electron is 9.109389 • 10–33 g.Compare the number 9 to the number

1,000,000,000,000,000,000,000,000,000,000,000,and this gives some idea of the ratio between anelectron’s mass and a gram. The large number—1 followed by 33 zeroes—is many trillions oftimes longer than the age of the universe in seconds.

IONS AND CHEMICAL BONDS.

Electrons, though very small, are exceedinglypowerful, and they are critical to both physicaland chemical processes. The negative charge ofan electron, designated by the symbol 1- or -1, isenough to counteract the positive charge (1+ or+1) of a proton, even though the proton hasmuch greater mass. As a result, atoms (becausethey possess equal numbers of protons and elec-trons) have an electric charge of zero.

Sometimes an atom will release an electron,or a released electron will work its way into thestructure of another atom. In either case, anatom that formerly had no electric chargeacquires one, becoming an ion. In the first of theinstances described, the atom that has lost anelectron or electrons becomes a positivelycharged ion, or cation. In the second instance, anatom that gains an electron or electrons becomesa negatively charged ion, or anion.

One of the first forms of chemical bondingdiscovered was ionic bonding, in which electronsclearly play a role. In fact, electrons are critical tovirtually all forms of chemical bonding, whichrelates to the interactions between electrons ofdifferent atoms.

MOVEMENT OF THE ELECTRON

ABOUT THE NUCLEUS. If the size of thenucleus were compared to a grape, the edge ofthe atom itself would form a radius of about amile. Because it is much smaller than the nucle-


Electrons


us, an electron might be depicted in this scenarioas a speck of dust—but a speck of dust withincredible energy, which crosses the mile radiusof this enlarged atom at amazing speeds. In anitem of matter that has been frozen to absolutezero, an electron moves relatively little. On theother hand, it can attain speeds comparable tothat of the speed of light: 186,000 mi (299,339km) per second.

The electron does not move around thenucleus as a planet orbits the Sun—a model ofelectron behavior that was once accepted, butwhich has since been overturned. On the otherhand, as we shall see, the electron does not sim-ply “go where it pleases”: it acts in accordancewith complex patterns described by the quantumtheory of physics. Indeed, to some extent thebehavior of the electron is so apparently erraticthat the word “pattern” seems hardly to describeit. However, to understand the electron clearly,one has to set aside all ideas about how objectsbehave in the physical world.

Emerging Models of Electron Behavior

The idea that matter is composed of atoms orig-inated in ancient Greece, but did not take holduntil early in the nineteenth century, with the

atomic theory of English chemist John Dalton(1766-1844). In the years that followed, numer-ous figures—among them Russian chemistDmitri Mendeleev (1834-1907), father of theperiodic table of elements—contributed to theemerging understanding of the atomic model.

Yet Mendeleev, despite his awareness that themass of an atom differentiated one element fromanother, had no concept of subatomic particles.No one did: until very late in the nineteenth cen-tury, the atom might as well have been a hard-shelled ball of matter, for all that scientists under-stood about its internal structure. The electronwas the first subatomic particle discovered, in1897—nearly a century after the scientific begin-ning of atomic theory. Yet the first hints regard-ing its existence had begun to appear some 60years before.

DISCOVERY OF THE ELEC-

TRON. In 1838, British physicist and chemistMichael Faraday (1791-1867) was working with aset of electrodes—metal plates used to emit orcollect electric charge—which he had placed ateither end of an evacuated glass tube. (In otherwords, most of the air and other matter had beenremoved from the tube.) He applied a charge ofseveral thousand volts between the electrodes,and discovered that an electric current flowedbetween them.

2py

x

y

z

2px

x

y

z

2p2

x

y

z

TWO-LOBED ORBITALS.


Electrons


This seemed to suggest the existence of par-ticles carrying an electric charge, and fourdecades later, English physicist William Crookes(1832-1919) expanded on these findings with hisexperiments using an apparatus that came to beknown as a Crookes tube. As with Faraday’sdevice, the Crookes tube used an evacuated glasstube encased between two electrodes—a cathodeat the negatively charged end, and an anode atthe positively charged end. A wire led outside thebulb to an electric source, and when electricitywas applied to the electrodes, the cathodes emit-ted rays. Crookes concluded that the cathode rayswere particles with a negative electric charge thatcame from the metal in the cathode plate.

English physicist J. J. Thomson (1856-1940)hypothesized that the negatively charged parti-cles Crookes had observed were being emitted byatoms, and in 1897 he gave a name to these par-ticles: electrons. The discovery of the electronraised questions concerning its place in the atom:obviously, there had to be a counterbalancingpositive charge, and if so, from whence did itcome?

FROM PLUM PUDDINGS TO

PLANETS. Around the beginning of the

twentieth century, the prevailing explanation ofatomic structure was the “plum pudding model,”which depicted electrons as floating like raisins ina “pudding” of positive charges. This was over-turned by the discovery of the nucleus, and,subsequently, of the proton and neutron it contained.

In 1911, the great Danish physicist NielsBohr (1885-1962) studied hydrogen atoms, andconcluded that electrons move around the nucle-us in much the same way that planets movearound the Sun. This worked well when describ-ing the behavior of hydrogen, which, in its sim-plest form—the isotope protium—has only oneelectron and one proton, without any neutrons.The model did not work as well when applied toother elements, however, and within less thantwo decades Bohr’s planetary model was over-turned.

As it turns out, the paths of an electron’smovement around the nucleus are nothing likethat of a planet’s orbit—except inasmuch as bothmodels describe a relatively small object movingaround a relatively large one. The reality is muchmore complex, and to comprehend the secret ofthe electron’s apparently random behavior is

A COMPUTER-GENERATED MODEL OF A NEON ATOM. THE NUCLEUS, AT CENTER, IS TOO SMALL TO BE SEEN AT THIS

SCALE AND IS REPRESENTED BY THE FLASH OF LIGHT. SURROUNDING THE NUCLEUS ARE THE ATOM’S ELECTRON

ORBITALS: 1S (SMALL SPHERE), 2S (LARGE SPHERE), AND 2P (LOBED). (Photograph by Kenneth Eward/BioGrafx. National

Audubon Society Collection/Photo Researchers, Inc. Reproduced by permission.)


Electronstruly a mind-expanding experience. Yet Bohr isstill considered among the greatest scientists ofthe twentieth century: it was he, after all, whofirst explained the quantum behavior of elec-trons, examined below.


Quantum Theory and the Atom

Much of what scientists understand today aboutthe atom in general, and the electron in particu-lar, comes from the quantum theory introducedby German physicist Max Planck (1858-1947).Planck showed that, at the atomic level, energy isemitted in tiny packets, or “quanta.” Applyingthis idea to the electron, Bohr developed an ideaof the levels at which an electron moves aroundthe nucleus. Though his conclusions led him tothe erroneous planetary model, Bohr’s explana-tion of energy levels still prevails.

As has been suggested, the interactionbetween electrons and protons is electromagnet-ic, and electromagnetic energy is emitted in theform of radiation, or a stream of waves and par-ticles. The Sun, for instance, emits electromag-netic radiation along a broad spectrum thatincludes radio waves, infrared light, visible light,ultraviolet light, x rays, and gamma rays. Theseare listed in ascending order of their energy lev-els, and the energy emitted can be analyzed interms of wavelength and frequency: the shorterthe wavelength, the greater the frequency and thegreater the energy level.

When an atom is at its ordinary energy level,it is said to be in a ground state, but when itacquires excess energy, it is referred to as being inan excited state. It may release some of that ener-gy in the form of a photon, a particle of electro-magnetic radiation. The amount of energyinvolved can be analyzed in terms of the wave-lengths of light the atom emits in the form ofphotons, and such analysis reveals some surpris-ing things about the energy levels of atoms.

If one studies the photons emitted by anatom as it moves between a ground state and anexcited state, one discovers that it emits only cer-tain kinds of photons. From this, Bohr conclud-ed that the energy levels of an atom do not exist

on a continuum; rather, there are only certainenergy levels possible for an atom of a given ele-ment. The energy levels are therefore said to bequantized.

The Wave Mechanical Model

In everyday terms, quantization can be comparedto the way that a person moves up a set of stairs:by discrete steps. If one step directly followsanother, there is no step in between, nor is thereany gradual way of moving from step to step, asone would move up a ramp. The movement ofelectrons from one energy level to another is nota steady progression, like the movement of a per-son up a ramp; rather, it is a series of quantumsteps, like those a person makes when climbing aset of stairs.

The idea of quantization was ultimatelyapplied to describing the paths that an electronmakes around the nucleus, but this requiredsome clarification along the way. It had beenbelieved that an electron could move throughany point between the nucleus and the edge ofthe atom (again, like a ramp), but it later becameclear that the electrons could only move alongspecific energy levels. As we have seen, Bohrbelieved that these corresponded to the orbits ofplanets around the Sun; but this explanationwould be discarded in light of new ideas thatemerged in the 1920s.

During the early part of that decade,French physicist Louis de Broglie (1892-1987)and Austrian physicist Erwin Schrödinger(1887-1961) introduced what came to beknown as the wave mechanical model, alsoknown as the particle-wave hypothesis. Becauselight appeared to have the properties of bothparticles and waves, they reasoned, electrons(possessing electromagnetic energy as they did)might behave in the same fashion. In otherwords, electrons were not just particles: in somesense, they were waves as well.

UNDERSTANDING ORBITALS.

The wave mechanical model depicted the move-ment of electrons, not as smooth orbits, but asorbitals—regions in which there is the highestprobability that an electron will be found. Anorbital is nothing like the shape of a solar system,but, perhaps ironically, it can be compared to thephotographs astronomers have taken of galaxies.In most of these photographs, one sees an area of



Electrons intense light emitted by the stars in the center.Further from this high-energy region, the distri-bution of stars (and hence of light) becomesincreasingly less dense as one moves from thecenter of the galaxy to the edges.

Replace the center of the galaxy with thenucleus of the atom, and the stars with electrons,and this is an approximation of an orbital. Just asa galaxy looks like a cloud of stars, scientists usethe term electron cloud to describe the patternformed by orbitals. The positions of electronscannot be predicted; rather, it is only possible toassign probabilities as to where they will be. Nat-urally, they are most drawn to the positivecharges in the nucleus, and hence an orbitaldepicts a high-density region of probabilities atthe center—much like the very bright center of agalaxy.

The further away from the nucleus, the lessthe probability that an electron will be in thatposition. Hence in models of an orbital, the dotsare concentrated at the center, and become lessdense the further away from the nucleus they are.As befits the comparison to a cloud, the edge ofan orbital is fuzzy. Contrary to the earlier beliefthat an atom was a clearly defined little pellet ofmatter, there is no certainty regarding the exactedge of a given atom; rather, scientists define thesphere of the atom as the region encompassing90% of the total electron probability.

THE MYSTERIOUS ELECTRON.

As complex as this description of electron behav-ior may seem, one can rest assured that the reali-ty is infinitely more complex than this simplifiedexplanation suggests. Among the great mysteriesof the universe is the question of why an electronmoves as it does, or even exactly how it does so.Nor does a probability model give us any way ofknowing when an electron will occupy a particu-lar position.

In fact, as German physicist WernerHeisenberg (1901-1976) showed with whatcame to be known as the Heisenberg Uncertain-ty Principle, it is impossible to know both thespeed of an electron and its precise position atthe same time. This, of course, goes againstevery law of physics that prevailed until about1920, and in fact quantum theory offers anentirely different model of reality than the oneaccepted during the seventeenth, eighteenth,and nineteenth centuries.

Orbitals

Given the challenges involved in understandingelectron behavior, it is amazing just how muchscientists do know about electrons—particularlywhere energy levels are concerned. This, in turn,makes possible an understanding of the periodictable that would astound Mendeleev.

Every element has a specific configuration ofenergy levels that becomes increasingly complexas one moves along the periodic table. In thepresent context, these configurations will beexplained as simply as possible, but the reader isencouraged to consult a reliable chemistry text-book for a more detailed explanation.

PRINCIPAL ENERGY LEVELS

AND SUBLEVELS. The principal energylevel of an atom indicates a distance that an elec-tron may move away from the nucleus. This isdesignated by a whole-number integer, begin-ning with 1 and moving upward: the higher thenumber, the further the electron is from thenucleus, and hence the greater the energy in theatom. Each principal energy level is divided intosublevels corresponding to the number n of theprincipal energy level: thus principal energy level1 has one sublevel, principal energy level 2 hastwo, and so on.

The simplest imaginable atom, a hydrogenatom in a ground state, has an orbital designatedas 1s1. The s indicates that an electron at energylevel 1 can be located in a region described by asphere. As for the significance of the superscript1, this will be explained shortly.

Suppose the hydrogen atom is excitedenough to be elevated to principal energy level 2.Now there are two sublevels, 2s (for now we willdispense with the superscript 1) and 2p. A porbital is rather like the shape of a figure eight,with its center of gravity located on the nucleus,and thus unlike the s sublevel, p orbitals can havea specific directional orientation. Depending on whether it is oriented along an x-, y-, or z-axis, orbitals in sublevel p are designated as 2px,2py, or 2pz.

If the hydrogen atom is further excited, andtherefore raised to principal energy level 3, it nowhas three possible sublevels, designated as s, p,and d. Some of the d orbitals can be imagined astwo figure eights at right angles to one another,once again with their centers of gravity along thenucleus of the atom. Because of their more com-



Electronsplex shape, there are five possible spatial orienta-tions for orbitals at the d sublevel.

Even more complex is the model of an atomat principal level 4, with four sublevels—s,p, d,and f, which has a total of seven spatial orienta-tions. Obviously, things get very, very complex atincreased energy levels. The greater the energylevel, the further the electron can move from thenucleus, and hence the greater the possible num-ber of orbitals and corresponding shapes.

ELECTRON SPIN AND THE

PAULI EXCLUSION PRINCIPLE.

Every electron spins in one of two directions, andthese are indicated by the symbols ↑ and ↓.According to the Pauli exclusion principle,named after the Austrian-Swiss physicist Wolf-gang Pauli (1900-1958), no more than two elec-trons can occupy the same orbital, and those twoelectrons must have spins opposite one another.

This explains the use of the superscript 1,which indicates the number of electrons in agiven orbital. This number is never greater thantwo: hence, the electron configuration of heliumis written as 1s2. It is understood that these twoelectrons must be spinning in opposite direc-tions, but sometimes this is indicated by anorbital diagram showing both an upward- anddownward-pointing arrow in an orbital that hasbeen filled, or only an upward-pointing arrow inan orbital possessing just one electron.

Electron Configuration andthe Periodic Table

THE FIRST 18 ELEMENTS. As onemoves up the periodic table from atomic number1 (hydrogen) to 18 (argon), a regular patternemerges. The orbitals are filled in a neat progres-sion: from helium (atomic number 2) onward, allof principal level 1 is filled; beginning with beryl-lium (atomic number 4), sublevel 2s is filled;from neon (atomic number 10), sublevel 2p—and hence principal level 2 as a whole—is filled,and so on. The electron configuration for neon,thus, is written as 1s22s22p6.

Note that if one adds together all the super-script numbers, one obtains the atomic numberof neon. This is appropriate, of course, sinceatomic number is defined by the number of pro-tons, and an atom in a non-ionized state has anequal number of protons and electrons. In notic-ing the electron configurations of an element,pay close attention to the last or highest principal

energy level represented. These are the valenceelectrons, the ones involved in chemical bonding.By contrast, the core electrons, or the ones thatare at lower energy levels, play no role in thebonding of atoms.

SHIFTS IN ELECTRON CONFIG-

URATION PATTERNS. After argon, how-ever, as one moves to the element occupying thenineteenth position on the periodic table—potassium—the rules change. Argon has an elec-tron configuration of 1s22s22p63s23p6, and by thepattern established with the first 18 elements,potassium should begin filling principal level 3d.Instead, it “skips” 3d and moves on to 4s. Theelement following argon, calcium, adds a secondelectron to the 4s level.

After calcium, the pattern again changes.Scandium (atomic number 21) is the first of thetransition metals, a group of elements on theperiodic table in which the 3d orbitals are filled.This explains why the transition metals are indi-cated by a shading separating them from the restof the elements on the periodic table.

But what about the two rows at the very bot-tom of the chart, representing groups of elementsthat are completely set apart from the periodictable? These are the lanthanide and actinideseries, which are the only elements that involve fsublevels. In the lanthanide series, the seven 4forbitals are filled, while the actinide series reflectsthe filling of the seven 5f orbitals.

As noted earlier, the patterns involved in thef sublevel are ultra-complex. Thus it is not sur-prising than the members of the lanthanideseries, with their intricately configured valenceelectrons, were very difficult to extract from oneanother, and from other elements: hence theirold designation as the “rare earth metals.” How-ever, there are a number of other factors—relat-ing to electrons, if not necessarily electron con-figuration—that explain why one element bondsas it does to another.

CHANGES IN ATOMIC SIZE. Withthe tools provided by the basic discussion of elec-trons presented in this essay, the reader is encour-aged to consult the essays on Chemical Bonding,as well as The Periodic Table of the Elements,both of which explore the consequences of elec-tron arrangement in chemistry. Not only areelectrons the key to chemical bonding, under-standing their configurations is critical to anunderstanding of the periodic table.



Electrons


One of the curious things about the period-ic table, for instance, is the fact that the sizes ofatoms decrease as one moves from left to rightacross a row or period, even though the sizesincrease as one moves from top to bottom alonga column or group. The latter fact—the increaseof atomic size in a group, as a function of increas-ing atomic number—is easy enough to explain:the higher the atomic number, the higher theprincipal energy level, and the greater the dis-tance from the nucleus to the furthest probabili-ty range.

On the other hand, the decrease in sizeacross a period (row) is a bit more challenging tocomprehend. However, all the elements in a peri-od have their outermost electrons at a particularprincipal energy level corresponding to the num-

ber of the period. For instance, the elements on

period 5 all have principal energy levels 1

through 5. Yet as one moves along a period from

left to right, there is a corresponding increase in

the number of protons within the nucleus. This

means a stronger positive charge pulling the elec-

trons inward; therefore, the “electron cloud” is

drawn ever closer toward the increasingly power-

ful charge at the center of the atom.

WHERE TO LEARN MORE

“Chemical Bond” (Web site). <http://www.science.uwaterloo.ca/~cchieh/cact/c120/chembond.html>(May 18, 2001).

“The Discovery of the Electron” (Web site).<http://www.aip.org/history/electron/> (May 18,2001).

ANION: The negative ion that results

when an atom gains one or more electrons.

ANODE: An electrode at the positively

charged end of a supply of electric current.

ATOM: The smallest particle of an

element.






CATHODE: An electrode at the negative-

ly charged end of a supply of electric cur-

rent.

CATION: The positive ion that results

when an atom loses one or more electrons.

ELECTRODE: A structure, often a metal

plate or grid, that conducts electricity, and

which is used to emit or collect electric

charge.

ELECTRON: A negatively charged parti-

cle in an atom.

ELECTRON CLOUD: A term used to

describe the pattern formed by orbitals.

EXCITED STATE: A term describing the

characteristics of an atom that has

acquired excess energy.

GROUND STATE: A term describing

the state of an atom at its ordinary energy

level.







neutrons.



in the nucleus of an atom, alongside

protons.




KEY TERMS


Electrons


Ebbing, Darrell D.; R. A. D. Wentworth; and James P.

Birk. Introductory Chemistry. Boston: Houghton Mif-

flin, 1995.

Gallant, Roy A. The Ever-Changing Atom. New York:

Benchmark Books, 1999.

Goldstein, Natalie. The Nature of the Atom. New York:

Rosen Publishing Group, 2001.

“Life, the Universe, and the Electron” (Web site).

<http://www.iop.org/Physics/Electron/Exhibition/

(May 18, 2001).


“Valence Shell Electron Pair Repulsion (VSEPR).”<http://www.shef.ac.uk/~chem/vsepr/chime/vsepr.html> (May 18, 2001).

“What Are Electron Microscopes?” (Web site).<http://www.unl.edu/CMRAcfem/em.htm> (May 18,2001).


ORBITAL: A pattern of probabilities

regarding the regions that an electron can

occupy within an atom in a particular

energy state. The orbital, complex and

imprecise as it may seem, is a much more

accurate depiction of electron behavior

than the model once used, which depicted

electrons moving in precisely defined

orbits around the nucleus, rather as planets

move around the Sun.



order of atomic number. Vertical columns

within the periodic table indicate groups

or “families” of elements with similar

chemical characteristics.

PHOTON: A particle of electromagnetic

radiation.

PRINCIPAL ENERGY LEVEL: A value

indicating the distance that an electron

may move away from the nucleus of an

atom. This is designated by a whole-num-

ber integer, beginning with 1 and moving

upward. The higher the number, the fur-

ther the electron is from the nucleus, and

hence the greater the energy in the atom.


in an atom.

QUANTIZATION: A term describing any

property that has only certain discrete val-

ues, as opposed to values distributed along

a continuum. The quantization of an atom

means that it does not have a continuous

range of energy levels; rather, it can exist

only at certain levels of energy from the

ground state through various excited

states.





waves.

KEY TERMS CONTINUED



ISOTOPESIsotopes

CONCEPTIsotopes are atoms of the same element that havedifferent masses due to differences in the numberof neutrons they contain. Many isotopes are sta-ble, meaning that they are not subject to radioac-tive decay, but many more are radioactive. Thelatter, also known as radioisotopes, play a signifi-cant role in modern life. Carbon-14, for instance,is used for estimating the age of objects within arelatively recent span of time—up to about 5,000years—whereas geologists and other scientistsuse uranium-238 to date minerals of an age on ascale with that of the Earth. Concerns overnuclear power and nuclear weapons testing in theatmosphere have heightened awareness of thedangers posed by certain kinds of radioactive iso-topes, which can indeed be hazardous to humanlife. However, the reality is that people are sub-jected to considerably more radiation from non-nuclear sources.

HOW IT WORKS

Atoms and Elements

The elements are substances that cannot be bro-ken down into other matter by chemical means,and an atom is the fundamental particle in anelement. As of 2001, there were 112 known ele-ments, 88 of which occur in nature; the rest werecreated in laboratories. Due to their high levels ofradioactivity, they exist only for extremely shortperiods of time. Whatever the number of ele-ments—and obviously that number will increaseover time, as new elements are synthesized—thesame number of basic atomic structures exists inthe universe.

What distinguishes one element fromanother is the number of protons, subatomicparticles with a positive electric charge, in thenucleus, or center, of the atom. The number ofprotons, whatever it may be, is unique to an ele-ment. Thus if an atom has one proton, it is anatom of hydrogen, because hydrogen has anatomic number of 1, as shown on the periodictable of elements. If an atom has 109 protons, onthe other hand, it is meitnerium. (Meitnerium,synthesized at a German laboratory in 1982, isthe last element on the periodic table to havebeen assigned a name as of 2001.)

THE NUCLEUS AND ELEC-

TRONS. Together with protons in the nucleusare neutrons, which exert no charge. The discovery of these particles, integral to the for-mation of isotopes, is discussed below. Thenucleus, with a diameter about 1/10,000 that ofthe atom itself, makes up only a tiny portion ofthe atom’s volume, but the vast majority of itsmass. Thus a change in the mass of a nucleus, asoccurs when an isotope is formed, is reflected bya noticeable change in the mass of the atom itself.

Far from the nucleus (in relative terms, ofcourse), at the perimeter of the atom, are theelectrons, which have a negative electric charge.Whereas the protons and neutrons have aboutthe same mass, the mass of an electron is lessthan 0.06% of either a proton or neutron.Nonetheless, electrons play a highly significantrole in chemical reactions and chemical bonding.Just as isotopes are the result of changes in thenumber of neutrons, ions—atoms that are eitherpositive or negative in electric charge—are theresult of changes in the number of electrons.


IsotopesUnless it loses or gains an electron, thusbecoming an ion, an atom is neutral in charge,and it maintains this electric-charge neutrality byhaving an equal number of protons and elec-trons. There is, however, no law of the universestating that an atom must have the same numberof neutrons as it does protons and electrons:some do, but this is far from universal, as we shall see.

Neutrons

The number of neutrons is variable within anelement precisely because they exert no charge,and thus while their addition or removal changesthe mass, it does not affect the electric charge ofthe atom. Therefore, whereas the importance ofthe proton and the electron is very clear to any-one who studies atomic behavior, neutrons, onthe other hand, might seem at first glance asthough they are only “along for the ride.” Yet theyare all-important to the formation of isotopes.

Not surprisingly, given their lack of electriccharge, neutrons were the last of the three majorsubatomic particles to be discovered. Englishphysicist J. J. Thomson (1856-1940) identifiedthe electron in 1897, and another English physi-cist, Ernest Rutherford (1871-1937), discoveredthe proton in 1914. Rutherford’s discovery over-turned the old “plum pudding” model, wherebyatoms were depicted as consisting of electronsfloating in a positively charged cloud, rather likeraisins in an English plum pudding. As Ruther-ford showed, the atom must have a nucleus—yetprotons alone could not account for the mass ofthe nucleus.

There must be something else at the heart ofthe atom, and in 1932, yet another English physi-cist, James Chadwick (1891-1974), identifiedwhat it was. Working with radioactive material,he found that a certain type of subatomic parti-cle could penetrate lead. All types of radiationknown at the time were stopped by the lead, andtherefore Chadwick reasoned that this particlemust be neutral in charge. In 1932, he won theNobel Prize in physics for his discovery of theneutron.

NEUTRONS AND NUCLEAR

FUSION. Neutrons played a critical role in thedevelopment of the atomic bomb during the1940s. In nuclear fission, atoms of uranium arebombarded with neutrons. The result is that theuranium nucleus splits in half, releasing huge


amounts of energy. As it does so, it emits severalextra neutrons, which split more uranium nuclei,creating still more energy and setting off a chainreaction.

This explains the destructive power in anatomic bomb, as well as the constructivepower—providing energy to homes and busi-nesses—in a nuclear power plant. Whereas thechain reaction in an atomic bomb becomes anuncontrolled explosion, in a nuclear plant, thereaction is slowed and controlled. One of themeans used to do this is by the application of“heavy water,” which, as we shall see, is watermade with a hydrogen isotope.

Isotopes: The Basics

Two atoms may have the same number of pro-tons, and thus be of the same element, yet differin their number of neutrons. Such atoms arecalled isotopes, atoms of the same element hav-ing different masses. The name comes from theGreek phrase isos topos, meaning “same place”:because they have the same atomic number,

A MUSHROOM CLOUD RISES AFTER THE DETONATION OF

A HYDROGEN BOMB BY FRANCE IN A 1968 TEST. DEU-TERIUM IS A CRUCIAL PART OF THE DETONATING DEVICE

FOR HYDROGEN BOMBS. (Bettmann/Corbis. Reproduced by per-

mission.)


Isotopes


isotopes of the same element occupy the sameposition on the periodic table.

Also called nuclides, isotopes are representedsymbolically as follows: �

ma�S, where S is the sym-

bol of the element, a is the atomic number, and mis the mass number—the sum of protons andneutrons in the atom’s nucleus. For the stable sil-ver isotope designated as �

94

37�Ag, for instance, Ag is

the element symbol; 47 its atomic number; and93 the mass number. From this, it is easy to dis-cern that this particular stable isotope has 46neutrons in its nucleus.

Because the atomic number of any elementis established, sometimes isotopes are represent-ed simply with the mass number, thus: 93Ag. Theymay also be designated with a subscript notationindicating the number of neutrons, so that this

information can be obtained at a glance withouthaving to do the arithmetic. For the silver isotopeshown here, this is written as �

94

37�Ag46. Isotopes can

also be indicated by simple nomenclature: forinstance, carbon-12 or carbon-13.

Stable and Unstable Isotopes

Radioactivity is a term describing a phenomenonwhereby certain materials are subject to a form ofdecay brought about by the emission of high-energy particles or radiation. Forms of particlesor energy emitted in radiation include alpha par-ticles (positively charged helium nuclei); betaparticles (either electrons or subatomic particlescalled positrons); or gamma rays, which occupythe highest energy level in the electromagnetic

ENRICO FERMI.


Isotopesradiation emitted by the Sun. Radioactivity willbe discussed below, but for the present, the prin-cipal concern is with radioactive properties as adistinguishing factor between the two varieties ofisotope.

Isotopes are either stable or unstable. Theunstable variety, known as radioisotopes, aresubject to radioactive decay, but in this context,“decay” does not mean what it usually does. Aradioisotope does not “rot”; it decays by turninginto another isotope of the same element—oreven into another element entirely. (For example,uranium-238 decays by emitting alpha particles,ultimately becoming lead-206.) A stable isotope,on the other hand, has already become what it isgoing to be, and will not experience furtherdecay.

Most elements have between two and six sta-ble isotopes. On the other hand, a few ele-ments—for example, technetium—have no sta-ble isotopes. Twenty elements, among them gold,fluorine, sodium, aluminum, and phosphorus,have only one stable isotope each. The elementwith the most stable isotopes is easy to rememberbecause its name is almost the same as its num-ber of stable isotopes: tin, with 10.

As for unstable isotopes, there are over1,000, some of which exist in nature, but most ofwhich have been created synthetically in labora-tories. This number is not fixed; in any case, it isnot necessarily important, because many of thesehighly radioactive isotopes last only for fractionsof a second before decaying to form a stable iso-tope. Yet radioisotopes in general have so manyuses, in comparison to stable isotopes, that theyare often referred to simply as “isotopes.”

Understanding Isotopes

Before proceeding with a discussion of isotopesand their uses, it is necessary to address a pointraised earlier, when it was stated that some atomsdo have the same numbers of neutrons and pro-tons, but that this is far from universal. In fact,nuclear stability is in part a function of neutron-to-proton ratio.

Stable nuclei with low atomic numbers (upto about 20) have approximately the same num-ber of neutrons and protons. For example, themost stable and abundant form of carbon is car-bon-12, with six protons and six neutrons.Beyond atomic number 20 or so, however, thenumber of neutrons begins to grow: in other

words, the lowest mass number is increasinglyhigh in comparison to the atomic number.

For example, uranium has an atomic num-ber of 92, but the lowest mass number for a ura-nium isotope is not 184, or 92 multiplied by two;it is 218. The ratio of neutrons to protons neces-sary for a stable isotope creeps upward along theperiodic table: tin, with an atomic number of 50,has a stable isotope with a mass number of 120,indicating a 1.4 to 1 ratio of neutrons to protons.For mercury-200, the ratio is 1.5 to 1.

The higher the atomic number, by defini-tion, the greater the number of protons in thenucleus. This means that more neutrons arerequired to “bind” the nucleus together. In fact,all nuclei with 84 protons or more (i.e., startingat polonium and moving along the periodictable) are radioactive, for the simple reason thatit is increasingly difficult for the neutrons towithstand the strain of keeping so many protonsin place.

One can predict the mode of radioactivedecay by noting whether the nucleus is neutron-rich or neutron-poor. Neutron-rich nucleiundergo beta emission, which decreases thenumbers of protons in the nucleus. Neutron-poor nuclei typically undergo positron emissionor electron capture, the first of these being moreprevalent among the lighter nuclei. Elementswith atomic numbers of 84 or greater generallyundergo alpha emission, which decreases thenumbers of protons and neutrons by two each.


Deuterium and Tritium

Only three isotopes are considered significantenough to have names of their own, as opposedto being named after a parent atom (for example,carbon-12, uranium-238). These are protium,deuterium, and tritium, all three isotopes ofhydrogen. Protium, or 1H, is hydrogen in its mostbasic form—one proton, no neutrons—and thename “protium” is only applied when necessaryto distinguish it from the other two isotopes.Therefore we will focus primarily on the two others.

Deuterium, designated as 2H, is a stable iso-tope, whereas tritium—3H—is radioactive. Both,in fact, have chemical symbols (D and T respec-



Isotopes tively), just as though they were elements on theperiodic table. What makes these two so special?They are, as it were, “the products of a goodhome”—in other words, their parent atom is themost basic and plentiful element in the universe.Indeed, the vast majority of the universe ishydrogen, along with helium, which is formed bythe fusion of hydrogen atoms. If all atoms werenumbers, then hydrogen would be 1; but ofcourse, this is more than a metaphor, since itsatomic number is indeed 1.

Ordinary hydrogen or protium, as noted,consists of a single proton and a single electron,the simplest possible atomic form possible. Itssimplicity has made it a model for understandingthe atom, and therefore when physicists discov-ered the existence of two hydrogen atoms thatwere just a bit more complex, they wereintrigued.

Just as hydrogen represented the standardagainst which atoms could be measured, scien-tists reasoned, deuterium and tritium could offervaluable information regarding stable and unsta-ble isotopes respectively. Furthermore, the pro-nounced tendency of hydrogen to bond withother substances—it almost never appears byitself on Earth—presented endless opportunitiesfor study regarding hydrogen isotopes in associa-tion with other elements.

ISOLATION OF DEUTERIUM.

Deuterium is sometimes called “heavy hydro-gen,” and its nucleus—with one proton and oneneutron—is called a deuteron. It was first isolat-ed in 1931 by American chemist Harold ClaytonUrey (1893-1981), who was awarded the 1934Nobel Prize in Chemistry for his discovery.

Serving at that time as a professor of chem-istry at Columbia University in New York City,Urey started with the assumption that any hydro-gen isotopes other than protium must exist invery minute quantities. This assumption, in turn,followed from an awareness that hydrogen’s aver-age atomic mass—measured in atomic massunits—was only slightly higher than 1. Theremust be, as Urey correctly reasoned, a very smallquantity of “heavy hydrogen” on Earth.

To separate deuterium, Urey collected a rela-tively large sample of liquid hydrogen: 4.2 quarts(4 l). Then he allowed the liquid to evaporatevery slowly, predicting that the more abundantprotium would evaporate more quickly than the

isotope whose existence he had hypothesized.After all but 0.034 oz (1 ml) of the sample hadevaporated, he submitted the remainder to aform of analysis called spectroscopy, adding aburst of energy to the atoms and then analyzingthe light spectrum they emitted for evidence ofdiffering varieties of atom.

CHARACTERISTICS AND USES

OF DEUTERIUM. Deuterium, with anatomic mass of 2.014102 amu, is almost exactlytwice as heavy as protium, which has an atomicmass of 1.007825. Its melting point, or the tem-perature at which it changes from a solid to a liq-uid -426°F (-254°C), is much higher than forprotium, which melts at -434°F (-259°C). Thesame relationship holds for its boiling point, orthe temperature at which it changes from a liquidto its normal state on Earth, as a gas: -417°F (-249°C), as compared to -423°F (-253°C) forprotium. Deuterium is also much, much lessplentiful than protium: protium represents99.985% of all the hydrogen that occurs natural-ly, meaning that deuterium accounts for just0.015%.

Often, deuterium is applied as a tracer, anatom or group of atoms whose participation in achemical, physical, or biological reaction can beeasily observed. Radioisotopes are most oftenused as tracers, precisely because of theirradioactive emissions; deuterium, on the otherhand, is effective due to its almost 2:1 mass ratioin comparison to protium. In addition, it bondswith other atoms in a fashion slightly differentfrom that of protium, and this contrast makes itspresence easier to trace.

Its higher boiling and melting points meanthat when deuterium is combined with oxygen toform “heavy water” (D2O), the water likewise hashigher boiling and melting points than ordinarywater. Heavy water is often used in nuclear fis-sion reactors to slow down the fission process, orthe splitting of atoms.

DEUTERIUM IN NUCLEAR FU-

SION. Deuterium is also applied in a type ofnuclear reaction much more powerful that fis-sion: fusion, or the joining of atomic nuclei. TheSun produces energy by fusion, a thermonuclearreaction that takes places at temperatures ofmany millions of degrees Celsius. In solar fusion,it appears that two protium nuclei join to form asingle deuteron.



IsotopesDuring the period shortly after World WarII, physicists developed a means of duplicatingthe thermonuclear fusion process. The result wasthe hydrogen bomb—more properly called afusion bomb—whose detonating device was acompound of lithium and deuterium called lithi-um deuteride. Vastly more powerful than the“atomic” (that is, fission) bombs dropped by theUnited States over Japan in 1945, the hydrogenbomb greatly increased the threat of worldwidenuclear annihilation in the postwar years.

Yet the power that could destroy the worldalso has the potential to provide safe, abundantfusion energy from power plants—a dream thatas yet remains unrealized. Among the approach-es being attempted by physicists studying nuclearfusion is a process in which two deuterons arefused. The result is a triton, the nucleus of tri-tium, along with a single proton. The triton anddeuteron would then be fused to create a heliumnucleus, with a resulting release of vast amountsof energy.

TRITIUM. Whereas deuterium has a sin-gle neutron, tritium—as its mass number of 3indicates—has two. And just as deuterium hasapproximately twice the mass of protium, tri-tium has about three times the mass, 3.016 amu.As is expected, the thermal properties of tritiumare different from those of protium. Again, themelting and boiling points are higher: thus tri-tium heavy water (T2O) melts at 40°F (4.5°C), ascompared with 32°F (0°C) for H2O.

Because it is radioactive, tritium is oftendescribed in terms of half-life, the length of timeit takes for a substance to diminish to one-half itsinitial amount. The half-life of tritium is 12.26years. As it decays, its nucleus emits a low-energybeta particle, and this results in the creation ofthe helium-3 isotope. Due to the low energy lev-els involved, the radioactive decay of tritiumposes little danger to humans.

Like deuterium, tritium is applied in nuclearfusion, though due to its scarcity, it is usuallycombined with deuterium. Furthermore, tritiumdecay requires that hydrogen bombs containingthe radioisotope be recharged periodically. Also,like deuterium, tritium is an effective tracer.Sometimes it is released in small quantities intogroundwater as a means of monitoring subter-ranean water flow. It is also used as a tracer inbiochemical processes.

Separating Isotopes

As noted in the discussion of deuterium, tritiumcan only be separated from protium due to thedifferences in mass. The chemical properties ofisotopes with the same parent element makethem otherwise indistinguishable, and hencepurely chemical means cannot be used to sepa-rate them.

Physicists working on the Manhattan Pro-ject, the U.S. effort to develop atomic weaponryduring World War II, were faced with the need toseparate 235U from 238U. Uranium-238 is far moreabundant, but what they wanted was the urani-um-235, highly fissionable and thus useful in theprocesses they were attempting.

Their solution was to allow a gaseous urani-um compound to diffuse, or separate, the urani-um through porous barriers. Because uranium-238 was heavier, it tended to move more slowlythrough the barriers, much like grains of rice get-ting caught in a sifter. Another means of separat-ing isotopes is by mass spectrometry.

Radioactivity

One of the scientists working on the ManhattanProject was Italian physicist Enrico Fermi (1901-1954), who used radium and beryllium powderto construct a neutron source for making newradioactive materials. Fermi and his associatessucceeded in producing radioisotopes of sodium,iron, copper, gold, and numerous other elements.As a result of Fermi’s work, for which he won the1938 Nobel Prize in Physics, scientists have beenable to develop radioactive versions of virtuallyall elements.

Interestingly, the ideas of radioactivity, fis-sion reactions, and fusion reactions collectivelyrepresent the realization of a goal sought by themedieval alchemists: the transformation of oneelement into another. The alchemists, forerun-ners of chemists, believed they could transformordinary metals into gold by using variouspotions—an impossible dream. Yet as noted inthe preceding paragraph, among the radioiso-topes generated by Fermi’s neutron source wasgold. The “catch,” of course, is that this gold wasunstable; furthermore, the amount of energy andhuman mental effort required to generate it faroutweighed the monetary value of the gold itself.

Radioactivity is, in the modern imagination,typically associated with fallout from nuclear



Isotopes war, or with hazards resulting from nuclearpower—hazards that, as it turns out, have beengreatly exaggerated. Nor is radioactivity alwaysharmful to humans. For instance, with its appli-cations in medicine—as a means of diagnosingand treating thyroid problems, or as a treatmentfor cancer patients—it can actually save lives.

HAZARDS ASSOCIATED WITH

RADIOACTIVITY. It is a good thing thatradiation, even the harmful variety known asionizing radiation, is not fatal in small doses,because every person on Earth is exposed tosmall quantities of radiation every year. About82% of this comes from natural sources, and18% from manmade sources. Of course, somepeople are at much greater risk of radiationexposure than others: coal miners are exposed tohigher levels of the radon-222 isotope presentunderground, while cigarette smokers ingestmuch higher levels of radiation than ordinarypeople, due to the polonium-210, lead-210, andradon-222 isotopes present in the nitrogen fertil-izers used to grow tobacco.

Nuclear weapons, as most people know, pro-duce a great deal of radioactive pollution. How-ever, atmospheric testing of nuclear armamentshas long been banned, and though the isotopesreleased in such tests are expected to remain inthe atmosphere for about a century, they do notconstitute a significant health hazard to mostAmericans. (It should be noted that nations notinclined to abide by international protocolsmight still conduct atmospheric tests in defianceof the test bans.) Nuclear power plants, despitethe great deal of attention they have receivedfrom the media and environmentalist groups, donot pose the hazard that has often been claimed:in fact, coal- and oil-burning power plants areresponsible for far more radioactive pollution inthe United States.

This is not to say that nuclear energy posesno dangers, as the disaster at Chernobyl in theformer Soviet Union has shown. In April 1986,an accident at a nuclear reactor in what is nowthe Ukraine killed 31 workers immediately, andultimately led to the deaths of some 10,000 peo-ple. The fact that the radiation was allowed tospread had much to do with the secretive tacticsof the Communist government, which attempted to cover up the problem rather than evacuate the area.

Another danger associated with nuclearpower plants is radioactive waste. Spent fuel rodsand other waste products from these plants haveto be dumped somewhere, but it cannot simplybe buried in the ground because it will create acontinuing health hazard through the water sup-ply. No fully fail-safe storage system has beendeveloped, and the problem of radioactive wasteposes a continuing threat due to the extremelylong half-lives of some of the isotopes involved.

Dating Techniques

In addition to their uses in applications related tonuclear energy, isotopes play a significant role indating techniques. The latter may sound like asubject that has something to do with romance,but it does not: dating techniques involve the useof materials, including isotopes, to estimate theage of both organic and inorganic materials.

Uranium-238, for instance, has a half-life of4.47 • 109 years, which is nearly the age of Earth;in fact, uranium-dating techniques have beenused to determine the planet’s age, which is esti-mated at about 4.7 billion years. As noted else-where in this volume, potassium-argon dating,which involves the isotopes potassium-40 andargon-40, has been used to date volcanic layers ineast Africa. Because the half-life of potassium-40is 1.3 billion years, this method is useful for dat-ing activities that are distant in the human scaleof time, but fairly recent in geological terms.

Another dating technique is radiocarbondating, used for estimating the age of things thatwere once alive. All living things contain carbon,both in the form of the stable isotope carbon-12and the radioisotope carbon-14. While a plant oranimal is living, there is a certain proportionbetween the amounts of these two isotopes in theorganism’s body, with carbon-12 being far moreabundant. When the organism dies, however, itceases to acquire new carbon, and the carbon-14present in the body begins to decay into nitro-gen-14. The amount of nitrogen-14 that has beenformed is thus an indication of the amount oftime that has passed since the organism was alive.

Because it has a half-life of 5,730 years, car-bon-14 is useful for dating activities within thespan of human history, though it is not withoutcontroversy. Some scientists contend, forinstance, that samples may be contaminated bycarbon from the surrounding soils, thus affectingratios and leading to inaccurate dates.



Isotopes



ment. Atoms are made up of protons, neu-

trons, and electrons. Atoms that have the

same number of protons—that is, are of

the same element—but differ in number of

neutrons are known as isotopes.












a large sample.


one kind of atom. Hence an element can-

not be chemically broken into other sub-

stances.


cles in an atom, which spin around the

protons and neutrons that make up the

atom’s nucleus.



amount.







neutrons. This results in a difference of

mass. Isotopes may be either stable or

unstable. The latter type is known as a

radioisotope.




protons.




The plural of “nucleus” is nuclei.

NUCLIDES: Another name for isotopes.












KEY TERMS


Isotopes


WHERE TO LEARN MORE

“Carbon 14 Dating Calculator” (Web site). <http://www.museum.mq.edu.au/eegypt2/carbdate.html> (May15, 2001).

Ebbing, Darrell D.; R. A. D. Wentworth; and James P.Birk. Introductory Chemistry. Boston: Houghton Mif-flin, 1995.

“Exploring the Table of Isotopes” (Web site). <http://ie.lbl.gov/education/isotopes.htm> (May 15, 2001).

Goldstein, Natalie. The Nature of the Atom. New York:Rosen Publishing Group, 2001.

“The Isotopes” (Web site). <http://chemlab.pc.maricopa.edu/periodic/isotopes.html> (May 15, 2001).

“Isotopes” University of Colorado Department of Physics(Web site). <http://www.colorado.edu/physics/2000/isotopes/index.html> (May 15, 2001).

Milne, Lorus Johnson and Margery Milne. Understand-ing Radioactivity. Illustrated by Bill Hiscock. NewYork: Atheneum, 1989.

Smith, Norman F. Millions and Billions of Years Ago: Dat-ing Our Earth and Its Life. New York: F. Watts, 1993.

“Stable Isotope Group.” Martek Biosciences (Web site).<http://www.martekbio.com/frmain.htm> (May 15,2001).

“Tracking with Isotopes” (Web site). <http://whyfiles.org/083isotope/2.html> (May 15, 2001).








beings.


phenomenon whereby certain materials

are subject to a form of decay brought

about by the emission of high-energy par-

ticles or radiation, including alpha parti-

cles, beta particles, or gamma rays.

RADIOISOTOPE: An isotope subject to

the decay associated with radioactivity. A

radioisotope is thus an unstable isotope.

TRACER: An atom or group of atoms

whose participation in a chemical, physi-

cal, or biological reaction can be easily

observed. Radioisotopes are often used as

tracers.

KEY TERMS CONTINUED



IONS AND IONIZATIONIons and Ionization

CONCEPTAtoms have no electric charge; if they acquire

one, they are called ions. Ions are involved in a

form of chemical bonding that produces

extremely strong bonds between metals, or

between a metal and a nonmetal. These sub-

stances, of which table salt is an example, are

called ionic compounds. Ionization is the process

whereby electrons are removed from an atom or

molecule, as well as the process whereby an ionic

substance, such as salt, is dissociated into its

component ions in a solution such as water.

There are several varieties of ionization, includ-

ing field ionization, which almost everyone has

experienced in the form of static electricity. Ion

exchange, or the replacement of one ion by

another, is used in applications such as water

purification, while chemists and physicists use

ions in mass spectrometry, to discover mass and

structural information concerning atoms and

molecules. Another example of ions at work

(and a particularly frightening example at that) is

ionizing radiation, associated with the radioac-

tive decay following a nuclear explosion.

HOW IT WORKS

Ions: Positive and Negative

Atoms have no electric charge, because they

maintain an equal number of protons (positively

charged subatomic particles) and electrons, sub-

atomic particles with a negative charge. In certain

situations, however, the atom may lose or gain

one or more electrons and acquire a net charge,

becoming an ion.

Aluminum, for instance, has an atomicnumber of 13, which tells us that an aluminumatom will have 13 protons. Given the fact thatevery proton has a positive charge, and that mostatoms tend to be neutral in charge, this meansthat there are usually 13 electrons, with a nega-tive charge, present in an atom of aluminum. Yetlike all metals, aluminum is capable of formingan ion by losing electrons—in this case, three.

CATIONS. Initially, the aluminum atomhad a charge of +13 + (–13) = 0; in other words,its charge was neutral due to the equal numbersof protons and electrons. When it becomes anion, it loses 3 electrons, leaving behind only 10.Now the charge is +13 + (–10) = +3. Thus theremaining aluminum ion is said to have a netpositive charge of 3, represented as +3 or 3+.Chemists differ as to whether they represent theplus sign (or the minus sign, in the case of a neg-atively charged ion) before or after the number.Because both systems of notation are used, thesewill be applied interchangeably throughout thecourse of this essay.

When a neutral atom loses one or more elec-trons, the result is a positively charged ion, orcation (pronounced KAT-ie-un). Cations areusually represented by a superscript number andplus sign: Al+3 or Al3+, for instance, represents thealuminum cation described above. A cation isnamed after the element of which it is an ion:thus the ion we have described is either called thealuminum ion, or the aluminum cation.

ANIONS. When a neutrally charged atomgains electrons, acquiring a negative charge as aresult, this type of ion is known as an anion (AN-ie-un). Anions can be represented symbolicallyin much the same way as cations: Cl-, for


Ions andIonization


instance, is an anion of chlorine that forms whenit acquires an electron, thus assuming a netcharge of –1. Note that the 1 is not represented inthe superscript notation, much as people do notwrite 101. In both cases, the 1 is assumed, but anynumber higher than 1 is shown.

The anion described here is never called achlorine anion; rather, anions have a specialnomenclature. If the anion represents, as was thecase here, a single element, it is named by addingthe suffix -ide to the name of the original ele-ment name: chloride. Such is the case, forinstance, with a deadly mixture of carbon andnitrogen (CN-), better known as cyanide.

Most often the -ide suffix is used, but in thecase of most anions involving more than one ele-ment (polyatomic anions), as well as with oxyan-ions (anions containing oxygen), the rules canget fairly complicated. The general principles fornaming anions are as follows:

• -ide: A single element with a negativecharge. Note, however, that both hydroxide(OH–) and cyanide (CN–) also receive the -ide suffix, even though they involve morethan one element.

• -ate: An oxyanion with the normal numberof oxygen atoms, a number that depends on

the nature of the compound. Examplesinclude oxalate (C2O4

- 2) or chlorate (ClO3

-).

• -ite: An oxyanion containing 1 less oxygenthan normal. Examples include chlorite(ClO2

–).

• hypo____ite: An oxyanion with 2 less oxy-gens than normal, but with the normalcharge. An example is hypochlorite, orClO–.

• per____ate: An oxyanion with 1 more oxy-gen than normal, but with the normalcharge. Perchlorate, or ClO4

–, is an example.

• thio-: An anion in which sulfur has replacedan oxygen. Thus, SO4

–2 is called sulfate,whereas S2O3

–2 is called thiosulfate.

Elements and Ion Charges

As one might expect, given the many differencesamong families of elements on the periodic table,different elements form ions in different ways.Yet precisely because many of these can begrouped into families, primarily according to thecolumn or group they occupy on the periodictable, it is possible to predict the ways in whichthey will form ions. The table below provides afew rules of thumb. (All group numbers refer tothe North American version of the periodic table;

A COMMON FORM OF FIELD IONIZATION IS STATIC ELECTRICITY. HERE, A GIRL PLACES HER HAND ON A STATIC ELEC-TRICITY GENERATOR. (Paul A. Souders/Corbis. Reproduced by permission.)


Ions andIonization


see Periodic Table of Elements essay for an expla-nation of the differences between this and theIUPAC version.)

• Alkali metals (Group 1) form 1+ cations.For example, the ion of lithium (Li) isalways Li+.

• Alkaline earth metals (Group 2) form 2+cations. Thus, beryllium (Be), for instance,forms a Be2+ ion.

• Most Group 3 metals (aluminum, gallium,and indium) form 3+ cations. The cation ofaluminum, thus, is designated as Al3+.

• Group 6 nonmetals and metalloids (oxy-gen, sulfur, selenium, and tellurium) form2– anions. Oxygen, in its normal ionizedstate, is shown as O2–.

• Halogens (Group 7) form 1– anions. Fluo-rine’s anion would therefore be designatedas Fl–.

The metals always form positive ions, orcations; indeed, one of the defining characteris-tics of a metal is that it tends to lose electrons.However, the many elements of the transitionmetals family form cations with a variety of dif-ferent charges; for this reason, there is no easyway to classify the ways in which these elementsform cations.

Likewise, it should be evident from the

above table that nonmetals, such as oxygen or

fluorine, gain electrons to form anions. This, too,

is a defining characteristic of this broad grouping

of elements. The reasons why these elements—

both metals and nonmetals—behave as they do

are complex, involving the numbers of valence

electrons (the electrons involved in chemical

bonding) for each group on the periodic table, as

well as the octet rule of chemical bonding,

whereby elements typically bond so that each

atom has eight valence electrons.


Ionic Bonds, Compounds,and Solids

German chemist Richard Abegg (1869-1910),

whose work with noble gases led to the discovery

of the octet rule, hypothesized that atoms com-

bine with one another because they exchange

electrons in such a way that both end up with

eight valence electrons. This was an early model

of ionic bonding, which results from attractions

between ions with opposite electric charges:

THE DAMAGED REACTOR AT THE CHERNOBYL NUCLEAR PLANT IN THE FORMER SOVIET UNION. THE 1986 ACCIDENT

AT THE PLANT RELEASED IONIZING RADIATION INTO THE ATMOSPHERE. (AP/Wide World Photos. Reproduced by permission.)


Ions andIonization

when they bond, these ions “complete” oneanother.

As noted, ionic bonds occur when a metalbonds with a nonmetal, and these bonds areextremely strong. Salt, for instance, is formed byan ionic bond between the metal sodium (a +1cation) and the nonmetal chlorine, a -1 anion.Thus Na+ joins with Cl- to form NaCl, or tablesalt. The strength of the bond in salt is reflectedby its melting point of 1,472°F (800°C), which ismuch higher than that of water at 32°F (0°C).

In ionic bonding, two ions start out with dif-ferent charges and end up forming a bond inwhich both have eight valence electrons. In acovalent bond, as is formed between nonmetals,two atoms start out as most atoms do, with a netcharge of zero. Each ends up possessing eightvalence electrons, but neither atom “owns” them;rather, they share electrons. Today, chemistsunderstand that most bonds are neither purelyionic nor purely covalent; rather, there is a widerange of hybrids between the two extremes. Thedegree to which elements attract one another is afunction of electronegativity, or the relative abil-ity of an atom to attract valence electrons.

IONIC COMPOUNDS AND

SOLIDS. Not only is salt formed by an ionicbond, but it is an ionic compound—that is, acompound containing at least one metal andnonmetal, in which ions are present. Salt is alsoan example of an ionic solid, or a crystalline solidthat contains ions. A crystalline solidis a type ofsolid in which the constituent parts are arrangedin a simple, definite geometric pattern that isrepeated in all directions. There are three types ofcrystalline solid: a molecular solid (for example,sucrose or table sugar), in which the moleculeshave a neutral electric charge; an atomic solid (adiamond, for instance, which is pure carbon);and an ionic solid.

Salt is not formed of ordinary molecules, inthe way that water or carbon dioxide are. (Bothof these are molecular compounds, though nei-ther is a solid at room temperature.) The internalstructure of salt can be depicted as a repeatingseries of chloride anions and sodium cationspacked closely together like oranges in a crate.Actually, it makes more sense to picture them ascantaloupes and oranges packed together, withthe larger chloride anions as the cantaloupes andthe smaller sodium cations as the oranges.

If a grocer had to pack these two fruitstogether in a crate, he would probably put downa layer of cantaloupes, then follow this with alayer of oranges in the spaces between the can-taloupes. This pattern would be repeated as thecrate was packed, with the smaller oranges fillingthe spaces. Much the same is true of the way thatchlorine “cantaloupes” and sodium “oranges” arepacked together in salt.

This close packing of positive and negativecharges helps to form a tight bond, and thereforesalt must be heated to a high temperature beforeit will melt. Solid salt does not conduct electrici-ty, but when melted, it makes for an extremelygood conductor. When it is solid, the ions aretightly packed, and thus there is no freedom ofmovement to carry the electric charge along; butwhen the structure is disturbed by melting,movement of ions is therefore possible.

Ionization and IonizationEnergy

Water is not a good conductor, although it willcertainly allow an electric current to flowthrough it, which is why it is dangerous to oper-ate an electrical appliance near water. We havealready seen that salt becomes a good conductorwhen melted, but this can also be achieved bydissolving it in water. This is one type of ioniza-tion, which can be defined as the process inwhich one or more electrons are removed froman atom or molecule to create an ion, or theprocess in which an ionic solid, such as salt, dis-sociates into its component ions upon being dis-solved in a solution.

The amount of energy required to achieveionization is called ionization energy or ioniza-tion potential. When an atom is at its normalenergy level, it is said to be in a ground state. Atthat point, electrons occupy their normal orbitalpatterns. There is always a high degree of attrac-tion between the electron and the positivelycharged nucleus, where the protons reside. Theenergy required to move an electron to a higherorbital pattern increases the overall energy of theatom, which is then said to be in an excited state.

The excited state of the atom is simply a stepalong the way toward ionizing it by removingthe electron. “Step” is an appropriate metaphor,because electrons do not simply drift along acontinuum from one energy level to the next, theway a person walks gently up a ramp. Rather,



Ions andIonization

they make discrete steps, like a person climbinga flight of stairs or a ladder. This is one of the keyprinciples of quantum mechanics, a cutting-edge area of physics that also has numerousapplications to chemistry. Just as we speak of asudden change as a “quantum leap,” electronsmake quantum jumps from one energy level toanother.

Due to the high attraction between the elec-tron and the nucleus, the first electron to beremoved is on the outermost orbital—that is,one of the valence electrons. This amount ofenergy is called the first ionization energy. Toremove a second electron will be considerablymore difficult, because now the atom is a cation,and the positive charge of the protons in thenucleus is greater than the negative charge of theelectrons. Hence the second ionization energy,required to remove a second electron, is muchhigher than the first ionization energy.

IONIZATION ENERGIES OF ELE-

MENTS AND COMPOUNDS. First andsecond ionization energy levels for given ele-ments have been established, though it should benoted that hydrogen, because it has only oneelectron, has only a first ionization energy. Ingeneral, the figures for ionization energy increasefrom left to right along a period or row on theperiodic table, and decrease from top to bottomalong a column or group.

The reason ionization energy increases alonga period is that nonmetals, on the right side ofthe table, have higher ionization energies thanmetals, which are on the left side. Ionizationenergies decrease along a group because elementslower on the table have higher atomic numbers,which means more protons and thus more elec-trons. It is therefore easier for them to give upone of their electrons than it is for an elementwith a lower atomic number—just as it would beeasier for a millionaire to lose a dollar than itwould be for a person earning minimum wage.

For molecules in compounds, the ionizationenergies are generally related to those of the ele-ments whose atoms make up the molecule. Justas elements with fewer electrons are typically lessinclined to give up one, so are molecules withonly a few atoms. Thus the ionization energy ofcarbon dioxide (CO2), with just three atoms, isrelatively high. Conversely, in larger molecules aswith larger atoms, there are more electrons to

give up, and therefore it is easier to separate oneof these from the molecules.

IONIZATION PROCESSES. Anumber of methods are used to produce ions formass spectrometry (discussed below) or otherapplications. The most common of these meth-ods is electron impact, produced by bombardinga sample of gas with a stream of fast-movingelectrons. Though easier than some other meth-ods, this one is not particularly efficient, becauseit supplies more energy than is needed to removethe electron. An electron gun, usually a heatedtungsten wire, produces huge amounts of elec-trons, which are then fired at the gas. Becauseelectrons are so small, this is rather like using arapid-fire machine gun to kill mosquitoes: it isalmost inevitable that some of the mosquitoeswill get hit, but plenty of the rounds will fire intothe air without hitting a single insect.

Another ionization process is field ioniza-tion, in which ionization is produced by subject-ing a molecule to a very intense electric field.Field ionization occurs in daily life, when staticelectricity builds up on a dry day: the small sparkthat jumps from the tip of your finger when youtouch a doorknob is actually a stream of elec-trons. In field ionization as applied in laborato-ries, fine, sharpened wires are used. This processis much more efficient than electron impact ion-ization, employing far less energy in relation tothat required to remove the electrons. Because itdeposits less energy into the ion source, the par-ent ion, it is often used when it is necessary not todamage the ionized specimen.

Chemical ionization employs a method sim-ilar to that of electron impact ionization, exceptthat instead of electrons, a beam of positivelycharged molecular ions is used to bombard andionize the sample. The ions used in this bom-bardment are typically small molecules, such asthose in methane, propane, or ammonia.Nonetheless, the molecular ion is much largerthan an electron, and these collisions are highlyreactive, yet they tend to be much more efficientthan electron impact ionization. Many massspectrometers use a source capable of both elec-tron impact and chemical ionization.

Ionization can also be supplied by electro-magnetic radiation, with wavelengths shorterthan those of visible light—that is, ultravioletlight, x rays, or gamma rays. This process, calledphotoionization, can ionize small molecules,



Ions andIonization

such as those of oxygen (O2). Photoionizationoccurs in the upper atmosphere, where ultravio-let radiation from the Sun causes ionization ofoxygen and nitrogen (N2) in their molecularforms.

In addition to these other ionizationprocesses, ionization can be produced in a fairlysimple way by subjecting atoms or molecules tothe heat from a flame. The temperatures, howev-er, must be several thousand degrees to be effec-tive, and therefore specialized flames, such as anelectrical arc, spark, or plasma, are typically used.

MASS SPECTROMETRY. As not-ed,a number of the ionization methods describedabove are used in mass spectrometry, a means ofobtaining structure and mass information con-cerning atoms or molecules. In mass spectrome-try, ionized particles are accelerated in a curvedpath through an electromagnetic field. The fieldwill tend to deflect lighter particles from thecurve more easily than heavier ones. By the timethe particles reach the detector, which measuresthe ratio between mass and charge, the ions willhave been separated into groups according totheir respective mass-to-charge ratios.

When molecules are subjected to mass spec-trometry, fragmentation occurs. Each moleculebreaks apart in a characteristic fashion, and thismakes it possible for a skilled observer to inter-pret the mass spectrum of the particles generat-ed. Mass spectrometry is used to establish valuesfor ionization energy, as well as to ascertain themass of substances when that mass is not known.It can also be used to determine the chemicalmakeup of a substance. Mass spectrometry isapplied by chemists, not only in pure research,but in applications within the environmental,pharmaceutical, and forensic (crime-solving)fields. Chemists for petroleum companies use it to analyze hydrocarbons, as do scientists work-ing in areas that require flavor and fragrance analysis.

Ion Exchange

The process of replacing one ion with another ofa similar charge is called ion exchange. When anionic solution—for example, salt dissolved inwater—is placed in contact with a solid contain-ing ions that are only weakly bonded within itscrystalline structure, ion exchange is possible.Throughout the exchange, electric neutrality ismaintained: in other words, the total number of

positive charges in the solid and solution equalsthe total number of negative charges. The onlything that changes is the type of ion in the solidand in the solution.

There are natural ion exchangers, such aszeolites, a class of minerals that contain alu-minum, silicon, oxygen, and a loosely held cationof an alkali metal or alkaline earth metal (Group1 and Group 2 of the periodic table, respectively).When the zeolite is placed in an ionic solution,exchange occurs between the loosely held zeolitecation and the dissolved cation. In synthetic ionexchangers, used in laboratories, a charged groupis attached to a rigid structural framework. Oneend of the charged group is permanently fixed tothe frame, while a positively or negativelycharged portion kept loose at the other endattracts other ions in the solution.

RESINS AND SEMIPERMEABLE

MEMBRANES. Anion resins and cationresins are both solid materials, but in an anionresin, the positive ions are tightly bonded, whilethe negative ions are loosely bonded. The nega-tive ions will exchange with negative ions in thesolution. The reverse is true for a cation resin, inwhich the negative ions are tightly bonded, whilethe loosely bonded positive ions exchange withpositive ions in the solution. These resins haveapplications in scientific studies, where they areused to isolate and collect various types of ionicsubstances. They are also used to purify water, byremoving all ions.

Semipermeable membranes are natural orsynthetic materials with the ability to selectivelypermit or retard the passage of charged anduncharged molecules through their surfaces. Inthe cells of living things, semipermeable mem-branes regulate the balance between sodium andpotassium cations. Kidney dialysis, whichremoves harmful wastes from the urine ofpatients with kidneys that do not function prop-erly, is an example of the use of semipermeablemembranes to purify large ionic molecules. Whenseawater is purified by removal of salt, or other-wise unsafe water is freed of its impurities, thewater is passed through a semipermeable mem-brane in a process known as reverse osmosis.

Ionizing Radiation

We have observed a number of ways in whichionization is useful, and indeed part of nature.But ionizing radiation, which causes ionization



Ions andIonization


ANION: The negatively charged ion that

results when an atom gains one or more

electrons. An anion (pronounced “AN-ie-

un”) of a single element is named by

adding the suffix -ide to the name of the

original element—hence, “chloride.” Other

rules apply for more complex anions.

CATION: The positively charged ion that

results when an atom loses one or more

electrons. A cation (pronounced KAT-ie-

un) is named after the element of which it

is an ion and, thus, is called, for instance,

the aluminum ion or the aluminum cation.

CRYSTALLINE SOLID: A type of solid

in which the constituent parts have a sim-

ple and definite geometric arrangement

that is repeated in all directions. Types of

crystalline solids include atomic solids,

ionic solids, and molecular solids.


cles in an atom. The number of electrons

and protons in an atom is the same, thus

canceling out one another. When an atom

loses or gains one or more electrons, how-

ever—thus becoming an ion—it acquires a

net electric charge.

EXCITED STATE: A term describing the

characteristics of an atom that has

acquired excess energy.


the characteristics of an atom that is at its

ordinary energy level.



thus has a net electric charge. There are two

types of ions: anions and cations.

ION EXCHANGE: The process of

replacing one ion with another of similar

charge.

IONIC BONDING: A form of chemical

bonding that results from attractions

between ions with opposite electric

charges.

IONIC COMPOUND: A compound

(two or more elements chemically bonded

to one another), in which ions are present.

Ionic compounds contain at least one metal

and nonmetal, joined by an ionic bond.

IONIC SOLID: A form of crystalline

solid that contains ions. When mixed with

a solvent such as water, ions from table

salt—an example of an ionic solid—move

freely throughout the solution, making it

possible to conduct an electric current.

IONIZATION: A term that refers to two

different processes. In one kind of ioniza-

tion, one or more electrons are removed

from an atom or molecule to create an ion.

A second kind of ionization occurs when

an ionic solid, such as salt dissociates into

its component ions upon being dissolved

in a solution.

IONIZATION ENERGY: The amount of

energy required to achieve ionization in

which one or more electrons are removed

from an atom or molecule. Another term

for this is ionization potential.





regarding the position of an electron for an

atom in a particular energy state.

OXYANION: An anion containing

oxygen.

POLYATOMIC ANION: An anion

involving more than one element.


in an atom.

KEY TERMS


Ions andIonization

in the substance through which it passes, isextremely harmful. It should be noted that not allradiation is unhealthy: after all, Earth receivesheat and light from the Sun by means of radia-tion. Ionizing radiation, on the other hand, is thekind of radiation associated with radioactive fall-out from nuclear warfare, and with nuclear dis-asters such as the one at Chernobyl in the formerSoviet Union in 1986.

Whereas thermal radiation from the Sun canbe harmful if one is exposed to it for too long,ionizing radiation is much more so, and involvesmuch greater quantities of energy deposited perarea per second. In ionizing radiation, an ioniz-ing particle knocks an electron off an atom ormolecule in a living system (that is, human, ani-mal, or plant), freeing an electron. The moleculeleft behind becomes a free radical, a highly reac-tive group of atoms with unpaired electrons.These can spawn other free radicals, inducingchemical changes that can cause cancer andgenetic damage.

WHERE TO LEARN MORE

“The Atom.” Thinkquest (Web site). <http://library.thinkquest.org/17940/texts/atom/atom.html> (May18, 2001).

“Explore the Atom” CERN—European Organization forNuclear Research (Web site). <http://public.web.cern.ch/Public/SCIENCE/Welcome.html> (May 18,2001).

Gallant, Roy A. The Ever-Changing Atom. New York:Benchmark Books, 1999.

“Ion Exchange Chromatography” (Web site). <http://ntri.tamuk.edu/fplc/ion.html> (June 1, 2001).

“Just After Big Bang, Colliding Gold Ions Exploded.” Uni-Sci: Daily University Science News (Web site).<http://unisci.com/stories/20012/0501012.htm>(June 1, 2001).


“Portrait of the Atom” (Web site). <http://www.inetarena.com/~pdx4d/snelson/Portrait.html> (May18, 2001).

“A Science Odyssey: You Try It: Atom Builder.” PBS—Pub-lic Broadcasting System (Web site). <http://www.pbs.org/wgbh/aso/tryit/atom/> (May 18, 2001).

“Table of Inorganic Ions and Rules for Naming InorganicCompounds.” Augustana College (Web site.)<http://inst.augie.edu/~dew/242mat.htm> (June 1,2001).

“Virtual Chemistry Lab.” University of Oxford Depart-ment of Chemistry (Web site). <http://www.chem.ox.ac.uk/vrchemistry/complex/menubar.html> (June 1,2001).




MOLECULESMolecules

CONCEPTPrior to the nineteenth century, chemists pur-sued science simply by taking measurements,before and after a chemical reaction, of the sub-stances involved. This was an external approach,rather like a person reaching into a box and feel-ing of the contents without actually being able tosee them. With the evolution of atomic theory,chemistry took on much greater definition: forthe first time, chemists understood that thematerials with which they worked were interact-ing on a level much too small to see. The effects,of course, could be witnessed, but the activitiesthemselves involved the interactions of atoms inmolecules. Just as an atom is the most basic par-ticle of an element, a molecule is the basic parti-cle of a compound. Whereas there are only about90 elements that occur in nature, many millionsof compounds are formed naturally or artificial-ly. Hence the study of the molecule is at least asimportant to the pursuit of modern chemistry asthe study of the atom. Among the most impor-tant subjects in chemistry are the ways in whichatoms join to form molecules—not just thenumbers and types of atoms involved, but theshape that they form together in the molecularstructure.

HOW IT WORKS

Introduction to the Molecule

Sucrose or common table sugar, of course, isgrainy and sweet, yet it is made of three elementsthat share none of those characteristics. The for-mula for sugar is C12H22O1 1, meaning that each

molecule is formed by the joining of 12 carbonatoms, 22 hydrogens, and 11 atoms of oxygen.Coal is nothing like sugar—for one thing, it is asblack as sugar is white, yet it is almost pure car-bon. Carbon, at least, is a solid at room tempera-ture, like sugar. The other two components ofsugar, on the other hand, are gases, and highlyflammable ones at that.

The question of how elements react to oneanother, producing compounds that are alto-gether unlike the constituent parts, is one of themost fascinating aspects of chemistry and,indeed, of science in general. Combined in otherways and in other proportions, the elements insugar could become water (H2O), carbon dioxide(CO2), or even petroleum, which is formed bythe joining of carbon and hydrogen.

Two different compounds of hydrogen andoxygen serve to further illustrate the curiositiesinvolved in the study of molecules. As noted,hydrogen and oxygen are both flammable, yetwhen they form a molecule of water, they can beused to extinguish most fires. On the other hand,when two hydrogens join with two oxygens toform a molecule of hydrogen peroxide (H2O2),the resulting compound is quite different fromwater. In relatively high concentrations, hydro-gen peroxide can burn the skin, and in still high-er concentrations, it is used as rocket fuel. Andwhereas water is essential to life, pure hydrogenperoxide is highly toxic.

THE QUESTION OF MOLECU-

LAR STRUCTURE. It is not enough, how-ever, to know that a certain combination ofatoms forms a certain molecule, because mole-cules may have identical formulas and yet bequite different substances. In English, for


Molecules


instance, there is the word “rose.” Simply seeing

the word, however, does not tell us whether it is a

noun, referring to a flower, or a verb, as in “she

rose through the ranks.” Similarly, the formula of

a compound does not necessarily tell what it is,

and this can be crucial.

For instance, the formula C2H6O identifies

two very different substances. One of these is

ethyl alcohol, the type of alcohol found in beer

and wine. Note that the elements involved are the

same as those in sugar, though the proportions

are different: in fact, some aspects of the body’s

reaction to ethyl alcohol are not so different from

its response to sugar, since both lead to unhealthy

weight gain. In reasonable small quantities, of

course, ethyl alcohol is not toxic, or at least only

mildly so; yet methyl ether—which has an iden-

tical formula—is a toxin.

But the distinction is not simply an external

one, as simple as the difference between beer and

a substance such as methyl ether, sometimes used

as a refrigerant. To put it another way, the exter-

nal difference reflects an internal disparity:

though the formulas for ethyl alcohol and methyl

ether are the same, the arrangements of the

atoms within the molecules of each are not. The

substances are therefore said to be isomers.

In fact C2H6O is just one of three types offormula for a compound: an empirical formula,or one that shows the smallest possible whole-number ratio of the atoms involved. By contrast,a molecular formula—a formula that indicatesthe types and numbers of atoms involved—shows the actual proportions of atoms. If the for-mula for glucose, a type of sugar (C6H12O6 ), wererendered in empirical form, it would be CH2O,which would reveal less about its actual struc-ture. Most revealing of all, however, is a structur-al formula—a diagram that shows how the atomsare bonded together, complete with lines repre-senting covalent bonds. (Structural formulassuch as those that apply the Couper or Lewis sys-tems are discussed in the Chemical Bondingessay, which also examines the subject of covalentbonds.)

Chemists involved in the area of stereo-chemistry, discussed below, attempt to developthree-dimensional models to show how atomsare arranged in a molecule. Such models for ethylalcohol and methyl ether, for instance, wouldreveal that they are quite different, much as thetwo definitions of rose mentioned above illus-trate the two distinctly different meanings.Because stereochemistry is a highly involved andcomplex subject, it can only be touched uponvery briefly in this essay; nonetheless, an under-

MODELS OF VARIOUS ELEMENTS.


Molecules


standing of a molecule’s actual shape is critical tothe work of a professional chemist.

MOLECULES AND COM-

POUNDS. A molecule can be most properlydefined as a group of atoms joined in a specificstructure. A compound, on the other hand, is asubstance made up of more than one type ofatom—in other words, more than one type ofelement. Not all compounds are composed ofdiscrete molecules, however. For instance, tablesalt (NaCl) is an ionic compound formed by end-lessly repeating clusters of sodium and chlorinethat are not, in the strictest sense of the word,molecules.

Salt is an example of a crystalline solid, or asolid in which the constituent parts are arrangedin a simple, definite geometric pattern repeatedin all directions. There are three kinds of crys-talline solids, only one of which has a trulymolecular structure. In an ionic solid such astable salt, ions (atoms, or groups of atoms, withan electric charge) bond a metal to a nonmetal—in this case, the metal sodium and the nonmetalchlorine. Another type of crystalline solid, anatomic solid, is formed by atoms of one elementbonding to one another. A diamond, made ofpure carbon, is an example. Only the third typeof crystalline solid is truly molecular in structure:

a molecular solid—sugar, for example—is one inwhich the molecules have a neutral electriccharge.

Not all solids are crystalline; nor, of course,are all compounds solids: water, obviously, is aliquid at room temperature, while carbon diox-ide is a gas. Nor is every molecule composed ofmore than one element. Oxygen, for instance, isordinarily diatomic, meaning that even in its ele-mental form, it is composed of two atoms thatjoin in an O2 molecule. It is obvious, then, thatthe defining of molecules is more complex thanit seems. One can safely say, however, that thevast majority of compounds are made up of mol-ecules in which atoms are arranged in a definitestructure.

In the essay that follows, we will discuss theways atoms join to form molecules, a subjectexplored in more depth within the ChemicalBonding essay. (In addition, compounds them-selves are examined in somewhat more detailwithin the Compounds essay.) We will alsobriefly examine how molecules bond to othermolecules in the formation of solids and liquids.First, however, a little history is in order: as notedin the introduction to this essay, chemists did notalways possess a clear understanding of thenature of a molecule.

TWO HIGH SCHOOL STUDENTS POSE WITH A MODEL OF A DNA MOLECULE. (James A. Sugar/Corbis. Reproduced by permission.)


MoleculesA Brief History of the Molecule

In ancient and medieval times, early chemists—some of whom subscribed to an unscientific sys-tem known as alchemy—believed that one ele-ment could be transformed into another. Thusmany an alchemist devoted an entire career tothe vain pursuit of turning lead into gold. Thealchemists were at least partially right, however:though one element cannot be transformed intoanother (except by nuclear fusion), it is possibleto change the nature of a compound by alteringthe relations of the elements within it.

Modern understanding of the elementsbegan to emerge in the seventeenth century, butthe true turning point came late in the eighteenthcentury. It was then that French chemist AntoineLavoisier (1743-1794) defined an element as asimple substance that could not be separated intosimpler substances by chemical means. Aroundthe same time, another French chemist, Joseph-Louis Proust (1754-1826) stated that a givencompound always contained the same propor-tions of mass between elements. The ideas ofLavoisier and Proust were revolutionary at thetime, and these concepts pointed to a substruc-ture, invisible to the naked eye, underlying allmatter.

In 1803, English chemist John Dalton (1766-1844) defined that substructure by introducingthe idea that the material world is composed oftiny particles called atoms. Despite the enormousleap forward that his work afforded to chemists,Dalton failed to recognize that matter is notmade simply of atoms. Water, for instance, is notjust a collection of “water atoms”: clearly, there issome sort of intermediary structure in whichatoms are combined. This is the molecule, a con-cept introduced by Italian physicist Amedeo Avo-gadro (1776-1856).

AVOGADRO AND THE IDEA OF

THE MOLECULE. French chemist andphysicist Joseph Gay-Lussac (1778-1850) hadannounced in 1809 that gases combine to formcompounds in simple proportions by volume. AsGay-Lussac explained, the ratio, by weight,between hydrogen and oxygen in water is eight toone. The fact that this ratio was so “clean,” involv-ing whole numbers rather than decimals,intrigued Avogadro, who in 1811 proposed thatequal volumes of gases have the same number ofparticles if measured at the same temperature

and pressure. This, in turn, led him to thehypothesis that water is not composed simply ofatoms, but of molecules in which hydrogen andoxygen combine.

For several decades, however, chemists large-ly ignored Avogadro’s idea of the molecule. Onlyin 1860, four years after his death, was the con-cept resurrected by Italian chemist StanislaoCannizzaro (1826-1910). Of course, the under-standing of the molecule has progressed enor-mously in the years since then, and much of thisprogress is an outcome of advances in the studyof subatomic structure. Only in the early twenti-eth century did physicists finally identify theelectron, the negatively charged subatomic parti-cle critical to the bonding of atoms.


Molecular Mass

Just as the atoms of elements have a definitemass, so do molecules—a mass equal to that ofthe combined atoms in the molecule. The figuresfor the atomic mass of all elements are estab-lished, and can be found on the periodic table;therefore, when one knows the mass of a hydro-gen atom and an oxygen atom, as well as the factthat there are two hydrogens and one oxygen in amolecule of water, it is easy to calculate the massof a water molecule.

Individual molecules cannot easily be stud-ied; therefore, the mass of molecules is comparedby use of a unit known as the mole. The molecontains 6.022137 � 1023 molecules, a figureknown as Avogadro’s number, in honor of theman who introduced the concept of the mole-cule. When necessary, it is possible today to studyindividual molecules, or even atoms and sub-atomic particles, using techniques such as massspectrometry.

Bonding Within Molecules

Note that the mass of an atom in a molecule doesnot change; nor, indeed, do the identities of theindividual atoms. An oxygen atom in water is thesame oxygen atom in sugar, or in any number ofother compounds. With regard to compounds, itshould be noted that these are not the same thingas a mixture, or a solution. Sugar or salt can bedissolved in water at the appropriate tempera-



Moleculestures, but the resulting solution is not a com-pound; the substances are joined physically, butthey are not chemically bonded.

Chemical bonding is the joining, throughelectromagnetic force, of atoms representing dif-ferent elements. Each atom possesses a certainvalency, which determines its ability to bondwith atoms of other elements. Valency, in turn, isgoverned by the configuration of valence elec-trons at the highest energy level (the shell) of theatom.

While studying noble gases, noted for theirtendency not to bond, German chemist RichardAbegg (1869-1910) discovered that these gasesalways have eight valence electrons. This led tothe formation of the octet rule: most elements(with the exception of hydrogen and a few oth-ers) are inclined to bond in such a way that theyend up with eight valence electrons.

When a metal bonds to a nonmetal, this isknown as ionic bonding, which results fromattractions between ions with opposite electriccharges. In ionic bonding, two ions start out withdifferent charges and form a bond in which bothhave eight valence electrons. Nonmetals, howev-er, tend to form covalent bonds. In a covalentbond, two atoms start out as most atoms do, witha net charge of zero. Each ends up possessingeight valence electrons, but neither atom “owns”them; rather, they share electrons.

ELECTRONEGATIVITY. Not all ele-ments bond covalently in the same way. Each hasa certain value of electronegativity—the relativeability of an atom to attract valence electrons.Elements capable of bonding are assigned anelectronegativity value ranging from a minimumof 0.7 for cesium to a maximum of 4.0 for fluo-rine. The greater the electronegativity value, thegreater the tendency of an element to attractvalence electrons.

When substances of differing electronegativ-ity values form a covalent bond, this is describedas polar covalent bonding. Water is an example ofa molecule with a polar covalent bond. Becauseoxygen has a much higher electronegativity (3.5)than hydrogen (2.1), the electrons tend to gravi-tate toward the oxygen atom. By contrast, mole-cules of petroleum, a combination of carbon andhydrogen, tend to be nonpolar, because carbon(2.5) and hydrogen have very similar electroneg-ativity values.

A knowledge of electronegativity values canbe used to make predictions concerning bondpolarities. Bonds that involve atoms whose elec-tronegativities differ by more than 2 units aresubstantially ionic, whereas bonds betweenatoms whose electronegativities differ by lessthan 2 units are polar covalent. If the atoms havethe same or similar electronegativity values, thebond is covalent.

Attractions Between Molecules

The energy required to pull apart a molecule isknown as bond energy. Covalent bonds thatinvolve hydrogen are among the weakest bondsbetween atoms, and hence it is relatively easy toseparate water into its constituent parts, hydro-gen and oxygen. (This is sometimes done by electrolysis, which involves the use of an electriccurrent to disperse atoms.) Double and triplecovalent bonds are stronger, but strongest of all isan ionic bond. The strength of the bond energyin salt, for instance, is reflected by its meltingpoint of 1,472°F (800°C), much higher than thatof water, at 32°F (0°C).

Bond energy relates to the attractionbetween atoms in a molecule, but in consideringvarious substances, it is also important to recog-nize the varieties of bonds between molecules—that is, intermolecular bonding. For example, thepolar quality of a water molecule gives it a greatattraction for ions, and thus ionic substancessuch as salt and any number of minerals dissolveeasily in water. On the other hand, we have seenthat petroleum is essentially nonpolar, and there-fore, an oil molecule offers no electric charge tobond it with a water molecule. For this reason, oiland water do not mix.

The bonding between water molecules isknown as a dipole-dipole attraction. This type ofintermolecular bond can be fairly strong in theliquid or solid state, though it is only about 1% asstrong as a covalent bond within a molecule.When a substance containing molecules joinedby dipole-dipole attraction is heated to become agas, the molecules spread far apart, and thesebonds become very weak. On the other hand,when hydrogen bonds to an atom with a highvalue of electronegativity (fluorine, for example),the dipole-dipole attraction between these mole-cules is particularly strong. This is known ashydrogen bonding.



Molecules


Even a nonpolar molecule, however, must

have some attraction to other nonpolar mole-

cules. The same is true of helium and the other

noble gases, which are highly nonattractive but

can be turned into liquids or even solids at

extremely low temperatures. The type of inter-

molecular attraction that exists in such a situa-

tion is described by the term London dispersion

forces. The name has nothing to do with the cap-

ital of England: it is a reference to German-

American physicist Fritz Wolfgang London(1900-1954), who in the 1920s studied the mole-cule from the standpoint of quantum mechanics.

Because electrons are not uniformly distrib-uted around the nucleus of an atom at every pos-sible moment, instantaneous dipoles are formedwhen most of the electrons happen to be on oneside of an atom. Of course, this only happens foran infinitesimal fraction of time, but it serves tocreate a weak attraction. Only at very low tem-



gadro (1776-1856), equal to 6.022137 x


number of molecules in a mole.

BOND ENERGY: The energy required

to pull apart the atoms in a chemical bond.

CHEMICAL BONDING: The joining,

through electromagnetic force, of atoms

representing different elements.




COVALENT BONDING: A type of

chemical bonding in which two atoms

share valence electrons.

DIATOMIC: A term describing an ele-

ment that exists as molecules composed of

two atoms.

DIPOLE-DIPOLE ATTRACTION: A

form of intermolecular bonding between

molecules formed by a polar covalent

bond.


cle in an atom.

ELECTRONEGATIVITY: The relative

ability of an atom to attract valence

electrons.

EMPIRICAL FORMULA: A chemical

formula that shows the smallest possible

whole-number ratio of the atoms involved.

Compare with molecular formula and

structural formula.

HYDROGEN BONDING: A kind of

dipole-dipole attraction between mole-

cules formed of hydrogen along with an

element having a high electronegativity.

INTERMOLECULAR BONDING: The

bonding that exists between molecules.

This is not to be confused with chemical

bonding, the bonding of atoms within a

molecule.





bonding resulting from attractions


charges.

ISOMERS: Substances having the same

chemical formula, but which are chemical-

ly dissimilar due to differences in the

arrangement of atoms.

LONDON DISPERSION FORCES: A

term describing the weak intermolecular

bond between molecules that are not

formed by a polar covalent bond.

KEY TERMS


Molecules


peratures do London dispersion forces becomestrong enough to result in the formation of asolid. (Thus, for instance, oil and rubbing alcoholfreeze only at low temperatures.)

Molecular Structure

The Couper system and Lewis structures, dis-cussed in the Chemical Bonding essay, provide ameans of representing the atoms that make up amolecule. Though Lewis structures show the dis-

tribution of valence electrons, they do not repre-sent the three-dimensional structure of the mol-ecule. As noted earlier in this essay, the structureis highly important, because two compoundsmay be isomers, meaning that they have the sameproportions of the same elements, yet are differ-ent substances.

Stereochemistry is the realm of chemistrydevoted to the three-dimensional arrangementof atoms in a molecule. One of the most impor-



ly speaking, Avogadro’s number of mole-

cules; however, in the more precise SI defi-

nition, a mole is equal to the number of

carbon atoms in 12.01 g of carbon.

MOLECULAR FORMULA: A chemical

formula that indicates the types and num-

bers of atoms involved, showing the actual

proportions of atoms in a molecule. Com-

pare with empirical formula and structural

formula.

MOLECULAR SOLID: A form of crys-

talline solid—a solid in which the con-

stituent parts have a simple and definite

geometric arrangement repeated in all

directions—in which the molecules have a

neutral electric charge. Table sugar

(sucrose) is an example.


but not always representing more than one

element, joined in a structure.

Compounds are typically made up of

molecules.

OCTET RULE: A term describing the

distribution of valence electrons that takes

place in chemical bonding for most ele-

ments, which end up with eight valence

electrons.

POLAR COVALENT BONDING: The

type of chemical bonding between atoms

that have differing values of electronegativ-

ity. Water molecules are an example of a

polar covalent bond.

SHELL: The orbital pattern of the

valence electrons at the outside of an atom.

STEREOCHEMISTRY: The area of

chemistry devoted to the three-dimension-

al arrangement of atoms in a molecule.

STRUCTURAL FORMULA: A diagram

that shows how the atoms are bonded

together, complete with lines representing

covalent bonds. Compare with empirical

formula and molecular formula.

VALENCE ELECTRONS: Electrons

that occupy the highest energy levels in an

atom. These are the only electrons involved

in chemical bonding.

VALENCY: The property of the atom of

one element that determines its ability to

bond with atoms of other elements.

VSEPR (VALENCE SHELL ELEC-

TRON PAIR REPULSION) MODEL: A

means of representing the three-dimen-

sional structure of atoms in a molecule.

KEY TERMS CONTINUED


Molecules tant methods used is known as the VSEPR model(valence shell electron pair repulsion). In bond-ing, elements always share at least one pair ofelectrons, and the VSEPR model begins with theassumption that the electron pairs must be as farapart as possible to minimize their repulsion,since like charges repel.

VSEPR structures can be very complex, andthe rules governing them will not be discussedhere, but a few examples can be given. If there arejust two electron pairs in a bond between threeatoms, the structure of a VSEPR model is likethat of a stick speared through a ball, with twoother balls attached at each end. The “ball” is anatom, and the “stick” represents the electronpairs. In water, there are four electron pairs, butstill only three atoms and two bonds. In order tokeep the electron pairs as far apart as possible,the angle between the two hydrogen atomsattached to the oxygen is 109.5°.

WHERE TO LEARN MORE

Basmajian, Ronald; Thomas Rodella; and Allen E. Breed.Through the Molecular Maze: A Helpful Guide to the


Elements of Chemistry for Beginning Life Science Stu-dents. Merced, CA: Bioventure Associates, 1990.

Burnie, David. Microlife. New York: DK Publishing, 1997.

“Common Molecules” (Web site). <http://www.recipnet.indiana.edu/common/common.html> (June 2,2001).

Cooper, Christopher. Matter. New York: DK Publishing,2000.

Mebane, Robert C. and Thomas R. Rybolt. Adventureswith Atoms and Molecules, Book V: Chemistry Experi-ments for Young People. Springfield, NJ: Enslow Pub-lishers, 1995.

“The Molecules of Life” (Web site). <http://biop.ox.ac.uk/www/mol_of_life/Molecules_of_Life.html> (June 2,2001).

“Molecules of the Month.” University of Oxford (Web site).<http://www.chem.ox.ac.uk/mom/> (June 2, 2001).

“Molecules with Silly or Unusual Names” (Web site).<http://www.bris.ac.uk/Depts/Chemistry/MOTM/silly/sillymols.htm> (June 2, 2001).

“Theory of Atoms in Molecules” (Web site). <http://www.chemistry.mcmaster.ca/faculty/bader/aim/> (June 2,2001).



117

SCIENCE OF EVERYDAY THINGSreal-l ife chemistry

ELEMENTSELEMENTS

ELEMENTS

PERIODIC TABLE OF ELEMENTS

FAMILIES OF ELEMENTS



ELEMENTSElements

CONCEPTThe elements are at the heart of chemistry, andindeed they are at the heart of life as well. Everyphysical substance encountered in daily life iseither an element, or more likely a compoundcontaining more than one element. There aremillions of compounds, but only about 100 ele-ments, of which only 88 occur naturally onEarth. How can such vast complexity be createdfrom such a small amount of elements? Considerthe English alphabet, with just 26 letters, withwhich an almost infinite number of things can besaid or written. Even more is possible with theelements, which greatly outnumber the letters ofthe alphabet. Despite the relatively great quantityof elements, however, just two make up almostthe entire mass of the universe—and these twoare far from abundant on Earth. A very smallnumber of elements, in fact, are essential to lifeon this planet, and to the existence of humanbeings.

HOW IT WORKS

The Focal Point of Chemistry

Like physics, chemistry is concerned with basic,underlying processes that explain how the uni-verse works. Indeed, these two sciences, alongwith astronomy and a few specialized fields, arethe only ones that address phenomena both onthe Earth and in the universe as a whole. By con-trast, unless or until life on another planet is dis-covered, biology has little concern with existencebeyond Earth’s atmosphere, except inasmuch asthe processes and properties in outer space affectastronauts.

While physics and chemistry address manyof the same fundamental issues, they do so invery different ways. To make a gross generaliza-tion—subject to numerous exceptions, butnonetheless useful in clarifying the basic differ-ence between the two sciences—physicists areconcerned with external phenomena, andchemists with internal ones.

For instance, when physicists and chemistsstudy the interactions between atoms, unlessphysicists focus on some specialized area ofatomic research, they tend to treat all atoms asmore or less the same. Chemists, on the otherhand, can never treat atoms as though they arejust undifferentiated particles colliding in space.The difference in structure between one kind ofatom and another, in fact, is the starting-point ofchemical study.

Structure of the Atom

An atom is the fundamental particle in a chemi-cal element, or a substance that cannot be brokendown into another substance by chemical means.Clustered at the center, or nucleus, of the atomare protons, which have a positive electric charge,and neutrons, which possess no charge.

Spinning around the nucleus are electrons,which are negatively charged. The vast majorityof the atom’s mass is made up by the protons andneutrons, which have approximately the samemass, whereas that of the electron is much small-er. The mass of an electron is about 1/1836 thatof proton, and 1/1839 that of a neutron.

It should be noted that the nucleus, thoughit constitutes most of the atom’s mass, is only atiny portion of the atom’s volume. If the nucleuswere the size of a grape, in fact, the electronswould, on average, be located about a mile away.


Elements


Isotopes

Atoms of the same element always have the same

number of protons, and since this figure is

unique for a given element, each element is

assigned an atomic number equal to the numberof protons in its nucleus. Two atoms may havethe same number of protons, and thus be of thesame element, yet differ in their number of neu-trons. Such atoms are called isotopes.

THE MOST COMMON AND/OR IMPORTANT CHEMICAL ELEMENTS.


ElementsIsotopes are generally represented as follows:�ma�S, where S is the symbol of the element, a is the

atomic number, and m is the mass number—thesum of protons and neutrons in the atom’s nucle-us. For the stable silver isotope designated as�94

37�Ag, for instance, Ag is the element symbol (dis-

cussed below); 47 its atomic number; and 93 themass number. From this, it is easy to discern thatthis particular stable isotope has 46 neutrons inits nucleus.

Because the atomic number of any elementis established, sometimes isotopes are represent-ed simply with the mass number thus: 93Ag. Theymay also be designated with a subscript notationindicating the number of neutrons, so that onecan obtain this information at a glance withouthaving to do the arithmetic. For the silver isotopeshown here, this is written as �

94

37�Ag46. Isotopes are

sometimes indicated by simple nomenclature aswell: for instance, carbon-12 or carbon-13.

Ions

The number of electrons in an atom is usuallythe same as the number of protons, and thusmost atoms have a neutral charge. In certain sit-uations, however, the atom may lose or gain oneor more electrons and acquire a net charge,becoming an ion. The electrons are not “lost”when an atom becomes an ion: they simply goelsewhere.

Aluminum (Al), for instance, has an atomicnumber of 13, which tells us that an aluminumatom will have 13 protons. Given the fact thatevery proton has a positive charge, and that mostatoms tend to be neutral in charge, this meansthat there are usually 13 electrons, with a nega-tive charge, present in an atom of aluminum.Aluminum may, however, form an ion by losingthree electrons.

CATIONS. After its three electrons havedeparted, the remaining aluminum ion has a netpositive charge of 3, represented as +3. How dowe know this? Initially the atom had a charge of+13 + (–13) = 0. With the exit of the 3 electrons,leaving behind only 10, the picture changes: nowthe charge is +13 + (–10) = +3.

When a neutral atom loses one or more elec-trons, the result is a positively charged ion, orcation (pronounced KAT-ie-un). Cations arerepresented by a superscript number and plussign after the element symbol: Al3+, for instance,


represents the aluminum cation described above.(Some chemists represent this with the plus signbefore the number—for example, Al+3.) A cationis named after the element of which it is an ion:thus the ion we have described is called either thealuminum ion, or the aluminum cation.

ANIONS. When a neutrally charged atomgains electrons, and as a result acquires a negativecharge, this type of ion is known as an anion(AN-ie-un). Anions can be represented symboli-cally in much the same way of cations: C1-, forinstance, is an anion of chlorine that forms whenit acquires an electron, thus assuming a netcharge of –1. Note that the 1 is not represented inthe superscript notation, much as people do notwrite 101. In both cases, the 1 is assumed, where-as any number higher than 1 is shown.

The anion described here is never called achlorine anion; rather, anions have a specialnomenclature. If the anion represents, as is thecase here, a single element, it is named by addingthe suffix -ide to the name of the original ele-ment name: hence it would be called chloride.Other anions involve more than one element,and in these cases other rules apply for designat-ing names. A few two-element anions use the -ideending; such is the case, for instance, with a dead-ly mixture of carbon and nitrogen (CN–), betterknown as cyanide.

For anions involving oxygen, there may bedifferent prefixes and suffixes, depending on therelative number of oxygen atoms in the anion.


The Periodic Table

Note that an ion is never formed by a change inthe number of protons: that number, as notedearlier, is a defining characteristic of an element.If all we know about a particular atom is that ithas one proton, we can be certain that it is anatom of hydrogen. Likewise, if an atom has 79protons, it is gold.

Knowing these quantities is not a matter ofmemorization: rather, one can learn this andmuch more by consulting the periodic table ofelements. The periodic table is a chart, present invirtually every chemistry classroom in the world,


Elements showing the elements arranged in order of atom-ic number. Elements are represented by boxescontaining the atomic number, element symbol,and average atomic mass, in atomic mass units,for that particular element. Vertical columnswithin the periodic table indicate groups or“families” of elements with similar chemicalcharacteristics.

These groups include alkali metals, alkalineearth metals, halogens, and noble gases. In themiddle of the periodic table is a wide range ofvertical columns representing the transition met-als, and at the bottom of the table, separatedfrom it, are two other rows for the lanthanidesand actinides.

An Overview of the Elements

As of 2001, there 112 elements, of which 88 occurnaturally on Earth. (Some sources show 92 natu-rally occurring elements; however, a few of theelements with atomic numbers below 92 havenot actually been found in nature.) The otherswere created synthetically, usually in a laborato-ry, and because these are highly radioactive, theyexist only for fractions of a second. The numberof elements thus continues to grow, but these“new” elements have little to do with the dailylives of ordinary people. Indeed, this is true evenfor some of the naturally occurring elements: fewpeople who are not chemically trained, forinstance, are able to identify thulium, which hasan atomic number of 69.

Though an element can theoretically exist asa gas, liquid, or a solid, in fact the vast majority ofelements are solids. Only 11 elements—the sixnoble gases, along with hydrogen, nitrogen, oxy-gen, fluorine, and chlorine—exist in the gaseousstate at a normal temperature of about 77°F(25°C). Just two are liquids at normal tempera-ture: mercury, a metal, and the non-metalbromine. (The metal gallium becomes liquid atjust 85.6°F, or 29.76°C.) The rest are all solids.

The noble gases are monatomic, meaningthat they exist purely as single atoms. So too arethe “noble metals,” such as gold, silver, and plat-inum. “Noble” in this context means “set apart”:noble gases and noble metals are known for theirtendency not to react to, and hence not to bondwith, other elements. On the other hand, a num-ber of other elements are described as diatomic,meaning that two atoms join to form a molecule.

Names of the Elements

ELEMENT SYMBOL. As noted above,the periodic table includes the element symbol orchemical symbol—a one-or two-letter abbrevia-tion for the name of the element. Many of theseare simple one-letter designations: O for oxygen,or C for carbon. Others are two-letter abbrevia-tions, such as Ne for neon or Si for silicon. Notethat the first letter is always capitalized, and thesecond is always lowercase.

In many cases, the two-letter symbols indi-cate the first and second letters of the element’sname, but this is far from universal. Cadmium,for example, is abbreviated Cd, while platinum isPt. In other cases, the symbol seems to havenothing to do with the element name as it is nor-mally used: for instance, Au for gold or Fe foriron.

Many of the one-letter symbols indicate ele-ments discovered early in history. For instance,carbon is represented by C, and later “C” ele-ments took two-letter designations: Ce for ceri-um, Cr for chromium, and so on. But many ofthose elements with apparently strange symbolswere among the first discovered, and this is pre-cisely why the symbols make little sense to a per-son who does not recognize the historical originsof the name.

HISTORICAL BACKGROUND ON

SOME ELEMENT NAMES. For manyyears, Latin was the language of communicationbetween scientists from different nations; hencethe use of Latin names such as aurum (“shiningdawn”) for gold, or ferrum, the Latin word foriron. Likewise, lead (Pb) and sodium (Na) aredesignated by their Latin names, plumbum andnatrium, respectively.

Some chemical elements are named forGreek or German words describing properties ofthe element—for example, bromine (Br), whichcomes from a Greek word meaning “stench.” Thename of cobalt comes from a German termmeaning “underground gnome,” because minersconsidered the metal a troublemaker. The namesof several elements with high atomic numbersreflect the places where they were originally dis-covered or created: francium, germanium,americium, californium.

Americium and californium, with atomicnumbers of 95 and 98 respectively, are amongthose elements that do not occur naturally, but



Elementswere created artificially. The same is true of sev-eral elements named after scientists—amongthem einsteinium, after Albert Einstein (1879-1955), and nobelium after Alfred Nobel (1833-1896), the Swedish inventor of dynamite whoestablished the Nobel Prize.

Abundance of Elements

IN THE UNIVERSE. The first two ele-ments on the periodic table, hydrogen and heli-um, represent 99.9% of the matter in the entireuniverse. This may seem astounding, but Earth isa tiny dot within the vastness of space, andhydrogen and helium are the principal elementsin stars.

All elements, except for those created artifi-cially, exist both on Earth and throughout theuniverse. Yet the distribution of elements onEarth is very, very different from that in otherplaces—as well it should be, given the fact thatEarth is the only planet known to support life.Hydrogen, for instance, constitutes only about0.87%, by mass, of the elements found in theplanet’s crust, waters, and atmosphere. As forhelium, it is not even among the 18 most abun-dant elements on Earth.

ON EARTH. That great element essentialto animal life, oxygen, is by far the most plentifulon Earth, representing nearly half—49.2%—ofthe total mass of atoms found on this planet.(Here the term “mass” refers to the known ele-mental mass of the planet’s atmosphere, waters,and crust; below the crust, scientists can onlyspeculate, though it is likely that much of Earth’sinterior consists of iron.) Together with silicon(25.7%), oxygen accounts for almost exactlythree-quarters of the elemental mass of Earth.Add in aluminum (7.5%), iron (4.71%), calcium(3.39%), sodium (2.63%), potassium (2.4%),and magnesium (1.93%), and these eight ele-ments make up about 97.46% of Earth’s material.

In addition to hydrogen, whose distributionis given above, nine other elements account for atotal of 2% of Earth’s composition: titanium(0.58%), chlorine (0.19%), phosphorus (0.11%),manganese (0.09%), carbon (0.08%), sulfur(0.06%), barium (0.04%), nitrogen (0.03%), andfluorine (0.03%). The remaining 0.49% is madeup of various elements.

Elements In the Human Body

Fans of science-fiction are familiar with thephrase “carbon-based life form,” which is used,for instance, by aliens in sci-fi movies to describehumans. In fact, the term is a virtual redundancyon Earth, since all living things contain carbon.

Essential though it is to life, carbon as acomponent of the human body takes secondplace to oxygen, which is an even larger propor-tion of the body’s mass—65.0%—than it is of theEarth. Carbon accounts for 18%, and hydrogenfor 10%, meaning that these three elements makeup 93% of the body’s mass. Most of the remain-der is taken up by 10 other elements: nitrogen(3%), calcium (1.4%), phosphorus (1.0%), mag-nesium (0.50%), potassium (0.34%), sulfur(0.26%), sodium (0.14%), chlorine (0.14%), iron(0.004%), and zinc (0.003%).

TRACE ELEMENTS. As small as theamount of zinc is in the human body, there arestill other elements found in even smaller quan-tities. These are known as trace elements, becauseonly traces of them are present in the body. Yetthey are essential to human well-being: withoutenough iodine, for instance, a person can devel-


THIS WOMAN’S GOITER IS PROBABLY THE RESULT OF A

LACK OF IODINE IN HER DIET. (Lester V. Bergman/Corbis. Repro-

duced by permission.)


Elements


op a goiter, a large swelling in the neck area.Chromium helps the body metabolize sugars,which is why people concerned with losingweight and/or toning their bodies through exer-cise may take a chromium supplement.

Even arsenic, lethal in large quantities, is atrace element in the human body, and medicinesfor treating illnesses such as the infection knownas “sleeping sickness” contain tiny amounts ofarsenic. Other trace elements include cobalt, cop-

per, fluorine, manganese, molybdenum, nickel,selenium, silicon, and vanadium.

MAINTAINING AND IMPROVING

HEALTH. Though these elements are presentin trace quantities within the human body, thatdoes not mean that exposure to large amounts ofthem is healthy. Arsenic, of course, is a goodexample; so too is aluminum. Aluminum is pres-ent in unexpected places: in baked goods, forinstance, where a compound containing alu-



An anion (pronounced “AN-ie-un”) of an

element is never called, for instance, the

chlorine anion. Rather, for an anion

involving a single element, it is named by





ment. An atom can exist either alone or in

combination with other atoms in a mole-

cule. Atoms are made up of protons, neu-

trons, and electrons. An atom that loses or

gains one or more electrons, and thus has a

net charge, is an ion. Atoms that have the

same number of protons—that is, are of

the same element—but differ in number of

neutrons are known as isotopes.







elements are listed on the periodic table in





a large sample. If a substance is a com-

pound, the average atomic mass of all

atoms in a molecule of that substance must

be added together to yield the average

molecular mass of that substance.



A cation (pronounced KAT-ie-un) is

named after the element of which it is an

ion and thus is called, for instance, the alu-

minum ion or the aluminum cation.

CHEMICAL SYMBOL: Another term

for element symbol.






two atoms. This is in contrast to monatom-

ic elements.





very small portion of the atom’s mass. The

number of electrons and protons is the

same, thus canceling out one another. But

when an atom loses or gains one or more

electrons, however—thus becoming an

ion—it acquires a net electric charge.

KEY TERMS


Elements


minum (baking powder) is sometimes used inthe leavening process; or even in cheeses, as anaid to melting when heated. The relatively highconcentrations of aluminum in these products,as well as fluorine in drinking water, has raisedconcerns among some scientists.

Generally speaking, an element is healthy forthe human body in proportion to its presence inthe body. With trace elements and others that arefound in smaller quantities, however, it is some-

times possible and even advisable to increase the

presence of those elements by taking dietary sup-

plements. Hence a typical multivitamin contains

calcium, iron, iodine, magnesium, zinc, seleni-

um, copper, chromium, manganese, molybde-

num, boron, and vanadium.

For most of these, recommended daily

allowances (RDA) have been established by the

federal government. Usually, people do not take


one kind of atom. Unlike compounds, ele-

ments cannot be chemically broken into

other substances.

ELEMENT SYMBOL: A one- or two-

letter abbreviation for the name of an ele-

ment. These may be a single capitalized let-

ter (O for oxygen), or a capitalized letter

followed by a lowercase one (Ne for neon).

Sometimes the second letter is not the sec-

ond letter of the element name (for exam-

ple, Pt for platinum). In addition, the sym-

bol may refer to an original Greek or Latin

name, rather than the name used now, is

the case with Au (aurum) for gold.

ION: An atom that has lost or gained one

or more electrons, and thus has a net elec-

tric charge.




neutrons.


and neutrons in an atom’s nucleus. Where

an isotope is represented, the mass number

is placed above the atomic number to the

left of the element symbol.

MOLECULE: A group of atoms, usu-ally

but not always representing more



molecules.

MONATOMIC: A term describing an ele-

ment that exists as single atoms. This in

contrast to diatomic elements.




protons.






order of atomic number, along with ele-

ment symbol and the average atomic mass


element. Vertical columns within the peri-

odic table indicate groups or “families”

of elements with similar chemical charac-

teristics.









KEY TERMS CONTINUED


Elements sodium as a supplement, though—Americansalready get more than their RDA of sodiumthrough salt, which is overly abundant in theAmerican diet.

WHERE TO LEARN MORE

Challoner, Jack. The Visual Dictionary of Chemistry. NewYork: DK Publishing, 1996.

“Classification of Elements and Compounds” (Web site).<http://dl.clackamas.cc.or.us/ch104-04/classifi.htm>(May 14, 2001).

“Elements and Compounds” (Web site).<http://wine1.sb.fsu.edu/chm1045/notes/Intro/Elements/Intro02.htm> (May 14, 2001).

“Elements, Compounds, and Mixtures.” Purdue UniversityDepartment of Chemistry. (Web site).

<http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch2/mix.html> (May 14, 2001).

“Elements, Compounds, and Mixtures” (Web site).<http://www.juniorcert.net/serve/cont.php3?pg=SC2CEC0339> (May 14, 2001).

Knapp, Brian J. Elements. Illustrated by David Woodroffeand David Hardy. Danbury, CT: Grolier Educational,1996.

“Matter—Elements, Compounds, and Mixtures” (Website). <http://chem.sci.gu.edu.au/help_desk/Matter.htm> (May 14, 2001).

Oxlade, Chris. Elements and Compounds. Chicago, IL:Heinemann Library, 2001.

Stwertka, Albert. A Guide to the Elements. New York:Oxford University Press, 1998.





PERIODIC TABLE OF ELEMENTSPeriodic Table of Elements

CONCEPTIn virtually every chemistry classroom on theplanet, there is a chart known as the periodictable of elements. At first glance, it looks like amere series of boxes, with letters and numbers inthem, arranged according to some kind of codenot immediately clear to the observer. The boxeswould form a rectangle, 18 across and 7 deep, butthere are gaps in the rectangle, particularly alongthe top. To further complicate matters, two rowsof boxes are shown along the bottom, separatedfrom one another and from the rest of the table.Even when one begins to appreciate all the infor-mation contained in these boxes, the periodictable might appear to be a mere chart, ratherthan what it really is: one of the most sophisti-cated and usable means ever designed for repre-senting complex interactions between the build-ing blocks of matter.

HOW IT WORKS

Introduction to the PeriodicTable

As a testament to its durability, the periodictable—created in 1869—is still in use today.Along the way, it has incorporated modificationsinvolving subatomic properties unknown to theman who designed it, Russian chemist DmitriIvanovitch Mendeleev (1834-1907). YetMendeleev’s original model, which we will dis-cuss shortly, was essentially sound, inasmuch as itwas based on the knowledge available to chemistsat the time.

In 1869, the electromagnetic force funda-mental to chemical interactions had only recent-

ly been identified; the modern idea of the atomwas less than 70 years old; and another threedecades were to elapse before scientists beganuncovering the substructure of atoms that causesthem to behave as they do. Despite these limita-tions in the knowledge available to Mendeleev,his original table was sound enough that it hasnever had to be discarded, but merely clarifiedand modified, in the years since he developed it.

The rows of the periodic table of elementsare called periods, and the columns are known asgroups. Each box in the table represents an ele-ment by its chemical symbol, along with itsatomic number and its average atomic mass inatomic mass units. Already a great deal has beensaid, and a number of terms need to beexplained. These explanations will require thelength of this essay, beginning with a little histor-ical background, because chemists’ understand-ing of the periodic table—and of the elementsand atoms it represents—has evolved consider-ably since 1869.

Elements and Atoms

An element is a substance that cannot be brokendown chemically into another substance. Anatom is the smallest particle of an element thatretains all the chemical and physical properties ofthe element, and elements contain only one kindof atom. The scientific concepts of both elementsand atoms came to us from the ancient Greeks,who had a rather erroneous notion of the ele-ment and—for their time, at least—a highlyadvanced idea of the atom.

Unfortunately, atomic theory died away inlater centuries, while the mistaken notion of four“elements” (earth, air, fire, and water) survived


PeriodicTable ofElements

In A New System of Chemical Philosophy(1808), Dalton put forward the idea that nature iscomposed of tiny particles, and in so doing headopted the Greek word atomos to describe thesebasic units. Drawing on Proust’s law of constantcomposition, Dalton recognized that the struc-ture of atoms in a particular element or com-pound is uniform, but maintained that com-pounds are made up of compound “atoms.” Infact, these compound atoms are really molecules,or groups of two or more atoms bonded to oneanother, a distinction clarified by Italian physicistAmedeo Avogadro (1776-1856).

Dalton’s and Avogadro’s contemporary,Swedish chemist Jons Berzelius (1779-1848),developed a system of comparing the mass ofvarious atoms in relation to the lightest one,hydrogen. Berzelius also introduced the system ofchemical symbols—H for hydrogen, O for oxy-gen, and so on—in use today. Thus, by the mid-dle of the nineteenth century, scientists under-stood vastly more about elements and atomsthan they had just a few decades before, and theneed for a system of organizing elements becameincreasingly clear. By mid-century, a number ofchemists had attempted to create just such anorganizational system, and though Mendeleev’swas not the first, it proved the most useful.

Mendeleev Constructs HisTable

By the time Mendeleev constructed his periodictable in 1869, there were 63 known elements. Atthat point, he was working as a chemistry profes-sor at the University of St. Petersburg, where hehad become acutely aware of the need for a wayof classifying the elements to make their relation-ships more understandable to his students. Hetherefore assembled a set of 63 cards, one foreach element, on which he wrote a number ofidentifying characteristics for each.

Along with the element symbol, discussedbelow, he included the atomic mass for the atomsof each. In Mendeleev’s time, atomic mass wasunderstood simply to be the collective mass of aunit of atoms—a unit developed by Avogadro,known as the mole—divided by Avogadro’snumber, the number of atoms or molecules in amole. With the later discovery of subatomic par-ticles, which in turn made possible the discoveryof isotopes, figures for atomic mass were clari-fied, as will also be discussed.


virtually until the seventeenth century, an erathat witnessed the birth of modern science. Yetthe ancients did know of substances later classi-fied as elements, even if they did not understandthem as such. Among these were gold, tin, cop-per, silver, lead, and mercury. These, in fact, aresuch an old part of human history that their dis-coverers are unknown. The first individual cred-ited with discovering an element was Germanchemist Hennig Brand (c. 1630-c. 1692), whodiscovered phosphorus in 1674.

MATURING CONCEPTS OF

ATOMS, ELEMENTS, AND MOLE-

CULES. The work of English physicist andchemist Robert Boyle (1627-1691) greatlyadvanced scientific understanding of the ele-ments. Boyle maintained that no substance wasan element if it could be broken down into othersubstances: thus air, for instance, was not an ele-ment. Boyle’s studies led to the identification ofnumerous elements in the years that followed,and his work influenced French chemistsAntoine Lavoisier (1743-1794) and Joseph-LouisProust (1754-1826), both of whom helped definean element in the modern sense. These men inturn influenced English chemist John Dalton(1766-1844), who reintroduced atomic theory tothe language of science.

DMITRI MENDELEEV. (Bettmann/Corbis. Reproduced by per-

mission.)


PeriodicTable of

Elements


In addition, Mendeleev also included figuresfor specific gravity—the ratio between the densi-ty of an element and the density of water—aswell as other known chemical characteristics ofan element. Today, these items are typically nolonger included on the periodic table, partly forconsiderations of space, but partly becausechemists’ much greater understanding of theproperties of atoms makes it unnecessary to clut-ter the table with so much detail.

Again, however, in Mendeleev’s time therewas no way of knowing about these factors. As faras chemists knew in 1869, an atom was an indi-visible little pellet of matter that could not becharacterized by terms any more detailed than itsmass and the ways it interacted with atoms of other elements. Mendeleev therefore ar-

ranged his cards in order of atomic mass, thengrouped elements that showed similar chemicalproperties.

CONFIDENT PREDICTIONS. AsMendeleev observed, every eighth element on thechart exhibits similar characteristics, and thus, heestablished columns whereby element number xwas placed above element number x + 8—forinstance, helium (2) above neon (10). The pat-terns he observed were so regular that for any“hole” in his table, he predicted that an elementto fill that space would be discovered.

Indeed, Mendeleev was so confident in thebasic soundness of his organizational system thatin some instances, he changed the figures for theatomic mass of certain elements because he wasconvinced they belonged elsewhere on the table.

Main–Group Elements

1

2

3

4

5

6

7

Peri

od

Main–Group Elements

*Lanthanides

Actinides

Inner–Transition Metals

†

112

Uubununbiium

(277)

71

Lulutetium

174.967

103

Lr

lawrencium

(262)

1

Hhydrogen

1.00794

3Lilithium

6.941

11Nasodium

22.989768

37

Rbrubidium

85.4678

55

Cscesium

132.90543

87

Frfrancium

(223) 88

Raradium

(226)

56

Babarium

137.327

38

Srstrontium

87.62

20

Cacalcium

40.078

12Mgmagnesium

24.3050

4Beberyllium

9.012182

II A2

I A1

21

Scscandium

44.955910

39

Yyttrium

88.90585

104

Rfrutherfordium

(261)

72

Hfhafnium

178.49

40

Zrzirconium

91.224

22

Tititanium

47.88 23

Vvanadium

50.9415

41

Nbniobium

92.90638

73

Tatantalum

180.9479

105

Dbdubnium

(262) 106

Sgseaborgium

(263)

74

Wtungsten

183.85

42

Momolybdenum

95.94

24

Crchromium

51.9961 25

Mnmanganese

54.9305

43

Tctechnetium

(98)

75

Rerhenium

186.207

107

Bhbohrium

(264) 108

Hshassium

(265)

76

Ososmium

190.2

44

Ruruthenium

101.07

26

Feiron

55.847

109

Mtmeitnerium

(268)

77

Iriridium

192.22

45

Rhrhodium

102.90550

27

Cocobalt

58.93320 28

Ninickel

58.69

46

Pdpalladium

106.42

78

Ptplatinum

195.08

110

Uunununnilium

(269) 111

Uuuunununium

(272)

79

Augold

196.96654

47

Agsilver

107.8682

29

Cucopper

63.546 30

Znzinc

65.39

48

Cdcadmium

112.411

80

Hgmercury

200.59 81

Tlthallium

204.3833

49

Inindium

114.82

31

Gagallium

69.723

13

Alaluminum

26.981539

B

5

boron

10.811

III A13

IV A14

C

6

carbon

12.011

14

Sisilicon

28.0855

32

Gegermanium

72.61

50

Sntin

118.710

82

Pblead

207.2

114

Uuqununquadium

(289)

83

Bibismuth

208.98037

51

Sbantimony

121.75

33

Asarsenic

74.92159

15

Pphosphorus

30.973762

N

7

nitrogen

14.00674

V A15

VI A16

O

8

oxygen

15.9994

16

Ssulfur

32.066

34

Seselenium

78.96

52

Tetellurium

127.60

84

Popolonium

(209)

116

Uuhununhexium

(289)

F

9

fluorine

18.9984032

17

Clchlorine

35.4527

35

Brbromine

79.904

53

Iiodine

126.90447

118

Uuoununoctium

(293)

86

Rnradon

(222)

54Xexenon

131.29

36

Krkrypton

83.80

18

Arargon

39.948

10

Neneon

20.1797

2

Hehelium

4.002602

VIII A18

VII A17

19

Kpotassium

39.0983

Ac

89

actinium

(227) 102

No

nobelium

(259)

70

Ybytterbium

173.0469

Tmthulium

168.93421

101

Md

mendelevium

(258)100

Fm

fermium

(257)

68

Ererbium

167.2667

Hoholmium

164.93032

99

Es

einsteinium

(252)98

Cfcalifornium

(251)

66

Dydysprosium

162.50

97

Bkberkelium

(247)96

Cmcurium

(247)

64

Gdgadolinium

157.25

95

Amamericium

(243)

63

Eueuropium

151.965

94

Puplutonium

(244)93

Npneptunium

(237)

61

Pmpromethium

(145) 62

Smsamarium

150.3660

Ndneodymium

144.24

92

Uuranium

238.028991

Paprotactinium

(231)

59

Prpraseodymium

140.90765

90

Ththorium

232.0381

58

Cecerium

140.115

La

57

lanthanum

138.9055

III B3

IV B4

V B5

VI B6

VII B7

II B12

I B118 10

Transition Metals

VIII B9

65

Tbterbium

158.92534

85

Atastatine

(210)

86

Rnradon

(222)Atomic numberSymbol

Name

Atomic weight

67-70*

89-102†

THE PERIODIC TABLE (IUPAC SYSTEM)



Later discoveries of isotopes, which in some casesaffected the average atomic mass considerably,confirmed his suppositions. Likewise the undis-covered elements he named “eka-aluminum,”“eka-boron,” and “eka-silicon” were later identi-fied as gallium, scandium, and germanium,respectively.


Subatomic Structures Clari-fy the Periodic Table

Over a period of 35 years, between the discoveryof the electron in 1897 and the discovery of theneutron in 1932, chemists’ and physicists’ under-standing of atomic structure changed complete-ly. The man who identified the electron was Eng-lish physicist J. J. Thomson (1856-1940). Theelectron is a negatively charged particle that con-tributes little to an atom’s mass; however, it has agreat deal to do with the energy an atom possess-es. Thomson’s discovery made it apparent thatsomething else had to account for atomic mass,as well as the positive electric charge offsettingthe negative charge of the electron.

Thomson’s student Ernest Rutherford(1871-1937)—for whom, incidentally, ruther-fordium (104 on the periodic table) is named—identified that “something else.” In a series ofexperiments, he discovered that the atom has anucleus, a center around which electrons move,and that the nucleus contains positively chargedparticles called protons. Protons have a mass1,836 times as great as that of an electron, andthus, this seemed to account for the total atomicmass.

ISOTOPES AND ATOMIC MASS.

Later, working with English chemist FrederickSoddy (1877-1956), Rutherford discovered thatwhen an atom emitted certain types of particles,its atomic mass changed. Rutherford and Soddynamed these atoms of differing mass isotopes,though at that point—because the neutron hadyet to be discovered—they did not know exactlywhat change had caused the change in mass. Cer-tain types of isotopes, Soddy and Rutherfordwent on to conclude, had a tendency to decay byemitting particles or gamma rays, moving(sometimes over a great period of time) towardstabilization. In the process, these radioactive iso-

topes changed into other isotopes of the sameelement—and sometimes even to isotopes ofother elements.

Soddy concluded that atomic mass, as meas-ured by Berzelius, was actually an average of themass figures for all isotopes within that element.This explained a problem with Mendeleev’s peri-odic table, in which there seemed to be irregular-ities in the increase of atomic mass from elementto element. The answer to these variations inmass, it turned out, related to the number of iso-topes associated with a given element: the greaterthe number of isotopes, the more these affectedthe overall measure of the element’s mass.

A CLEARER DEFINITION OF

ATOMIC NUMBER. Just a few years afterRutherford and Soddy discovered isotopes,Welsh physicist Henry Moseley (1887-1915)uncovered a mathematical relationship betweenthe amount of energy a given element emittedand its atomic number. Up to this point, the peri-odic table had assigned atomic number in orderof mass, beginning with the lightest element,hydrogen. Using atomic mass and other charac-teristics as his guides, Mendeleev had been ableto predict the discovery of new elements, butsuch predictions had remained problematic.Thanks to Moseley’s work, it became possible topredict the existence of undiscovered elementswith much greater accuracy.

As Moseley discovered, the atomic numbercorresponds to the number of positive charges inthe nucleus. Thus carbon, for instance, has anatomic number of 6 not because there are fivelighter elements—though this is also true—butbecause it has six protons in its nucleus. Theordering by atomic number happens to corre-spond to the ordering by atomic mass, but atom-ic number provides a much more precise meansof distinguishing elements. For one thing, atom-ic number is always a whole integer—1 forhydrogen, for instance, or 17 for chlorine, or 92for uranium. Figures for mass, on the other hand,are almost always rendered with whole numbersand decimal fractions (for example, 1.008 forhydrogen).

If atoms have no electric charge, meaningthat they have the same number of protons aselectrons, then why do chemists not say thatatomic number represents the number of pro-tons or electrons? The reason is that electrons caneasily be lost or gained by atoms to form ions,



PeriodicTable of

Elements

which have an electric charge. However, protonsare very hard to remove.

NEUTRONS AND ATOMIC MASS.

By 1932, scientists had come a long way towardunderstanding the structure of the atom. Notonly had the electron, nucleus, and proton beendiscovered, but the complex model of electronconfiguration (described later in this essay) hadbegun to evolve. Yet, one nagging questionremained: the mass of the protons in the nucleussimply could not account for the entire mass ofthe atom. Neither did the electrons make a sig-nificant contribution to mass.

Suppose a proton was “worth” $1,836, whilean electron had a value of only $1. In the “bankaccount” for deuterium, an isotope of hydrogen,there is $3,676, which poses a serious discrepan-cy in accounting. Because deuterium is a form ofhydrogen, it has one proton as well as one elec-tron, but that only accounts for $1,837. Wheredoes deuterium get the other $1,839? Thesenumbers are not chosen at random, as we shall see.

The answer to the problem of atomic masscame when English physicist James Chadwick(1891-1974) identified the neutron, a particlewith no electric charge, residing in the nucleusalongside the protons. Whereas the proton has amass 1,836 times as large as that of the electron,the neutron’s mass is slightly larger—1,839 timesthat of an electron. This made it possible to clar-ify the values of atomic mass, which up to thattime had been problematic, because a mole ofatoms representing one element is likely to con-tain numerous isotopes.

Average Atomic Mass

Today, the periodic table lists, along with chemi-cal symbol and atomic number, the averageatomic mass of each element. As its name sug-gests, the average atomic mass provides the aver-age value of mass—in atomic mass units(amu)—for a large sample of atoms. Accordingto Berzelius’s system for measuring atomic mass,1 amu should be equal to the mass of a hydrogenatom, even though that mass had yet to be meas-ured, since hydrogen almost never appears alonein nature. Today, in accordance with a 1960agreement among members of the internationalscientific community, measurements of atomic

mass take carbon-12, an isotope found in all liv-ing things, as their reference point.

It is inconvenient, to say the least, to measurethe mass of a single carbon-12 atom, or indeed ofany other atom. Instead, chemists use a largenumber of atoms, a value known as Avogadro’snumber, which in general is the number of atomsin a mole (abbreviated mol). Avogadro’s numberis defined as 6.02214199 • 1023, with an uncer-tainty of 4.7 • 1016. In other words, the number ofparticles in a mole could vary by as much as47,000,000,000,000,000 on either side of thevalue for Avogadro’s number. This might seemlike a lot, but in fact it is equal to only about 80parts per billion.

When 1 is divided by Avogadro’s number,the result is 1.66 • 10–24—the value, in grams, of 1amu. However, according to the 1960 agreement,1 amu is officially 1/12 the mass of a carbon-12atom, whose exact value (re-tested in 1998), is1.6653873 � 10–24 g. Carbon-12, sometimes rep-resented as (12/6)C, contains six protons and sixneutrons, so the value of 1 amu thus obtained is,in effect, an average of the mass for a proton andneutron.

Though atoms differ, subatomic particles donot. There is no such thing, for instance, as a“hydrogen proton”—otherwise, these subatomicparticles, and not atoms, would constitute thebasic units of an element. Given the unvaryingmass of subatomic particles, combined with thefact that the neutron only weighs 0.16% morethan a proton, the established value of 1 amuprovides a convenient means of comparing mass.This is particularly useful in light of the largenumbers of isotopes—and hence of varying fig-ures for mass—that many elements have.

ATOMIC MASS UNITS AND THE

PERIODIC TABLE. The periodic table as itis used today includes figures in atomic massunits for the average mass of each atom. As itturns out, Berzelius was not so far off in his useof hydrogen as a standard, since its mass is almostexactly 1 amu—but not quite. The value is actu-ally 1.008 amu, reflecting the presence of slightlyheavier deuterium isotopes in the average sampleof hydrogen

Figures increase from hydrogen along theperiodic table, though not by a regular pattern.Sometimes the increase from one element to thenext is by just over 1 amu, and in other cases, the




increase is by more than 3 amu. This only servesto prove that atomic number, rather than atomicmass, is a more straightforward means of order-ing the elements.

Mass figures for many elements that tend toappear in the form of radioactive isotopes areusually shown in parentheses. This is particular-ly true for elements with very, very high atomicnumbers (above 92), because samples of theseelements do not stay around long enough to bemeasured. Some have a half-life—the period inwhich half the isotopes decay to a stable form—of just a few minutes, and for others, the half-lifeis a fraction of a second. Therefore, atomic massfigures represent the mass of the longest-livedisotope.

Elements

As of 2001, there were 112 known elements, ofwhich about 90 occur naturally on Earth. Urani-um, with an atomic number of 92, was the lastnaturally occurring element discovered: hencesome sources list 92 natural elements. Othersources, however, subtract those elements with alower atomic number than uranium that werefirst created in laboratories rather than discov-ered in nature. In any case, all elements withatomic numbers higher than 92 are synthetic,meaning that they were created in laboratories.Of these 20 elements—all of which haveappeared only in the form of radioactive isotopeswith short half-lives—the last three have yet toreceive permanent names.

In addition, three other elements—designat-ed by atomic numbers 114, 116, and 118, respec-tively—are still on the drawing board, as it were,and do not yet even have temporary names. Thenumber of elements thus continues to grow, butthese “new” elements have little to do with thedaily lives of ordinary people. Indeed, this is trueeven for some of the naturally occurring ele-ments: for example, few people who are notchemically trained would be able to identifyyttrium, which has an atomic number of 39.

Though an element can exist theoretically asa gas, liquid, or a solid, in fact, the vast majorityof elements are solids. Only 11 elements exist inthe gaseous state at a normal temperature ofabout 77°F (25°C). These are the six noble gases;fluorine and chlorine from the halogen family; aswell as hydrogen, nitrogen, and oxygen. Just two

are liquids at normal temperature: mercury, ametal, and the nonmetal halogen bromine. Itshould be noted that the metal gallium becomesliquid at just 85.6°F (29.76°C); below that tem-perature, however, it— like the elements otherthan those named in this paragraph—is a solid.

Chemical Names and Symbols

For the sake of space and convenience, elementsare listed on the periodic table by chemical sym-bol or element symbol—a one-or two-letterabbreviation for the name of the element accord-ing to the system first developed by Berzelius.These symbols, which are standardized andunvarying for any particular element, greatly aidthe chemist in writing out chemical formulas,which could otherwise be quite cumbersome.

Many of the chemical symbols are simpleone-letter designations: H for hydrogen, O foroxygen, and F for fluorine. Others are two-letterabbreviations, such as He for helium, Ne forneon, and Si for silicon. Note that the first letteris always capitalized, and the second is alwayslowercase. In many cases, the two-letter symbolsindicate the first and second letters of the ele-ment’s name, but this is not nearly always thecase. Cadmium, for example, is abbreviated Cd,while platinum is Pt.

Many of the one-letter symbols indicate ele-ments discovered early in history. For instance,carbon is represented by C, and later “C” ele-ments took two-letter designations: Ce for ceri-um, Cr for chromium, and so on. Likewise, kryp-ton had to take the symbol Kr because potassiumhad already been assigned K. The association ofpotassium with K brings up one of the aspects ofchemical symbols most confusing to studentsjust beginning to learn about the periodic table:why K and not P? The latter had in fact alreadybeen taken by phosphorus, but then why not Po,assigned many years later instead to polonium?

CHEMICAL SYMBOLS BASED IN

OTHER LANGUAGES. In fact, potassi-um’s symbol is one of the more unusual exam-ples of a chemical symbol, taken from an ancientor non-European language. Soon after its discov-ery in the early nineteenth century, the elementwas named kalium, apparently after the Arabicqali or “alkali.” Hence, though it is known aspotassium today, the old symbol still stands.



PeriodicTable of

Elements

The use of Arabic in naming potassium isunusual in the sense that “strange” chemical sym-bols usually refer to Latin and Greek names.Latin names include aurum, or “shining dawn”for gold, symbolized as Au; or ferrum, the Latinword for iron, designated Fe. Likewise, lead (Pb)and sodium (Na) are represented by letters fromtheir Latin names, plumbum and natrium,respectively.

Some chemical elements are named forGreek or German words describing properties ofthe element. Consider, for instance, the halogens,collectively named for a Greek term meaning“salt producing.” Chloros, in Greek, describes asickly yellow color, and was assigned to chlorine;the name of bromine comes from a Greek wordmeaning “stink”; and that of iodine is a form of aGreek term meaning “violet-colored.” Astatine,last-discovered of the halogens and the rarest ofall natural elements, is so radioactive that it wasgiven a name meaning “unstable.” AnotherGreek-based example outside the halogen familyis phosphorus, or “I bring light”—appropriateenough, in view of its phosphorescent properties.

NAMES OF LATER ELEMENTS.

The names of several elements with high atomicnumbers—specifically, the lanthanides, thetransuranium elements of the actinide series, andsome of the later transition metals—have a num-ber of interesting characteristics. Several reflectthe places where they were originally discoveredor created: for example, germanium, americium,and californium. Other elements are named forfamous or not-so-famous scientists. Most peoplecould recognize einsteinium as being namedafter Albert Einstein (1879-1955), but the originof the name gadolinium—Finnish chemist JohanGadolin (1760-1852)—is harder for the averageperson to identify. Then of course there is ele-ment 101, named mendelevium in honor of theman who created the periodic table.

Two elements are named after women: curi-um after French physicist and chemist MarieCurie (1867-1934), and meitnerium after Austri-an physicist Lise Meitner (1878-1968). Curie, thefirst scientist to receive two Nobel Prizes—inboth physics and chemistry—herself discoveredtwo elements, radium and polonium. In keepingwith the trend of naming transuranium elementsafter places, she commemorated the land of herbirth, Poland, in the name of polonium. One ofCurie’s students, French physicist Marguerite

Perey (1909-1975), also discovered an elementand named it after her own homeland: francium.

Meitnerium, the last element to receive aname, was created in 1982 at the Gesellschaft fürSchwerionenforschung, or GSI, in Darmstadt,Germany, one of the world’s three leading centersof research involving transuranium elements.The other two are the Joint Institute for NuclearResearch in Dubna, Russia, and the University ofCalifornia at Berkeley, for which berkelium isnamed.

THE IUPAC AND THE NAMING

OF ELEMENTS. One of the researchersinvolved with creating berkelium was Americannuclear chemist Glenn T. Seaborg (1912-1999),who discovered plutonium and several othertransuranium elements. In light of his many con-tributions, the scientists who created element 106at Dubna in 1974 proposed that it be namedseaborgium, and duly submitted the name to theInternational Union of Pure and Applied Chem-istry (IUPAC).

Founded in 1919, the IUPAC is, as its namesuggests, an international body, and it oversees anumber of matters relating to the periodic table:the naming of elements, the assignment of chem-ical symbols to new elements, and the certifica-tion of a particular research team as the discover-ers of that element. For many years, the IUPACrefused to recognize the name seaborgium,maintaining that an element could not be namedafter a living person. The dispute over the ele-ment’s name was not resolved until the 1990s,but finally the IUPAC approved the name, andtoday seaborgium is included on the internation-al body’s official list.

Elements 110 through 112 had yet to benamed in 2001, and hence were still designatedby the three-letter symbols Uun, Uuu, and Uubrespectively. These are not names, but alphabeticrepresentations of numbers: un for 1, nil for 0,and bium for 2. Thus, the names are rendered asununnilium, unununium, and ununbium; theundiscovered elements 114, 116, and 118 arerespectively known as ununquadium, ununhexi-um, and ununoctium.

Layout of the Periodic Table

TWO SYSTEMS FOR LABELING

GROUPS. Having discussed the three items of




information contained in the boxes of the peri-odic table—atomic number, chemical symbol/name, and average atomic mass—it is now possi-ble to step back from the chart and look at itsoverall layout. To reiterate what was stated in theintroduction to the periodic table above, thetable is arranged in rows called periods, andcolumns known as groups. The deeper meaningof the periods and groups, however—that is, theway that chemists now understand them in lightof what they know about electron configura-tions—will require some explanation.

All current versions of the periodic tableshow seven rows—in other words, seven peri-ods—as well as 18 columns. However, the meansby which columns are assigned group numbersvaries somewhat. According to the system used inNorth America, only eight groups are numbered.These are the two “tall” columns on the left sideof the “dip” in the chart, as well as the six “tall”columns to the right of it. The “dip,” which spans10 columns in periods 4 through 7, is the regionin which the transition metals are listed. TheNorth American system assigns no group num-bers to these, or to the two rows set aside at thebottom, representing the lanthanide and actinideseries of transition metals.

As for the columns that the North Americansystem does number, this numbering may appearin one of four forms: either by Roman numerals;Roman numerals with the letter A (for example,IIIA); Hindu-Arabic numbers (for example, 3);or Hindu-Arabic numerals with the letter A.Throughout this book, the North American sys-tem of assigning Hindu-Arabic numerals with-out the letter A has been used. However, anattempt has been made in some places to includethe group designation approved by the IUPAC,which is used by scientists in Europe and mostparts of the world outside of North America.(Some scientists in North America are alsoadopting the IUPAC system.)

The IUPAC numbers all columns on thechart, so that instead of eight groups, there are18. The table below provides a means of compar-ing the North American and IUPAC systems.Columns are designated in terms of the elementfamily or families, followed in parentheses by theatomic numbers of the elements that appear atthe top and bottom of that column. The firstnumber following the colon is the number in theNorth American system (as described above, a

Hindu-Arabic numerical without an “A”), andthe second is the number in the IUPAC system.

Valence Electrons, Periods,and Groups

The merits of the IUPAC system are easy enoughto see: just as there are 18 columns, the IUPAClists 18 groups. Yet the North American system ismore useful than it might seem: the group num-ber in the North American system indicates thenumber of valence electrons, the electrons thatare involved in chemical bonding. Valence elec-trons also occupy the highest energy level in theatom—which might be thought of as the orbitfarthest from the nucleus, though in fact the real-ity is more complex.

A more detailed, though certainly far fromcomprehensive, discussion of electrons and ener-gy levels, as well as the history behind these dis-coveries, appears in the Electrons essay. In whatfollows, the basics of electron configuration willbe presented with the specific aim of making itclear exactly why elements appear in particularcolumns of the periodic table.

Element

Family

North

American IUPAC

Alkaline metals (4, 88)

Transition metals (21,89)

1

2

1

2

3

Hydrogen and

alkali metals (1, 87)

Halogens (9,85)

Noble gases (2,86)

Lanthanides

(58,71)

Actinides (90,103)

No number

group

assigned

in either

system

No number

group

assigned

in either

system

Transition metals (22,104) 4









Nonmetals and metals

(5,81)3 13

Nonmetals, metalloids,

and metal (6,82)4 14


and metal (7,83)5 15


(8,84)6

7

16

17

8 18



PeriodicTable of

Elements

PRINCIPAL ENERGY LEVELS

AND PERIODS. At one time, scientiststhought that electrons moved around a nucleusin regular orbits, like planets around the Sun. Infact the paths of an electron are much more com-plicated, and can only be loosely defined in termsof orbitals, a set of probabilities regarding thepositions that an electron is likely to occupy as itmoves around the nucleus. The pattern oforbitals is determined by the principal energylevel of the atom, which indicates a distance thatan electron may move away from the nucleus.

Principal energy level is designated by awhole-number integer, beginning with 1 andmoving upward: the higher the number, the fur-ther the electron is from the nucleus, and hencethe greater the energy in the atom. Each principalenergy level is divided into sublevels correspon-ding to the number n of the principal energylevel: thus, principal energy level 1 has one sub-level, principal energy level 2 has two, and so on.

The relationship between principal energylevel and period is relatively easy to demonstrate:the number n of a period on the periodic table isthe same as the number of the highest principalenergy level for the atoms on that row—that is,the principal energy level occupied by its valenceelectrons. Thus, elements on period 4 have ahighest principal energy level of 4, whereas thevalence electrons of elements on period 7 are atprincipal energy level 7. Note the conclusion thatthis allows us to draw: the further down the peri-odic table an element is positioned, the greaterthe energy in a single atom of that element. Notsurprisingly, most of the elements used innuclear power come from period 7, whichincludes the actinides.

VALENCE ELECTRON CONFIGU-

RATIONS AND GROUPS. Now to amore involved subject, whereby group number isrelated to valence electron configuration. Asmentioned earlier, the principal energy levels aredivided into sublevels, which are equal in num-ber to the principal energy level number: princi-pal energy level 1 has one sublevel, level 2 has twosublevels, and so on. As one might expect, withan increase in principal energy levels and sub-levels, there are increases in the complexity of theorbitals.

The four types of orbital patterns are desig-nated as s, p,d, and f. Two electrons can move inan s orbital pattern or shell, six in a p, 10 in a d,

and 14 in an f orbital pattern or shell. This saysnothing about the number of electrons that areactually in a particular atom; rather, the higherthe principal energy level and the larger the num-ber of sublevels, the greater the number of waysthat the electrons can move. It does happen to bethe case, however, that with higher atomic num-bers—which means more electrons to offset theprotons—the higher the energy level, the largerthe number of orbitals for those electrons.

Let us now consider a few examples ofvalence shell configurations. Hydrogen, with thesimplest of all atomic structures, has just oneelectron on principal energy level 1, so in effectits valence electron is also a core electron. Thevalence configuration for hydrogen is thus writ-ten as 1s1. Moving straight down the periodictable to francium (atomic number 87), which isin the same column as hydrogen, one finds that ithas a valence electron configuration of 7s1. Thus,although francium is vastly more complex andenergy-filled than hydrogen, the two elementshave the same valence-shell configuration;only the number of the principal energy level isdifferent.

Now look at two elements in Group 3(Group 13 in the IUPAC system): boron andthallium, which respectively occupy the top andbottom of the column, with atomic numbers of 5and 81. Boron has a valence-shell configurationof 2s22p1. This means its valence shell is at prin-cipal energy level 2, where there are two electronsin an s orbital pattern, and 2 in a p orbital pat-tern. Thallium, though it is on period 6, nonethe-less has the same valence-shell configuration:6s26p1.

Notice something about the total of thesuperscript figures for any element in Group 3 ofthe North American system: it is three. The sameis true in the other columns numbered on NorthAmerican charts, in which the total number ofelectrons equals the group number. Thus inGroup 7, the valence shell configuration isns2np5, where n is the principal energy level.There is only one exception to this: helium, inGroup 8 (the noble gases), has a valence shellconfiguration of 1s2. Were it not for the fact thatit clearly fits with the noble gases due to sharedproperties, helium would be placed next tohydrogen at the top of Group 2, where all theatoms have a valence-shell configuration of ns2.




Obviously the group numbers in the IUPACsystem do not correspond to the number ofvalence electrons, because the IUPAC chartincludes numbers for the columns of transitionmetals, which are not numbered in the NorthAmerican system. In any case, in both systemsthe columns contain elements that all have thesame number of electrons in their valence shells.Thus the term “group” can finally be defined inaccordance with modern chemists’ understand-ing, which incorporates electron configurationsof which Mendeleev was unaware. All the mem-bers of a group have the same number of valenceelectrons in the same orbital patterns, though atdifferent energy levels. (Once again, helium is thelone exception.)

Some Challenges of thePeriodic Table

IRREGULAR PATTERNS. Thegroups that are numbered in the North Americansystem are referred to as “representative” ele-ments, because they follow a clearly establishedpattern of adding valence shell electrons. By con-trast, the 40 elements listed in the “dip” at themiddle of the chart—the transition elements—do not follow such a pattern. This is why theNorth American system does not list them bygroup number, and also why neither system liststwo “branches” of the transition-metal family,the lanthanides and actinides.

Even within the representative elements,there are some challenges as far as electron con-figuration. For the first 18 elements—1 (hydro-gen) to 18 (argon)—there is a regular pattern oforbital filling. Beginning with helium (2)onward, all of principal level 1 is filled; then,beginning with beryllium (4), sublevel 2s beginsto fill. Sublevel 2p—and hence principal level 2 asa whole—becomes filled at neon (10).

After argon, as one moves to the elementoccupying the nineteenth position on the peri-odic table—potassium—the rules change.Argon, in Group 8 of the North American sys-tem, has a valence shell of 3s23p6, and by the pat-tern established with the first 18 elements, potas-sium should begin filling principal level 3d.Instead, it “skips” 3d and moves on to 4s. The ele-ment following argon, calcium, adds a secondelectron to the 4s sublevel.

After calcium, as the transition metals beginwith scandium (21), the pattern again changes:

indeed, the transition elements are defined by thefact that they fill the d orbitals rather than the porbitals, as was the pattern up to that point. Afterthe first period of transition metals ends withzinc (30), the next representative element—galli-um (31)—resumes the filling of the p orbitalrather than the d. And so it goes, all along thefour periods in which transition metals break upthe steady order of electron configurations.

As for the lanthanide and actinide series oftransitions metals, they follow an even moreunusual pattern, which is why they are set aparteven from the transition metals. These are theonly groups of elements that involve the highlycomplex f sublevels. In the lanthanide series, theseven 4f orbital shells are filled, while the actinideseries reflects the filling of the seven 5f orbitalshells.

Why these irregularities? One reason is thatas the principal energy level increases, the energylevels themselves become closer—i.e., there is lessdifference between the energy levels. The atom isthus like a bus that fills up: when there are just afew people on board, those few people (analo-gous to electrons) have plenty of room, but asmore people get on, the bus becomes increasing-ly more crowded, and passengers jostle againstone another. In the atom, due to differences inenergy levels, the 4s orbital actually has a lowerenergy than the 3d, and therefore begins to fillfirst. This is also true for the 6s and 4f orbitals.

CHANGES IN ATOMIC SIZE. Thesubject of element families is a matter unto itself,and therefore a separate essay in this book hasbeen devoted to it. The reader is encouraged toconsult the Families of Elements essay, whichdiscusses aspects of electron configuration aswell as the properties of various element families.

One last thing should be mentioned aboutthe periodic table: the curious fact that the sizesof atoms decreases as one moves from left toright across a row or period, even though thesizes increase as one moves from top to bottomalong a group. The increase of atomic size in agroup, as a function of increasing atomic num-ber, is easy enough to explain. The higher theatomic number, the higher the principal energylevel, and the greater the distance from the nucle-us to the furthest probability range for the electron.

On the other hand, the decrease in sizeacross a period is a bit more challenging to com-



PeriodicTable of

Elements




physical properties of the element.












a large sample.



gadro (1776-1856), equal to 6.022137 x


number of atoms or molecules in a mole.

CHEMICAL SYMBOL: A one- or two-


element.

COMPOUND: A substance made of two

or more elements that have bonded chem-

ically. These atoms are usually, but not

always, joined in molecules.


cle in an atom. The configurations of

valence electrons define specific groups on

the periodic table of elements, while the

principal energy levels of those valence

electrons define periods on the table.


one kind of atom, which cannot be chemi-

cally broken into other substances.

ELEMENT SYMBOL: Another term for

chemical symbol.

GROUPS: Columns on the periodic

table of elements. These are ordered

according to the numbers of valence elec-

trons in the outer shells of the atoms for

the elements represented.



amount.


gained one or more electrons, thus acquir-

ing a net electric charge.






unstable. The latter type, known as

radioisotopes, are radioactive.









but not always representing more than one

element, joined by chemical bonds. Com-

pounds are typically made of up mole-

cules.


has no electric charge. Neutrons, together

with protons, account for the majority of

average atomic mass. When atoms have the

same number of protons—and hence are

the same element—but differ in their

number of neutrons, they are called

isotopes.

KEY TERMS




prehend; however, it just takes a little explaining.

As one moves along a period from left to right,

there is a corresponding increase in the number

of protons within the nucleus. This means a

stronger positive charge pulling the electrons

inward. Therefore, the “cloud” of electrons is

drawn ever closer toward the increasingly power-

ful charge at the center of the atom, and the size

of the atom decreases because the electrons can-

not move as far away from the nucleus.

WHERE TO LEARN MORE

Challoner, Jack. The Visual Dictionary of Chemistry. New

York: DK Publishing, 1996.

“Elementistory” (Web site). <http://smallfry.dmu.ac.uk/chem/periodic/elementi.html> (May 22, 2001).

International Union of Pure and Applied Chemistry (Website). <http://www.iupac.org> (May 22, 2001).



“A Periodic Table of the Elements” Los Alamos NationalLaboratory (Web site). <http://pearl1.lanl.gov/periodic/> (May 22, 2001).

“The Pictorial Periodic Table” (Web site). <http://chem-lab.pc.maricopa.edu/periodic/periodic.html> (May22, 2001).




located. The nucleus accounts for the vast

majority of the average atomic mass.


regarding the regions that an electron can

occupy within an atom in a particular

energy state. The higher the principal ener-

gy level, the more complex the pattern of

orbitals.






element.

PERIODS: Rows of the periodic table of

elements. These represent successive prin-

cipal energy levels for the valence electrons

in the atoms of the elements involved.






upward. The higher the principal energy

level, the greater the energy in the atom,

and the more complex the pattern of

orbitals.


in an atom. The number of protons in the

nucleus of an atom is the atomic number

of an element.







decays to form another isotope— either of

the same element or another—until even-

tually it becomes stable. This stabilizing

process may take a few seconds, or many

years.



atom. These are the electrons involved in

chemical bonding.

KEY TERMS CONTINUED


PeriodicTable of

Elements

“Visual Elements” (Web site). <http://www.chemsoc.org/viselements/> (May 22, 2001).

WebElements (Web site). <http://www.webelements.com(May 22, 2001).




FAMILIES OF ELEMENTSFamilies of Elements

CONCEPTThe term “family” is used to describe elementsthat share certain characteristics—not only interms of observable behavior, but also withregard to atomic structure. All noble gases, forinstance, tend to be highly nonreactive: only afew of them combine with other elements, andthen only with fluorine, the most reactive of allsubstances. Fluorine is a member of anotherfamily, the halogens, which have so many sharedcharacteristics that they are grouped together,despite the fact that two are gases, two are solids,and one—bromine—is one of only two elementsthat appears at room temperature as a solid.Despite these apparent differences, commonelectron configurations identify the halogens as afamily. Families on the periodic table include, inaddition to noble gases and halogens, the alkalimetals, alkaline earth metals, transition metals,lanthanides, and actinides. The nonmetals form aloosely defined cross-family grouping, as do themetalloids.

HOW IT WORKS

The Basics of the PeriodicTable

Created in 1869, and modified several times sincethen, the periodic table of the elements devel-oped by Russian chemist Dmitri IvanovitchMendeleev (1834-1907) provides a highly usefulmeans of organizing the elements. Certainlyother organizational systems exist, butMendeleev’s table is the most widely used—andwith good reason. For one thing, it makes it pos-sible to see at a glance families of elements, many

of which either belong to the same group (col-umn) or the same period (row) on the table.

The periodic table is examined in depthwithin the essay devoted to that subject, andamong the specifics discussed in that essay arethe differing systems used for periodic-tablecharts in North America and the rest of theworld. In particular, the North American systemnumbers only eight groups, leaving 10 columnsunnumbered, whereas the other system—approved by the International Union of Pure andApplied Chemistry (IUPAC)—numbers all 18columns. Both versions of the periodic tableshow seven periods.

The groups numbered in the North Ameri-can system are the two “tall” columns on the leftside of the “dip” in the chart, as well as the six“tall” columns to the right of it. Group 1 in thissystem consists of hydrogen and the alkali met-als; Group 2, the alkaline earth metals; groups 3through 6, an assortment of metals, nonmetals,and metalloids; Group 7, halogens; and Group 8,noble gases. The “dip,” which spans 10 columnsin periods 4 through 7, is the region in which thetransition metals are listed. The North Americansystem assigns no group numbers to these, or tothe two rows set aside at the bottom, representingthe lanthanide and actinide series of transitionmetals.

The IUPAC system, on the other hand, offersthe obvious convenience of providing a numberfor each column. (Note that, like its North Amer-ican counterpart, the IUPAC chart provides nocolumn numbers for the lanthanides oractinides.) Furthermore, the IUPAC has behind itthe authority of an international body, foundedin 1919, which oversees a number of matters



relating to the periodic table: the naming of ele-

ments, the assignment of chemical symbols to

new elements, and the certification of a particu-

lar individual or research team as the discoverers

of that element. For these reasons, the IUPAC

system is coming into favor among North Amer-

ican chemists as well.

Despite the international acceptance of the

IUPAC system, as well as its merits in terms of

convenience, the North American system is gen-

erally the one used in this book. The reason, in

part, is that most American schools still use this

system; furthermore, there is a reasoning behind

the assignment of numbers to only eight groups,

as will be discussed. Where necessary or appro-

priate, however, group numbers in the IUPAC

system will be provided as well.

Principal Energy Levels

Group numbers in the North American system

indicate the number of valence electrons, or the

electrons that are involved in chemical bonding.

SIZE REPRESENTATION OF THE ATOMIC RADII OF THE MAIN-GROUP ELEMENTS.

Families ofElements


Families ofElements

Valence electrons also occupy the highest energylevel in the atom—which might be thought of asthe orbit farthest from the nucleus, though infact the term “orbit” is misleading when appliedto the ways an electron moves.

Electrons do not move around the nucleusof an atom in regular orbits, like planets aroundthe Sun; rather, their paths can only be looselydefined in terms of orbitals, a pattern of proba-bilities regarding the areas through which anelectron is likely to move. The pattern of orbitalsis determined by the principal energy level of theatom, which indicates the distance an electronmay move away from the nucleus.

Principal energy level is designated by awhole-number integer, beginning with 1 andmoving upward to 7: the higher the number, thefurther the electron is from the nucleus, andhence the greater the energy in the atom. Therelationship between principal energy level andperiod is relatively easy to demonstrate. Thenumber n of a period on the periodic table is thesame as the number of the highest principalenergy level for the atoms on that row—that is,the principal energy level occupied by its valenceelectrons. Thus, elements on period 1 have ahighest principal energy level of 1, and so on.

Valence Electron Con-figurations

When discussing families of elements, however,the periods or rows on the periodic table are notas important as the groups or columns. These aredefined by the valence electron configurations, asubject more complicated than principal energylevels—though the latter requires a bit moreexplanation in order to explain electron con-figurations.

Each principal energy level is divided intosublevels corresponding to the number n of theprincipal energy level: thus, principal energy level1 has one sublevel, principal energy level 2 hastwo, and so on. As one might expect, with anincrease in principal energy levels and sublevels,there are increases in the complexity of theorbitals.

ORBITAL PATTERNS. The four basictypes of orbital patterns are designated as s, p,d,and f. The s shape might be described as spheri-cal, though when talking about electrons, noth-ing is quite so neat: orbital patterns, remember,only identify regions of probability for the elec-

tron. In other words, in an s orbital, the total elec-tron cloud will probably end up being more orless like a sphere.

The p shape is like a figure eight around thenucleus, and the d like two figure eights meetingat the nucleus. Again, these and other orbital pat-terns do not indicate that the electron will neces-sary follow that path. What it means is that, ifyou could take millions of photographs of theelectron during a period of a few seconds, theresulting blur of images in a p orbital wouldsomewhat describe the shape of a figure eight.

The f orbital pattern is so complex that mostbasic chemistry textbooks do not even attempt toexplain it, and beyond f are other, even morecomplicated, patterns designated in alphabeticalorder: g, h, and so on. In the discussion that fol-lows, we will not be concerned with these, sinceeven for the lanthanides and the actinides, anatom at the ground state does not fill orbital pat-terns beyond an f.

SUBLEVELS AND ORBITAL

FILLING. Principal energy level 1 has only ans sublevel; 2 has an s and a p, the latter with threepossible orientations in space; 3 has an s, p, and d(five possible spatial orientations); and 4 has an s,p, d, and f (seven possible spatial orientations.)

According to the Pauli exclusion principle,only two electrons can occupy a single orbitalpattern—that is, the s sublevel or any one of thespatial orientations in p, d, and f—and those twoelectrons must be spinning in opposite direc-tions. Thus, two electrons can move in an sorbital pattern or shell, six in a p, 10 in a d, and14 in an f orbital pattern or shell. Valence shellconfigurations are therefore presented withsuperscript figures indicating the number ofelectrons in that orbital pattern—for instance, s1

for one electron in the s orbital, or d10, indicatinga d orbital that has been completely filled.


Representative Elements

Hydrogen (atomic number 1), with the simplestof all atomic structures, has just one electron onprincipal energy level 1, so, in effect, its valenceelectron is also a core electron. The valence con-figuration for hydrogen is thus written as 1s1. Itshould be noted, as described in the Electrons



Families ofElements

essay, that if a hydrogen atom (or any otheratom) is in an excited state, it may reach energylevels beyond its normal, or ground, state.

Moving straight down the periodic table tofrancium (atomic number 87), which is in thesame column as hydrogen, one finds that it has avalence electron configuration of 7s1. Thus,although francium is vastly more complex andenergy-filled than hydrogen, the two elementshave the same valence shell configuration; onlythe number of the principal energy level is differ-ent. All the elements listed below hydrogen inGroup 1 are therefore classified together as alkalimetals. Obviously, hydrogen—a gas—is not partof the alkali metal family, nor does it clearlybelong to any other family: it is the “lone wolf” ofthe periodic table.

Now look at two elements in Group 2, withberyllium (atomic number 4) and radium (88) atthe top and bottom respectively. Beryllium has avalence shell configuration of 2s2. This means itsvalence shell is at principal energy level 2, wherethere are two electrons on an s orbital pattern.Radium, though it is on period 7, nonetheless hasthe same valence shell configuration: 7s2. Thisdefines the alkaline earth metals family in termsof valence shell configuration.

For now, let us ignore groups 3 through 6—not to mention the columns between groups 2and 3, unnumbered in the North American sys-tem—and skip over to Group 7. All the elementsin this column, known as halogens, have valenceshell configurations of ns2np5. Beyond Group 7 isGroup 8, the noble gases, all but one of whomhave valence shell configurations of ns2np6. Theexception is helium, which has an s2 valence shell.This seems to put it with the alkaline earth met-als, but of course helium is not a metal. In termsof its actual behavior, it clearly belongs to thenoble gases family.

The configurations of these valence shellshave implications with regard to the ways inwhich elements bond, a subject developed atsome length in the Chemical Bonding essay. Herewe will consider it only in passing, to clarify thefact that electron configuration produces observ-able results. This is most obvious with the noblegases, which tend to resist bonding with mostother elements because they already have eightelectrons in their valence shell—the same num-ber of valence electrons that most other atomsachieve only after they have bonded.

From the Representative Elements to the TransitionElements

Groups 3 through 6, along with hydrogen andthe four families so far identified, constitute the44 representative or main-group elements. In 43of these 44, the number of valence shell electronsis the same as the group number in the NorthAmerican system. (Helium, which is in Group 8but has two valence electrons, is the lone excep-tion.) By contrast, the 40 elements listed in the“dip” at the middle of the chart—the transitionmetals—follow a less easily defined pattern. Thisis part of the reason why the North Americansystem does not list them by group number, andalso why neither system lists the two other fami-lies within the transition elements—the lan-thanides and actinides.

Before addressing the transition metals,however, let us consider patterns of orbital fill-ing, which also differentiate the representativeelements from the transition elements. Each suc-cessive representative element fills all the orbitalsof the elements that precede it (with some excep-tions that will be explained), then goes on to addone more possible electron configuration. Thetotal number of electrons—not just valence shellelectrons—is the same as the atomic number.Thus fluorine, with an atomic number of 9, has acomplete configuration of 1s22s22p5. Neon,directly following it with an atomic number of10, has a total configuration of 1s22s22p6. (Again,this is not the same as the valence shell configu-ration, which is contained in the last two sub-levels represented: for example, 2s22p6 for neon.)

The chart that follows shows the pattern bywhich orbitals are filled. Note that in severalplaces, the pattern of filling becomes “out oforder,” something that will be explained below.

Orbital Filling by Principal Energy Level

• 1s (2)

• 2s (2)

• 2p (6)

• 3s (2)

• 3p (6)

• 4s (2)

• 3d (10)

• 4p (6)

• 5s (2)

• 4d (10)

• 5p (6)



Families ofElements

• 6s (2)

• 4f (14)

• 5d (10)

• 6p (6)

• 7s (2)

• 5f (14)

• 6d (10)

PATTERNS OF ORBITAL FILL-

ING. Generally, the 44 representative elementsfollow a regular pattern of orbital filling, and thisis particularly so for the first 18 elements. Imag-ine a small amphitheater, shaped like a cone, withsmaller rows of seats at the front. These rows arealso designated by section, with the section num-ber being the same as the number of rows in thatsection.

The two seats in the front row comprise asection labeled 1 or 1s, and this is completelyfilled after helium (atomic number 2) enters theauditorium. Now the elements start filling sec-tion 2, which contains two rows. The first row ofsection 2, labeled 2s, also has two seats, and afterberyllium (4), it too is filled. Row 2p has 6 seats,and it is finally filled with the entrance of neon(10). Now, all of section 2 has been filled; there-fore, the eleventh element, sodium, starts fillingsection 3 at the first of its three rows. This row is3s—which, like all s rows, has only two seats.Thus, when element 13, aluminum, enters thetheatre, it takes a seat in row 3p, and eventuallyargon (18), completes that six-seat row.

By the pattern so far established, element 19(potassium) should begin filling row 3d by takingthe first of its 10 seats. Instead, it moves on to sec-tion 4, which has four rows, and it takes the firstseat in the first of those rows, 4s. Calcium (20)follows it, filling the 4s row. But when the nextelement, scandium (21), comes into the theatre,it goes to row 3d, where potassium “should have”gone, if it had continued filling sections in order.Scandium is followed by nine companions (thefirst row of transition elements) before anotherrepresentative element, gallium (31), comes intothe theatre. (For reasons that will not be dis-cussed here, chromium and copper, elements 24and 29, respectively, have valence electrons in4s—which puts them slightly off the transitionmetal pattern.)

According to the “proper” order of fillingseats, now that 3d (and hence all of section 3) isfilled, gallium should take a seat in 4s. But thoseseats have already been taken by the two preced-

ing representative elements, so gallium takes thefirst of six seats in 4p. After that row fills up atkrypton (36), it is again “proper” for the nextrepresentative element, rubidium (37), to take aseat in 4d. Instead, just as potassium skipped 3d,rubidium skips 4d and opens up section 5 by tak-ing the first of two seats in 5s.

Just as before, the next transition element—yttrium (39)—begins filling up section 4d, and isfollowed by nine more transition elements untilcadmium (48) fills up that section. Then, the rep-resentative elements resume with indium (49),which, like gallium, skips ahead to section 5p.And so it goes through the remainder of the periodic table, which ends with two representa-tive elements followed by the last 10 transition metals.

Transition Metals

Given the fact that it is actually the representativeelements that skip the d sublevels, and the transi-tion metals that go back and fill them, one mightwonder if the names “representative” and “transi-tion” (implying an interruption) should bereversed. However, remember the correlationbetween the number of valence shell electronsand group number for the representative ele-ments. Furthermore, the transition metals are theonly elements that fill the d orbitals.

This brings us to the reason why the lan-thanides and actinides are set apart even from thetransition metals. In most versions of the period-ic table, lanthanum (57) is followed by hafnium(72) in the transition metals section of the chart.Similarly, actinium (89) is followed by ruther-fordium (104). The “missing” metals—lan-thanides and actinides, respectively—are listed atthe bottom of the chart. There are reasons forthis, as well as for the names of these groups.

After the 6s orbital fills with the representa-tive element barium (56), lanthanum does whata transition metal does—it begins filling the 5dorbital. But after lanthanum, something strangehappens: cerium (58) quits filling 5d, and movesto fill the 4f orbital. The filling of that orbitalcontinues throughout the entire lanthanideseries, all the way to lutetium (71). Thus, lan-thanides can be defined as those metals that fillthe 4f orbital; however, because lanthanumexhibits similar properties, it is usually includedwith the lanthanides. Sometimes the term “lan-



Families ofElements


ACTINIDES: Those transition metals

that fill the 5f orbital. Because actinium—

which does not fill the 5f orbital—exhibits

characteristics similar to those of the

actinides, it is usually considered part of

the actinides family.

ALKALI METALS: All members, except

hydrogen, of Group 1 on the periodic table

of elements, with valence electron configu-

rations of ns1.

ALKALINE EARTH METALS: Group 2

on the periodic table of elements, with

valence electron configurations of ns2.



FAMILIES OF ELEMENTS: Related

elements, including the noble gases,

halogens, alkali metals, alkaline earth met-

als, transition metals,lanthanides, and

actinides. In addition, metals, nonmetals,

and metalloids form loosely defined fami-

lies. Other family designations—such as

carbon family—are sometimes used.


the state of an atom at its ordinary energy

level.






HALOGENS: Group 7 of the periodic

table of elements, with valence electron

configurations of ns2np5.




LANTHANIDES: The transition metals

that fill the 4forbital. Because lan-

thanum—which does not fill the 4f

orbital—exhibits characteristics similar to

those of the lanthanides, it is usually con-

sidered part of the lanthanide family.

MAIN-GROUP ELEMENTS: The 44

elements in Groups 1 through 8 on the

periodic table of elements, for which the

number of valence electrons equals the

group number. (The only exception is heli-

um.) The main-group elements, also called

representative elements, include the fami-

lies of alkali metals, alkali earth metals,

halogens, and noble gases, as well as other

metals, nonmetals, and metalloids.

METALLOIDS: Elements which exhibit

characteristics of both metals and non-

metals. Metalloids are all solids, but are not

lustrous or shiny, and they conduct heat

and electricity moderately well. The six

metalloids occupy a diagonal region

between the metals and nonmetals on the

right side of the periodic table. Sometimes

astatine is included with the metalloids,

but in this book it is treated within the

context of the halogens family.

METALS: A collection of 87 elements

that includes numerous families—the

alkali metals, alkaline earth metals, transi-

tion metals, lanthanides, and actinides, as

well as seven elements in groups 3 through

5. Metals, which occupy the left, center, and

part of the right-hand side of the periodic

table, are lustrous or shiny in appearance,

and malleable, meaning that they can be

molded into different shapes without

breaking. They are excellent conductors of

heat and electricity, and tend to form posi-

tive ions by losing electrons.

KEY TERMS


Families ofElements


thanide series” is used to distinguish the other 14

lanthanides from lanthanum itself.

A similar pattern occurs for the actinides.

The 7s orbital fills with radium (88), after which

actinium (89) begins filling the 6d orbital. Next

comes thorium, first of the actinides, which

begins the filling of the 5f orbital. This is com-

pleted with element 103, lawrencium. Actinides

can thus be defined as those metals that fill the 5f

orbital; but again, because actinium exhibits sim-

ilar properties, it is usually included with the

actinides.

Metals, Nonmetals, and Metalloids

The reader will note that for the seven families so

far identified, we have generally not discussed

them in terms of properties that can more easily

be discerned—such as color, phase of matter,

bonding characteristics, and so on. Instead, they

have been examined primarily from the stand-

point of orbital filling, which provides a solid

chemical foundation for identifying families.

Macroscopic characteristics, as well as the ways

that the various elements find application in

NOBLE GASES: Group 8 of the peri-

odic table of elements, all of whom (with

the exception of helium) have valence elec-

tron configurations of ns2np6.

NONMETALS: Elements that have a dull

appearance; are not malleable; are poor

conductors of heat and electricity; and

tend to gain electrons to form negative

ions. They are thus the opposite of metals

in most regards, as befits their name. Aside

from hydrogen, the other 18 nonmetals

occupy the upper right-hand side of the

periodic table, and include the noble gases,

halogens, and seven elements in groups 3

through 6.



atom in a particular energy state. The high-

er the principal energy level, the more

complex the pattern of orbitals. The four

types of orbital patterns are designated as s,

p, d, and f—each of which is more complex

than the one before.






element.


elements. These represent successive ener-

gy levels in the atoms of the elements

involved.









orbitals.

REPRESENTATIVE ELEMENTS: See

main-group elements.

TRANSITION METALS: A group of 40

elements, which are not assigned a group

number in the North American version of

the periodic table. These are the only ele-

ments that fill the d orbitals.



atom. These are the electrons involved in

chemical bonding.

KEY TERMS CONTINUED


Families ofElements

daily life, are discussed within essays devoted tothe various groups.

Note, also, that the families so far identifiedaccount for only 92 elements out of a total of 112listed on the periodic table: hydrogen; six alkalimetals; six alkaline earth metals; five halogens;six noble gases; 40 transition metals; 14 lan-thanides; and 14 actinides. What about the other20? Some discussions of element families assignthese elements, all of which are in groups 3through 6, to families of their own, which will bementioned briefly. However, because these “fam-ilies” are not recognized by all chemists, in thisbook the 20 elements of groups 3 through 6 aredescribed generally as metals, nonmetals, andmetalloids.

METALS AND NONMETALS. Met-als are lustrous or shiny in appearance, and mal-leable, meaning that they can be molded into dif-ferent shapes without breaking. They are excel-lent conductors of heat and electricity, and tendto form positive ions by losing electrons. On theperiodic table, metals fill the left, center, and partof the right-hand side of the chart. Thus it shouldnot come as a surprise that most elements (87, infact) are metals. This list includes alkali metals,alkaline earth metals, transition metals, lan-thanides, and actinides, as well as seven elementsin groups 3 through 6—aluminum, gallium,indium, thallium, tin, lead, and bismuth.

Nonmetals have a dull appearance; are notmalleable; are poor conductors of heat and elec-tricity; and tend to gain electrons to form nega-tive ions. They are thus the opposite of metals inmost regards, as befits their name. Nonmetals,which occupy the upper right-hand side of theperiodic table, include the noble gases, halogens,and seven elements in groups 3 through 5. Thesenonmetal “orphans” are boron, carbon, nitrogen,oxygen, phosphorus, sulfur, and selenium. Tothese seven orphans could be added an eighth,from Group 1: hydrogen. As with the metals, aseparate essay—with a special focus on the“orphans”—is devoted to nonmetals.

METALLOIDS AND OTHER

“FAMILY” DESIGNATIONS. Occupyinga diagonal region between the metals and non-metals are metalloids, elements which exhibitcharacteristics of both metals and nonmetals.They are all solids, but are not lustrous, and con-duct heat and electricity moderately well. The sixmetalloids are silicon, germanium, arsenic, anti-mony, tellurium, and polonium. Astatine issometimes identified as a seventh metalloid;however, in this book, it is treated as a member ofthe halogen family.

Some sources list “families” rather than col-lections of “orphan” metals, metalloids, and non-metals, in groups 3 through 6. These designa-tions are not used in this book; however, theyshould be mentioned briefly. Group 3 is some-times called the boron family; Group 4, the car-bon family; Group 5, the nitrogen family; andGroup 6, the oxygen family. Sometimes Group 5is designated as the pnictogens, and Group 6 asthe chalcogens.

WHERE TO LEARN MORE

Bankston, Sandy. “Explore the Periodic Table and Familiesof Elements” The Rice School Science Department(Web site). <http://www.ruf.rice.edu/~sandyb/Lessons/chem.html> (May 23, 2001).


“Elementistory” (Web site). <http://smallfry.dmu.ac.uk/chem/periodic/elementi.html> (May 22, 2001).

“Families of Elements” (Web site). <http://homepages.stuy.edu/~bucherd/ch23/families.html> (May 23,2001).


Maton, Anthea. Exploring Physical Science. Upper SaddleRiver, N.J.: Prentice Hall, 1997.


“The Pictorial Periodic Table” (Web site). <http://chemlab.pc.maricopa.edu/periodic/periodic.html>(May 22, 2001).


“Visual Elements” (Web site). <http://www.chemsoc.org/viselements/> (May 22, 2001).



149

SCIENCE OF EVERYDAY THINGSreal-l ife CHEMISTRY

METALSMETALS

METALS

ALKALI METALS

ALKALINE EARTH METALS

TRANSITION METALS

ACTINIDES

LANTHANIDES



METALSMetals

CONCEPTA number of characteristics distinguish metals,including their shiny appearance, as well as theirability to be bent into various shapes withoutbreaking. In addition, metals tend to be highlyefficient conductors of heat and electricity. Thevast majority of elements on the periodic tableare metals, and most of these fall into one of fivefamilies: alkali metals; alkaline earth metals; thevery large transition metals family; and the innertransition metal families, known as the lan-thanides and actinides. In addition, there areseven “orphans,” or metals that do not belong toa larger family of elements: aluminum, gallium,indium, thallium, tin, lead, and bismuth. Some ofthese may be unfamiliar to most readers; on theother hand, a person can hardly spend a daywithout coming into contact with aluminum.Tin is a well-known element as well, thoughmuch of what people call “tin” is not really tin.Lead, too, is widely known, and once was widelyused as well; today, however, it is known primari-ly for the dangers it poses to human health.

HOW IT WORKS

General Properties of Metals

Metals are lustrous or shiny in appearance, andmalleable or ductile, meaning that they can bemolded into different shapes without breaking.Despite their ductility, metals are extremelydurable and have high melting and boilingpoints. They are excellent conductors of heat andelectricity, and tend to form positive ions by los-ing electrons.

The bonds that metals form with each other,or with nonmetals, are known as ionic bonds,which are the strongest type of chemical bond.Even within a metal, however, there are extreme-ly strong bonds. Therefore, though it is easy toshape metals (that is, to slide the atoms in a metalpast one another), it is very difficult to separatemetal atoms. Their internal bonding is thus verystrong, but nondirectional.

Internal bonding in metals is described bythe electron sea model, which depicts metalatoms as floating in a “sea” of valence electrons,the electrons involved in bonding. These valenceelectrons are highly mobile within the crystallinestructure of the metal, and this mobility helps toexplain metals’ high conductivity. The ease withwhich metal crystals allow themselves to berearranged explains not only metals’ ductility,but also their ability to form alloys, a mixturecontaining one or more metals.

Abundance of Metals

Metals constitute a significant portion of the ele-ments found in Earth’s crust, waters, and atmos-phere.

The following list shows the most abundantmetals on Earth, with their ranking among allelements in terms of abundance. Many othermetals are present in smaller, though still signifi-cant, quantities.

• 3. Aluminum (7.5%)

• 4. Iron (4.71%)

• 5. Calcium (3.39%)

• 6. Sodium (2.63%)

• 7. Potassium (2.4%)

• 8. Magnesium (1.93%)


Metals


• 10. Titanium (0.58%)

• 13. Manganese (0.09%)

In addition, metals are also important com-ponents of the human body. The following listshows the ranking of the most abundant metalsamong the elements in the body, along with thepercentage that each constitutes. As with thequantities of metals on Earth, other metals arepresent in much smaller, but still critical,amounts.

• 5. Calcium (1.4%)

• 7. Magnesium (0.50%)

• 8. Potassium (0.34%)

• 10. Sodium (0.14%)

• 12. Iron (0.004%)

• 13. Zinc (0.003%)

Metals on the Periodic Table

On the periodic table of elements, metals fill theleft, center, and part of the right-hand side of thechart. The remainder of the right-hand section istaken by nonmetals, which have properties quitedifferent from (and often opposite to) those ofmetals, as well as metalloids, which display char-acteristics both of metals and nonmetals. Non-metallic elements are covered in essays on Non-metals; Metalloids; Halogens; Noble Gases; andCarbon. In addition, hydrogen, the one nonmetalon the left side of the periodic table—first amongthe elements, in the upper left corner—is alsodiscussed in its own essay.

The total number of metals on the periodictable is 87—in other words, about 80% of all

ALUMINUM IS THE MOST ABUNDANT METAL ON EARTH. THIS PHOTO SHOWS THE FIRST BLOBS OF ALUMINUM EVER

ISOLATED, A FEAT PERFORMED BY CHARLES HALL IN 1886. (James L. Amos/Corbis. Reproduced by permission.)


known elements. These include the six alkalimetals, which occupy Group 1 of the periodictable, as well as the six alkaline earth metals, orGroup 2. To the right of these are the 40 transi-tion metals, whose groups are unnumbered inthe North American version of the periodic table.Among the transition metals are two elements,lanthanum and actinium, often lumped in withthe families of inner transition metals that exhib-it similar properties. These are, respectively, the14 lanthanides and 14 actinides. In addition,there are the seven “orphan” metals of periods 3through 5, mentioned in the introduction to thisessay, which will be discussed below.

The various families of elements are defined,not so much by external characteristics as by theconfigurations of valence electrons in theiratoms’ shells. This subject will not be discussed inany depth here; instead, the reader is encouragedto consult essays on the specific metal families,which discuss the electron configurations ofeach.


Alkali Metals

The members of the alkali metal family are dis-tinguished by the fact that they have a valenceelectron configuration of s1 This means that theytend to bond easily with other substances. Alkalimetals tend to form positive ions, or cations, witha charge of 1+.

In contact with water, alkali metals form anegatively charged hydroxide ion (OH–). Whenalkali metals react with water, one hydrogenatom splits off from the water molecule to formhydrogen gas, while the other hydrogen atomremains bonded to the oxygen, forming hydrox-ide. Where the heavier members of the alkalimetal family are concerned, reactions can oftenbe so vigorous that the result is combustion oreven explosion. Alkali metals also react with oxy-gen to produce either an oxide, peroxide, orsuperoxide, depending on the particular memberof the alkali metal family involved.

Shiny and soft enough to be cut with a knife,the alkali metals are usually white, thoughcesium is a yellowish white. When placed in aflame, most of these substances produce charac-


teristic colors: lithium, for instance, glows brightred, and sodium an intense yellow. Heated potas-sium produces a violet color, rubidium a darkred, and cesium a light blue. This makes it possi-ble to identify the metals by color when heated—a useful trait, since they so often tend to be bond-ed with other elements.

The alkali metals, listed below by atomicnumber and chemical symbol, are:

• 3. Lithium (Li)

• 11. Sodium (Na)

• 19. Potassium (K)

• 37. Rubidium (Rb)

• 55. Cesium (Cs)

• 87. Francium (Fr)

Alkaline Earth Metals

Occupying Group 2 of the periodic table are thealkaline earth metals, which have a valence elec-tron configuration of s2. Like the alkali metals,the alkaline earth metals are known for their highreactivity—a tendency for bonds between atomsor molecules to be made or broken in such a waythat materials are transformed. Thus they are sel-

BEFORE THEY WERE DESTROYED BY TERRORIST ATTACKS

ON SEPTEMBER 11, 2001, THE TWIN TOWERS OF THE

WORLD TRADE CENTER IN NEW YORK CITY GLEAMED

WITH A CORROSION-RESISTANT COVERING OF ALUMINUM.(Bill Ross/Corbis. Reproduced by permission.)

Metals


Metals


dom found in pure form, but rather in com-pounds with other elements—for example, salts,a term that generally describes a compound con-sisting of a metal and a nonmetal.

The alkaline earth metals are also like thealkali metals in that they have the properties of abase as opposed to an acid. An acid is a substancethat, when dissolved in water, produces positiveions of hydrogen, designated symbolically as H+

ions. It is thus known as a proton donor. A base,on the other hand, produces negative hydroxideions when dissolved in water. These are designat-ed by the symbol OH-. Bases are therefore char-acterized as proton acceptors.

The alkaline earth metals are shiny, andmost are white or silvery in color. Like their“cousins” in the alkali metal family, they glowwith characteristic colors when heated. Calciumglows orange, strontium a very bright red, andbarium an apple green. Physically they are soft,though not as soft as the alkali metals; nor aretheir levels of reactivity as great as those of theirneighbors in Group 1.

The alkaline earth metals, listed below byatomic number and chemical symbol, are:

• 4. Beryllium (Be)

• 12. Magnesium (Mg)

• 20. Cadmium (Ca)

• 38. Strontium (Sr)

• 56. Barium (Ba)

• 88. Radium (Ra)

Transition Metals

The transition metals are the only elements that

fill the dorbitals. They are also the only elements

that have valence electrons on two different prin-

cipal energy levels. This sets them apart from the

representative or main-group elements, of which

the alkali metals and alkaline earth metals are a

part. The transition metals also follow irregular

patterns of orbital filling, which further distin-

guish them from the representative elements.

Other than that, there is not much that dif-

ferentiates the transition metals, a broad group-

ing that includes precious metals such as gold,

silver, and platinum, as well as highly functional

metals such as iron, manganese, and zinc. Many

of the transition metals, particularly those on

periods 4, 5, and 6, form useful alloys with one

another, and with other elements. Because of

their differences in valence electron configura-

tion, however, they do not always combine in the

same ways, even within an element. Iron, for

MOST “TIN” ROOFS NOW CONSIST OF AN IRON OR STEEL ROOF COVERED WITH ZINC, NOT TIN. (John Hulme; Eye Ubiqui-

tous/Corbis. Reproduced by permission.)


Metalsinstance, sometimes releases two electrons inchemical bonding, and at other times three.

GROUPING THE TRANSITION

METALS. There is no easy way to group thetransition metals, though some of these elementsare traditionally categorized together. These donot constitute “families” as such, but they do pro-vide useful ways to break down the otherwiserather daunting lineup of the transition metals.In two cases, there is at least a relationshipbetween group number on the periodic table andthe categories loosely assigned to a collection oftransition metals. Thus the “coinage metals”—copper, silver, and gold—all occupy Group 9 onthe IUPAC version of the periodic table. (Thesehave traditionally been associated with oneanother because their resistance to oxidation,combined with their malleability and beauty, hasmade them useful materials for fashioning coins.

Likewise the members of the “zinc group”—zinc, cadmium, and mercury—on Group 10 onthe IUPAC periodic table have often been associ-ated as a miniature unit due to common proper-ties. Members of the “platinum group”—plat-inum, iridium, osmium, palladium, rhodium,and ruthenium—occupy a rectangle on the table,corresponding to periods 5 and 6, and groups 6through 8. What actually makes them a “group,”however, is the fact that they tend to appeartogether in nature. Iron, nickel, and cobalt, foundalongside one another on Period 4, may begrouped together because they are all magneticto some degree.

The transition metals, listed below by atom-ic number and chemical symbol, are:

• 21. Scandium (Sc)

• 22. Titanium (Ti)

• 23. Vanadium (V)

• 24. Chromium (Cr)

• 25. Manganese (Mn)

• 26. Iron (Fe)

• 27. Cobalt (Co)

• 28. Nickel (Ni)

• 29. Copper (Cu)

• 30. Zinc (Zn)

• 39. Yttrium (Y)

• 40. Zirconium (Zr)

• 41. Niobium (Nb)

• 42. Molybdenum (Mo)

• 43. Technetium (Tc)

• 44. Ruthenium (Ru)

• 45. Rhodium (Rh)

• 46. Palladium (Pd)

• 47. Silver (Ag)

• 48. Cadmium (Cd)

• 57. Lanthanum (La)

• 72. Hafnium (Hf)

• 73. Tantalum (Ta)

• 74. Tungsten (W)

• 75. Rhenium (Re)

• 76. Osmium (Os)

• 77. Iridium (Ir)

• 78. Platinum (Pt)

• 79. Gold (Au)

• 80. Mercury (Hg)

• 89. Actinium (Ac)

• 104. Rutherfordium (Rf)

• 105. Dubnium (Db)

• 106. Seaborgium (Sg)

• 107. Bohrium (Bh)

• 108. Hassium (Hs)

• 109. Meitnerium (Mt)

• 110. Ununnilium (Uun)

• 111. Unununium (Uuu)

• 112. Ununbium (Uub)

The last nine, along with 11 of the actinides,are part of the list of metals known collectively asthe transuranium elements—that is, elementswith atomic numbers higher than 92. These haveall been produced artificially, in most cases with-in a laboratory. Note that the last three had notbeen named as of 2001: the “names” by whichthey are known simply mean, respectively, 110,111, and 112.

Lanthanides

The lanthanides are the first of two inner transi-tion metal groups. Note that in the list of transi-tion metals above, lanthanum (57) is followed byhafnium (72). The “missing” group of 14 ele-ments, normally shown at the bottom of theperiodic table, is known as the lanthanides, afamily which can be defined by the fact that allfill the 4f orbital. However, because lanthanum(which does not fill an f orbital) exhibits similarproperties, it is usually included with the lan-thanides.

Bright and silvery in appearance, many ofthe lanthanides—though they are metals—are sosoft they can be cut with a knife. They also tendto be highly reactive. Though they were onceknown as the “rare earth metals,” lanthanides



Metals only seemed rare because they were difficult toextract from compounds containing other sub-stances—including other lanthanides. Becausetheir properties are so similar, and because theytend to be found together in the same substances,the original isolation and identification of thelanthanides was an arduous task that took wellover a century.

The lanthanides (in addition to lanthanum,listed earlier with the transition metals), are:

• 58. Cerium (Ce)

• 59. Praseodymium (Pr)

• 60. Neodymium (Nd)

• 61. Promethium (Pm)

• 62. Samarium (Sm)

• 63. Europium (Eu)

• 64. Gadolinium (Gd)

• 65. Terbium (Tb)

• 66. Dysprosium (Dy)

• 67. Holmium (Ho)

• 68. Erbium (Er)

• 69. Thulium (Tm)

• 70. Ytterbium (Yb)

• 71. Lutetium (Lu)

Actinides

Just as lanthanum is followed on the periodictable by an element with an atomic number 15points higher, so actinium (89) is followed byrutherfordium (104). The elements in the “gap”are the other inner transition family, theactinides, which all fill the 5f orbital. These areplaced below the lanthanides at the bottom of thechart, and just as lanthanum is included with thelanthanides, even though it does not fill an forbital, so actinium is usually lumped in with theactinides due to similarities in properties.

Most of the actinides tend to be unstable orradioactive, the later ones highly so. The firstthree members of the group (not countingactinium) occur in nature, while the other 11 aretransuranium elements created artificially. Inmost cases, these were produced with a cyclotronor other machine for accelerating atoms, general-ly at the research center of the University of Cal-ifornia at Berkeley during a period from 1940 to1961. Einsteinium and fermium, however, wereby-products of nuclear testing at Bikini Atoll inthe south Pacific in 1952. The principal use of theactinides is in nuclear applications—weaponry

or power plants—though some of them also havespecialized uses as well.

The actinides (in addition to actinium, listedearlier with the transition metals), are:

• 90. Thorium (Th)

• 91. Protactinium (Pa)

• 92. Uranium (U)

• 93. Neptunium (Np)

• 94. Plutonium (Pu)

• 95. Americium (Am)

• 96. Curium (Cm)

• 97. Berkelium (Bk)

• 98. Californium (Cf)

• 99. Einsteinium (Es)

• 100. Fermium (Fm)

• 101. Mendelevium (Md)

• 102. Nobelium (No)

• 103. Lawrencium (Lr)

“Orphan” Metals

Seven other metals remain to be discussed. Theseare the “orphan” metals, which appear in groups3, 4, and 5 of the North American periodic table(13, 14, and 15 of the IUPAC table), and they arecalled “orphans” simply because none belongs toa clearly defined family. Sometimes aluminumand the three elements below it in Group 3—gal-lium, indium, and thallium—are lumped togeth-er as the “aluminum family.” This is not a widelyrecognized distinction, but will be used here onlyfor the purpose of giving some form of organiza-tion to the “orphan” elements.

Though they do not belong to families, the“orphan” metals are nonetheless highly impor-tant. Aluminum, after all, is the most abundantof all metals on Earth, and has wide applicationsin daily life. Tin, too, is widely known, as is lead.Somewhat less recognizable is bismuth, but itappears in a product with which most Americansare familiar: Pepto-Bismol.

The seven “orphans” are listed below, alongwith atomic number and chemical symbol:

• 13. Aluminum (Al)

• 31. Gallium (Ga)

• 49. Indium (In)

• 50. Tin (Sn)

• 81. Thallium (Tl)

• 82. Lead (Pb)

• 83. Bismuth (Bi)



MetalsAluminum

Named after alum, a salt known from ancienttimes and used by the Egyptians, Romans, andGreeks, aluminum is so reactive that it proveddifficult to isolate. Only in 1825 did Danishphysicist Hans Christian Ørsted (1777-1851) iso-late an impure version of aluminum metal, usinga complicated four-step process. Two years later,German chemist Friedrich Wöhler (1800-1882)obtained pure aluminum from a reaction ofmetallic potassium with aluminum chloride,or AlCl3.

For years thereafter, aluminum continued tobe so difficult to produce that it acquired the sta-tus of a precious metal. European kings displayedtreasures of aluminum as though they were goldor silver, and by 1855, it was selling for about$100,000 a pound. Then in 1886, French metal-lurgist Paul-Louis-Toussaint Héroult (1863-1914) and American chemist Charles MartinHall (1863-1914) independently developed aprocess that made aluminum relatively easy toextract by means of electrolysis. Thanks to whatbecame known as the Hall-Héroult process, theprice of aluminum dropped to around 30 cents apound.

America’s aluminum comes primarily frommines in Alabama, Arkansas, and Georgia, whereit often appears in a clay called kaolin, used inmaking porcelain. Aluminum can also be foundin deposits of feldspar, mica, granite, and otherclays. The principal sources of aluminum outsidethe United States include France, Surinam,Jamaica, and parts of Africa.

USES FOR ALUMINUM. Thoughaluminum is highly reactive, it does not corrodein moist air. Instead of forming rust the way irondoes, it forms a thin, hard, invisible coating ofaluminum oxide. This, along with its malleabili-ty, makes it highly useful for producing cans andother food containers. The World Trade CenterTowers in New York City gleam with a corrosion-resistant covering of aluminum. Likewise the ele-ment is used as a coating for mirrors: not only isit cheaper than silver, but unlike silver, it does nottarnish and turn black. Also, a thin coating ofmetallic aluminum gives Mylar balloons their sil-very sheen.

Aluminum conducts electricity about 60%as well as copper, meaning that it is an extraordi-narily good conductor, useful for transmittingelectrical power. Because it is much more light-

weight than copper and highly ductile, it is oftenused in high-voltage electric lines. Its conductiv-ity also makes aluminum a favorite material forkitchen pots and pans. Unlike copper, it is notknown to be toxic; however, because aluminumdissolves in strong acids, sometimes tomatoesand other acidic foods acquire a “metallic” tastewhen cooked in an aluminum pot.

Though it is soft in its pure form, whencombined with copper, manganese, silicon, mag-nesium, and/or zinc, aluminum can producestrong, highly useful alloys. These find applica-tion in automobiles, airplanes, bridges, highwaysigns, buildings, and storage tanks. Aluminum’shigh ductility also makes it the metal of choicefor foil wrappings—that is, aluminum foil. Com-pounds containing aluminum are used in a widerange of applications, including antacids;antiperspirants (potassium aluminum sulfatetends to close up sweat-gland ducts); and waterpurifiers.

Other Elements in the Alu-minum Group

GALLIUM. When Russian chemistDmitri Ivanovitch Mendeleev (1834-1907) creat-ed the periodic table in 1869, he predicted theexistence of several missing elements on the basisof “holes” between elements on the table. Amongthese was what he called “eka-aluminum,” dis-covered in the 1870s by French chemist PaulEmile Lecoq de Boisbaudran (1838-1912). Bois-baudran named the new element gallium, afterthe ancient name of his homeland, Gaul.

Gallium will melt if held in the hand, andremains liquid over a large range of tempera-tures—from 85.6°F (29.76°C) to 3,999.2°F(2,204°C). This makes it highly useful for ther-mometers with a large temperature range. Usedin a number of applications within the electron-ics industry, gallium is present in the compoundsgallium arsenide and gallium phosphide, appliedin lasers, solar cells, transistors, and other solid-state devices.

INDIUM AND THALLIUM. As Bois-baudran would later do, German mineralogistsFerdinand Reich (1799-1882) and HieronymusTheodor Richter (1824-1898) used a methodcalled spectroscopy to discover one of the alu-minum-group metals. Spectroscopy involvesanalyzing the frequencies of light (that is, the col-ors) emitted by an item of matter.



Metals However, Reich faced a problem whenstudying the zinc ores that he believed containeda new element: he was colorblind. Therefore hecalled on the help of Richter, his assistant. In1863, they discovered an element they calledindium because of the indigo blue color it emit-ted. Reich later tried to claim that he made thediscovery alone, but Richter’s role is indisputable.

Indium is used in making alloys, one ofwhich is a combination with gallium that is liq-uid at room temperature. Its alloys are used inbearings, dental amalgams, and nuclear controlroles. Various compounds are applied in solid-state electronics devices, and in electronicsresearch.

Two years before the discovery of indium,English physicist William Crookes (1832-1919)discovered another element using spectroscopicanalysis. This was thallium, named after theGreek word thallos (“green twig”), a reference toa lime-green spectral line that told Crookes hehad found a new element. Thallium is applied inelectronic equipment, and the compound thalli-um sulfate is used for killing ants and rodents.

Three Other “Orphans”

TIN. Known from ancient times, tin wascombined with copper to make bronze—an alloyso significant it defines an entire stage of humantechnological development. Tin’s chemical sym-bol, Sn, refers to the name the Romans gave it:stannum. Just as tin added to copper gave bronzeits strength, alloys of tin with copper and anti-mony created pewter, a highly useful, malleablematerial for pots and dinnerware popular fromthe thirteenth century through the early nine-teenth century.

Tin is found primarily in Bolivia, Malaysia,Indonesia, Thailand, Zaire, and Nigeria—a widegeographical range. Like aluminum today, it hasbeen widely used as a coating for other metals toretard corrosion. Hence the term “tin can,” refer-ring to a steel can coated with tin. Also, tin hasbeen used to coat iron or steel roofs, but becausezinc is cheaper, it is now most often the coatingof choice; nonetheless, these are still typicallycalled “tin roofs.”

LEAD. Due to knowledge of its toxicproperties, lead is not frequently used today, butthis was not always the case. The Romans, forinstance, used plumbum as they called it (hencethe chemical symbol Pb) as a material for making

water pipes. This is the root of our own wordplumber. (The same root explains the verbplumb or the noun plumb bob. A plumb bob is alead weight, hung from a string, used to ensurethat the walls of a structure are “plumb”—that is,perpendicular to the foundation.) Many histori-ans believe that plumbum in the Romans’ watersupply was one of the reasons behind the declineand fall of the Roman Empire.

The human body can only excrete very smallquantities of lead a day, and this is particularlytrue of children. Even in small concentrations,lead can cause elevation of blood pressure;reduction in the synthesis of hemoglobin, whichcarries oxygen from the lungs to the blood andorgans; and decreased ability to utilize vitamin Dand calcium for strengthening bones. Higherconcentrations of lead can effect the centralnervous system, resulting in decreased mentalfunctioning and hearing damage. Prolongedexposure can result in a coma or even death.

Before these facts became widely known inthe late twentieth century, lead was applied as aningredient in paint. In addition, it was used inwater pipes, and as an anti-knock agent in gaso-lines. Increased awareness of the health hazardsinvolved have led to a discontinuation of thesepractices. (Note that pencil “lead” is actuallygraphite, a form of carbon.)

Lead, however, is used for absorbing radia-tion—for instance, as a shield against X rays.Another important use for lead today is in stor-age batteries. Yet another “application” of leadrelates to its three stable isotopes (lead-206, -207,and -208), the end result of radioactive decay onthe part of various isotopes of uranium andother radioactive elements.

BISMUTH. During the Middle Ages,when a number of unscientific ideas prevailed,scholars believed that lead eventually became tin,tin eventually turned into bismuth, and bismuthdeveloped into silver. Hence they referred to bis-muth as “roof of silver.” Only in the sixteenthcentury did German metallurgist Georgius Agri-cola (George Bauer; 1494-1555) put forward theidea that bismuth was a separate element.

Bismuth appears in nature not only in itselemental form, but also in a number of com-pounds and ores. It is also obtained as a by-prod-uct from the smelting of gold, silver, copper, andlead. Because bismuth expands dramaticallywhen it cools, early printers used a bismuth alloy



Metals




which does not fill the f orbital—exhibits




ALKALI METALS: All members, ex-cept

hydrogen, of Group 1 on the periodic table

of elements, with valence electron configu-

rations of ns1.



valence electron configurations of ns2.

ALLOY: A mixture containing more than

one metal.

CONDUCTIVITY: The ability to conduct

heat or electricity. Metals are noted for

their high levels of conductivity.

CRYSTALLINE: A term describing an

internal structure in which the constituent

parts have a simple and definite geometric

arrangement, repeated in all directions.

Metals have a crystalline structure.

DUCTILE: Capable of being bent or

molded into various shapes without

breaking.

ELECTRON SEA MODEL: A widely

accepted model of internal bonding within

metals, which depicts metal atoms as

floating in a “sea” of valence electrons.






INNER TRANSITION METALS: The

lanthanides and actinides, both of which

fill the f orbitals. For this reason, they are

usually treated as distinct from the transi-

tion metals.











IUPAC SYSTEM: A version of the peri-

odic table of elements, authorized by the

International Union of Pure and Applied

Chemistry (IUPAC), which numbers all

groups on the table from 1 to 18. Thus in

the IUPAC system, in use primarily outside

of North America, both the representative

or main-group elements and the transition

metals are numbered.


that fill the 4f orbital. Because lan-

thanum—which does not fill the f




METALS: A collection of 87 elements

that includes numerous families—the

alkali metals, alkaline earth metals, transi-

tion metals, lanthanides, and actinides, as

well as seven elements in groups 3 through

5 of the North American system periodic

table. Metals are lustrous or shiny in

appearance, and ductile. They are excellent

conductors of heat and electricity, and tend

to form positive ions by losing electrons.

KEY TERMS


Metals


NORTH AMERICAN SYSTEM: A ver-

sion of the periodic table of elements that

only numbers groups of elements in which

the number of valence electrons equals the

group number. Hence the transition metals

are usually not numbered in this system.





complex the pattern of orbitals.





for that particular element. Elements are

arranged in groups or columns, as well as

in rows or periods, according to specific

aspects of their valence electron config-

urations.

PERIODS: Rows on the periodic table of



in the atoms of the elements involved.









orbitals.







decays to form another isotope until even-

tually (though this may take a long time) it

becomes stable.

REACTIVITY: The tendency for bonds

between atoms or molecules to be made or

broken in such a way that materials are

transformed.

REPRESENTATIVE OR MAIN-GROUP

ELEMENTS: The 44 elements in Groups

1 through 8 on the North American ver-

sion of periodic table of elements, for

which the number of valence electrons

equals the group number. (The only excep-

tion is helium.) The alkali metals and alka-

line earth metals are representative ele-

ments; all other metals are not.

SALT: Generally, a salt is any compound

that brings together a metal and a non-

metal. Salts are produced, along with water,

in the reaction of an acid and a base.


elements (counting lanthanum and actini-

um), which are not assigned a group num-

ber in the version of the periodic table of

elements known as the North American

system. These are the only elements that fill

the d orbital. In addition, the transition

metals—unlike the representative or main-

group elements—have their valence elec-

trons on two different principal energy lev-

els. Though the lanthanides and actinides

are known as inner transition metals, they

are usually treated separately.


that occupy the highest principal energy

level in an atom. These are the electrons

involved in chemical bonding.

KEY TERMS CONTINUED


Metalsto make metal type. When poured into forms cutwith the shapes of letters and other characters,the alloy produced symbols with clear, sharpedges.

Wood’s metal is a bismuth alloy involvingcadmium, tin, and lead, which melts at a temper-ature of 158°F (70°C). This makes it useful forautomatic sprinkler systems inside a building:when a fire breaks out, it melts the Wood’s metalseal, and turns on the sprinklers. Bismuth com-pounds are applied in cosmetics, paints, anddyes, and prior to the twentieth century, com-pounds of bismuth were widely used for treatingsexually transmitted diseases. Today, antibioticshave largely taken their place, but bismuth stillhas one well-known medicinal application: inbismuth subsalicylate, better known by the brandname Pepto-Bismol.

WHERE TO LEARN MORE

Aluminum Association (Web site). <http://www.aluminum.org> (May 28, 2001).

Kerrod, Robin. Matter and Materials. Illustrated by TerryHadler. Tarrytown, NY: Benchmark Books, 1996.

“Lead Programs.” Environmental Protection Agency (Website). <http://www.epa.gov/opptintr/lead/> (May 28,2001).

Mebane, Robert C. and Thomas R. Rybolt. Metals. Illus-trated by Anni Matsick. New York: Twenty-First Cen-tury Books, 1995.

“Metals Industry Guide.” About.com (Web site).<http://metals.about.com/industry/metals/mbody.htm> (May 28, 2001).

“Minerals and Metals.” Natural Resources of Canada(Web site). <http://www.nrcan.gc.ca/mms/school/e_mine.htm> (May 28, 2001).

Oxlade, Chris. Metal. Chicago, IL: Heinemann Library,2001.

Snedden, Robert. Materials. Des Plaines, IL: HeinemannLibrary, 1999.

WebElements (Web site). <http://www.webelements.com> (May 22, 2001).

Whyman, Kathryn. Metals and Alloys. Illustrated byLouise Nevett and Simon Bishop. New York: Glouces-ter Press, 1988.





ALKALI METALSAlkali Metals

CONCEPTGroup 1 of the periodic table of elements con-sists of hydrogen, and below it the six alkali met-als: lithium, sodium, potassium, rubidium,cesium, and francium. The last three areextremely rare, and have little to do with every-day life; on the other hand, it is hard to spend aday without encountering at least one of the firstthree—particularly sodium, found in table salt.Along with potassium, sodium is an importantcomponent of the human diet, and in com-pounds with other substances, it has an almostendless array of uses. Lithium does not have asmany applications, but to many people who havereceived it as a medication for bipolar disorder, itis quite literally a life-saver.

HOW IT WORKS

Electron Configuration of the Alkali Metals

In the essay on Families of Elements, there is alengthy discussion concerning the relationshipbetween electron configuration and the defini-tion of a particular collection of elements as a“family.” Here that subject will only be touchedupon lightly, inasmuch as it relates to the alkalimetals.

All members of Group 1 on the periodictable of elements have a valence electron config-uration of s1. This means that a single electron isinvolved in chemical bonding, and that this sin-gle electron moves through an orbital, or rangeof probabilities, roughly corresponding to asphere.

Most elements bond according to what isknown as the octet rule, meaning that when twoor more atoms are bonded, each has (or shares)eight valence electrons. It is for this reason thatthe noble gases, at the opposite side of the peri-odic table from the alkali metals, almost neverbond with other elements: they already haveeight valence electrons.

The alkali metals, on the other hand, arequite likely to find “willing partners,” since theyeach have just one valence electron. This bringsup one of the reasons why hydrogen, though it isalso part of Group 1, is not included as an alkalimetal. First and most obviously, it is not a metal;additionally, it bonds according to what is calledthe duet rule, such that it shares two electronswith another element.

Chemical and Physical Characteristics of the Alkali Metals

The term “alkali” (essentially the opposite ofan acid) refers to a substance that forms the neg-atively charged hydroxide ion (OH-) in contactwith water. On their own, however, alkali metalsalmost always form positive ions, or cations, witha charge of +1.

When alkali metals react with water, onehydrogen atom splits off from the water moleculeto form hydrogen gas, while the other hydrogenatom joins the oxygen to form hydroxide. Wherethe heavier members of the alkali metal familyare concerned, reactions can often be so vigorousthat the result is combustion or even explosion.Alkali metals also react with oxygen to produce


Alkali Metals


either an oxide, peroxide, or superoxide, depend-ing on the particular member of the alkali metalfamily involved.

Shiny and soft enough to be cut with a knife,the alkali metals are usually white (thoughcesium is more of a yellowish white). Whenplaced in a flame, most of these substances pro-duce characteristic colors: lithium, for instance,glows bright red, and sodium an intense yellow.Heated potassium produces a violet color, rubid-ium a dark red, and cesium a light blue. Thismakes it possible to identify the metals, whenheated, by color—a useful trait, since they are sooften inclined to be bonded with other elements.

MELTING AND BOILING

POINTS. As one moves down the rows or

periods of the periodic table, the mass of atoms

increases, as does the energy each atom possess-

es. Yet the amount of energy required to turn a

solid alkali metal into a liquid, or to vaporize a

liquid alkali metal, actually decreases with higher

atomic number. In other words, the higher the

atomic number, the lower the boiling and melt-

ing points.

The list below lists the six alkali metals in

order of atomic number, along with chemical

symbol, atomic mass, melting point, and boiling

point. Note that for francium, which is radioac-

tive, the figures given are for its most stable iso-

tope, francium-223 (223Fr)—which has a half-life

of only about 21 minutes.

ONE OF THE USES OF LITHIUM IS IN BATTERIES. THE NISSAN HYPERMINI IS AN ELECTRIC VEHICLE POWERED BY A

LITHIUM BATTERY. (Reuters NewMedia Inc./Corbis. Reproduced by permission.)


Alkali Metals


Atomic Number, Mass, and Melting and BoilingPoints of the Alkali Metals:

• 3. Lithium (Li), 6.941; Melting Point:356.9°F (180.5°C); Boiling Point: 2,457°F(1,347°C)

• 11. Sodium (Na), 22.99; Melting Point:208°F (97.8°C); Boiling Point: 1,621.4°F(883°C)

• 19. Potassium (K), 39.10; Melting Point:145.9°F (63.28°C); Boiling Point: 1,398.2°F(759°C)

• 37. Rubidium (Rb), 85.47; Melting Point:102.8°F (39.31°C); Boiling Point: 1,270.4°F(688°C)

• 55. Cesium (Cs), 132.9; Melting Point:83.12°F (28.4°C); Boiling Point: 1,239.8°F(671°C)

• 87. Francium (Fr), (223); Melting Point:80.6°F (27°C); Boiling Point: 1,250.6°F(677°C)

Abundance of Alkali Metals

Sodium and potassium are, respectively, the sixthand seventh most abundant elements on Earth,comprising 2.6% and 2.4% of the planet’s knownelemental mass. This may not seem like much,but considering the fact that just two elements—

oxygen and silicon—make up about 75%, andthat just 16 elements make up most of theremainder, it is an impressive share.

Lithium, on the other hand, is much lessabundant, and therefore, figures for its part ofEarth’s known elemental mass are measured inparts per million (ppm). The total lithium inEarth’s crust is about 17 ppm. Surprisingly,rubidium is more abundant, at 60 ppm; less sur-prisingly, cesium, with just 3 ppm, is very rare.Almost no francium is found naturally, except invery small quantities within uranium ores.


Lithium

Swedish chemist Johan August Arfvedson (1792-1841) discovered lithium in 1817, and named itafter the Greek word for “stone.” Four years later,another scientist named W. T. Brande succeededin isolating the highly reactive metal. Most of thelithium available on Earth’s crust is bound upwith aluminum and silica in minerals.

Since the time of its discovery, lithium hasbeen used in lubricants, glass, and in alloys of

ONE OF THE MOST STRIKING USES OF LITHIUM OCCURRED IN 1932, WITH THE DEVELOPMENT OF THE FIRST PARTI-CLE ACCELERATOR. SHOWN HERE ARE ITS INVENTORS, ERNEST WALTON (LEFT), ERNEST RUTHERFORD (CENTER),AND JOHN COCKCROFT. (Hulton-Deutsch Collection/Corbis. Reproduced by permission.)


Alkali Metals


lead, aluminum, and magnesium. In glass, it actsas a strengthening agent; likewise, metal alloysthat contain lithium tend to be stronger, yet lessdense. In 1994, physicist Jeff Dahn of SimonFraser University in British Columbia, Canada,developed a lithium battery. Not only was thebattery cheaper to produce than the traditionalvariety, Dahn and his colleagues announced, butthe disposal of used lithium batteries presentedless danger to the environment.

One of the most striking uses of lithiumoccurred in 1932, when English physicist John D.Cockcroft (1897-1967) and Irish physicist ErnestWalton (1903-1995) built the first particle accel-erator. By bombarding lithium atoms, they pro-duced highly energized alpha particles. This wasthe first nuclear reaction brought about by theuse of artificially accelerated particles—in otherwords, without the need for radioactive materialssuch as uranium-235. Cockcroft’s and Walton’sexperiment with lithium thus proved pivotal tothe later creation of the atomic bomb.

LITHIUM IN PSYCHIATRIC

TREATMENT. The most important applica-tion of lithium, however, is in treatment for thepsychiatric condition once known as manicdepression, today identified as bipolar disorder.Persons suffering from bipolar disorder tendtoward mood swings: during some periods thepatient is giddy (“manic,” or in a condition of“mania”), and during others the person is suici-dal. Indeed, prior to the development of lithiumas a treatment for bipolar disorder, as many asone in five patients with this condition commit-ted suicide.

Doctors do not know exactly how lithiumdoes what it does, but it obviously works:between 70% and 80% of patients with the bipo-lar condition respond well to treatment, and areable to go on with their lives in such a way thattheir condition is no longer outwardly evident.Lithium is also administered to patients who suf-fer unipolar depression and some forms of schiz-ophrenia.

EARLY MEDICINAL USES OF

LITHIUM. It is said that the great Greco-Roman physician Galen ( 129-c. 199) counseledpatients suffering from “mania” to bathe in, andeven drink the water from, alkaline springs. If so,he was nearly 2,000 years ahead of his time. Evenin the 1840s, not long after lithium was discov-ered, the mineral—mixed with carbonate or cit-

rate—was touted as a cure for insomnia, gout,epilepsy, diabetes, and even cancer.

None of these alleged cures proved a success;nor did a lithium chloride treatment adminis-tered in the 1940s as a salt substitute for patientson low-sodium diets. As it turned out, when notenough sodium is present, the body experiencesa buildup of sodium’s sister element, lithium.The result was poisoning, which in some casesproved fatal.

CADE’S BREAKTHROUGH. Thenin 1949, Australian psychiatrist John Cade dis-covered the value of lithium for psychiatric treat-ment. He approached the problem from anentirely different angle, experimenting with uricacid, which he believed to be a cause of manicbehavior. In administering the acid to guineapigs, he added lithium salts merely to keep theuric acid soluble—and was very surprised bywhat he discovered. The uric acid did not makethe guinea pigs manic, as he had expected;instead, they became exceedingly calm.

Cade changed the focus of his research, andtested lithium treatment on ten manic patients.Again, the results were astounding: one patientwho had suffered from an acute bipolar disorder(as it is now known) for five years was releasedfrom the hospital after three months of lithi-um treatment, and went on to lead a healthy,normal life.

Encouraged by the changes he had seen inpatients who received lithium, Cade published areport on his findings in the Medical Journal ofAustralia, but his work had little impact at thetime. Nor did the idea of lithium treatment meetwith an enthusiastic reception on the other sideof the Pacific: in the aftermath of the failedexperiments with lithium as a sodium substitutein the 1940s, stories of lithium poisoning werewidespread in the United States.

LITHIUM TODAY. Were it not for theefforts of Danish physician Mogens Schou, lithi-um might never have taken hold in the medicalcommunity. During the 1950s and 1960s, Schoucampaigned tirelessly for recognition of lithiumas a treatment for manic-depressive illness. Final-ly during the 1960s, the U.S. Food and DrugAdministration began conducting trials of lithi-um, and approved its use in 1974. Today some200,000 Americans receive lithium treatments.


Alkali Metals

A non-addictive and non-sedating medica-tion, lithium—as evidenced by the failed experi-ment in the 1940s—may still be dangerous inlarge quantities. It is absorbed quickly into thebloodstream and carried to all tissues in thebrain and body before passing through the kid-neys. Both lithium and sodium are excretedthrough the kidneys, and since sodium affectslithium excretion, it is necessary to maintain aproper quantity of sodium in the body. For thisreason, patients on lithium are cautioned toavoid a low-salt diet.

Sodium

Sodium compounds had been known for sometime prior to 1807, when English chemist SirHumphry Davy (1778-1829) succeeded in isolat-ing sodium itself. The element is represented by achemical symbol (Na), reflecting its Latin name,natrium. In its pure form, sodium has a bright,shiny surface, but in order to preserve thisappearance, it must be stored in oil: sodiumreacts quickly with oxygen, forming a white crustof sodium oxide.

Pure sodium never occurs in nature; instead,it combines readily with other substances to formcompounds, many of which are among the mostwidely used chemicals in industry. It is also high-ly soluble: thus whereas sodium and potassiumoccur in crystal rocks at about the same ratio,sodium is about 30 times more abundant in sea-water than its sister element.

OBTAINING SODIUM CHLO-

RIDE. Though the extraction of sodiuminvolves the use of a special process, the metal isplentiful in the form of sodium chloride—betterknown as table salt. In fact, the term salt in chem-istry refers generally to any combination of ametal with a nonmetal. More specifically, salts are(along with water) the product of reactionsbetween acids and bases.

Sodium chloride is so easy to obtain, andtherefore so cheap, that most industries makingother sodium compounds use it, simply separat-ing out the chloride (as described below) beforeadding other elements. The United States is theworld’s largest producer of sodium chloride,obtained primarily from brine, a term used todescribe any solution of sodium chloride inwater. Brine comes from seawater, subterraneanwells, and desert lakes, such as the Great Salt Lakein Utah. Another source of sodium chloride is

rock salt, created underground by the evapora-tion of long-buried saltwater seas.

Other top sodium-chloride-producingnations include China, Germany, Great Britain,France, India, and various countries in the formerSoviet Union. Salt may be cheap and plentiful forthe world in general, but there are places where itis a precious commodity. One such place is theSahara Desert, where salt caravans ply a brisktrade today, much as they have since ancient times.

ISOLATING SODIUM. Modernmethods for the production of sodium representan improvement in the technique Davy used in1807, although the basic principle is the same.Though several decades passed before electricitycame into widespread public use, scientists hadbeen studying its properties for years, and Davyapplied it in a process called electrolysis.

Electrolysis is the use of an electric currentto produce a chemical reaction—in this case, toseparate sodium from the other element or ele-ments with which it is combined. Davy first fusedor melted a sample of sodium chloride, thenelectrolyzed it. Using an electrode, a device thatconducts electricity and is used to emit or collectelectric charge, he separated the sodium chloridein such a way that liquid sodium metal collectedon the cathode, or negatively charged end. Mean-while, the gaseous chlorine was released throughthe anode, or the positively charged end.

The apparatus used for sodium separationtoday is known as the Downs cell, after its inven-tor, J. C. Downs. In a Downs cell, sodium chlo-ride and calcium chloride are combined in amolten mixture in which the presence of calciumchloride lowers the melting point of the sodiumchloride by more than 30%. When an electriccurrent is passed through the mixture, sodiumions move to the cathode, where they pick upelectrons to become sodium atoms. At the sametime, ions of chlorine migrate to the anode, los-ing electrons to become chlorine atoms.

Sodium is a low-density material that floatson water, and in the Downs cell, the molten sodi-um rises to the top, where it is drawn off. Thechlorine gas is allowed to escape through a ventat the top of the anode end of the cell, and theresulting sodium metal—that is, the elementalform of sodium—is about 99.8% pure.

USES FOR SODIUM CHLORIDE.

As indicated earlier, sodium chloride is by far themost widely known and commonly used sodium



Alkali Metals

compound—and this in itself is a distinction,given the fact that so many sodium compoundsare a part of daily life. Today people think of saltprimarily as a seasoning to enhance the taste offood, but prior to the development of refrigera-tion, it was vital as a preservative because it keptmicrobes away from otherwise perishable fooditems.

Salt does not merely improve the taste offood; it is an essential nutrient. Sodium com-pounds regulate transmission of signals throughthe nervous system, alter the permeability ofmembranes, and perform a number of other life-preserving functions. On the other hand, toomuch salt can aggravate high blood pressure.Thus, since the 1970s and 1980s, food manufac-turers have increasingly offered products low insodium, a major selling point for health-con-scious consumers.

OTHER SODIUM COMPOUNDS.

In addition to its widespread use in consumergoods, sodium chloride is the principal source ofsodium used in making other sodium com-pounds. These include sodium hydroxide, formanufacturing cellulose products such as film,rayon, soaps, and paper, and for refining pe-troleum. In its application as a cleaning solu-tion, sodium hydroxide is known as caustic sodaor lye.

Another widely used sodium compound issodium carbonate or, soda ash, applied in glass-making, paper production, textile manufactur-ing, and other areas, such as the production ofsoaps and detergents. Sodium also can be com-bined with carbon to produce sodium bicarbon-ate, or baking soda. Sodium sulfate, sometimesknown as salt cake, is used for making cardboardand kraft paper. Yet another widely used sodiumcompound is sodium silicate, or “water glass,”used in the production of soaps, detergents, andadhesives; in water treatment; and in bleachingand sizing of textiles.

Still other sodium compounds used byindustry and/or consumers include sodiumborate, or borax; sodium tartrate, or sal tartar;the explosive sodium nitrate, or Chilean salt-peter; and the food additive monosodium gluta-mate (MSG). Perhaps ironically, there are fewuses for pure metallic sodium. Once applied asan “anti-knock” additive in leaded gasoline,before those products were phased out for envi-ronmental reasons, metallic sodium is now used

as a heat-exchange medium in nuclear reactors.But its widest application is in the production ofthe many other sodium compounds used aroundthe world.

Potassium

In some ways, potassium is a strange substance,as evidenced by its behavior in response to water.As everyone knows, water tends to put out a fire,and most explosives, when exposed to sufficientquantities of water, become ineffective. Potassi-um, on the other hand, explodes in contact withwater and reacts violently with ice at tempera-tures as low as -148°F (-100°C). In a completereversal of the procedures normally followed formost substances, potassium is stored in kerosene,because it might burst into flames if exposed tomoist air!

Many aspects of potassium mirror thosealready covered with regard to sodium. The twohave a number of the same applications, and incertain situations, potassium is used as a sodiumsubstitute. Like sodium, potassium is neverfound alone in nature; instead, it comes primari-ly from sylvinite and carnalite, two ores contain-ing potassium chloride. Also, like sodium, potas-sium was first isolated in 1807 by Davy, using theprocess of electrolysis described above. A fewyears later, a German chemist dubbed the newlyisolated element “kalium,” apparently a deriva-tion of the Arabic qali, for “alkali”; hence the useof K as the chemical symbol for potassium.

USES FOR POTASSIUM. Potassiumhas another similarity with sodium; although itwas not isolated until the early nineteenth centu-ry, its compounds have been in use for many cen-turies. The Romans, for instance, used potassiumcarbonate, or potash, obtained from the ashes ofburned wood, to make soap. During the MiddleAges, the Chinese applied a form of saltpeter,potassium nitrate, in making gunpowder. And incolonial America, potash went into the produc-tion of soap, glass, and other products.

The production of just one ton of potashrequired the burning of several acres’ worth oftrees—a wasteful practice in more ways than one.Though there was no environmentalist move-ment in those days, financial concerns never goout of style. In order to save the money lost byusing up vast acres of timber, American industryin the nineteenth century sought another meansof making potash. The many similarities between



Alkali Metals

sodium and potassium provided a key, and thesubstitution of sodium carbonate for potassiumcarbonate saved millions of trees.

In 1847, German chemist Justus von Liebig(1803-1873) discovered potassium in living tis-

sues. As a result, scientists became aware of the

role this alkali metal plays in sustaining life:

indeed, potassium is present in virtually all living

cells. In the human body, potassium—which

accounts for only 0.4% of the body’s mass—is


IN THE SAHARA DESERT, SALT CARAVANS PLY A BRISK TRADE TODAY, MUCH AS THEY HAVE SINCE ANCIENT TIMES.HERE, TRADERS IN MALI LOAD A CAMEL FOR A SALT CARAVAN. (Nik Wheeler/Corbis. Reproduced by permission.)


Alkali Metals


ALKALI METALS: The elements in

Group 1 of the periodic table of elements,

with the exception of hydrogen. The alkali

metals, which include lithium, sodium,

potassium, rubidium, cesium, and franci-

um, all have one valence electron in the s1

orbital, and are highly reactive.

ANODE: An electrode at the positively

charged end of a supply of electric current.

BRINE: A term used to describe any

solution of sodium chloride in water.

CATHODE: An electrode at the negative-

ly charged end of a supply of electric cur-

rent.



All of the alkali metals tend to form +1

cations (pronounced KAT-ie-un).

ELECTRODE: A structure, often a metal

plate or grid, that conducts electricity, and

which is used to emit or collect electric

charge.

ELECTROLYSIS: The use of an electric

current to cause a chemical reaction.



amount.









unstable—that is, radioactive. Such is the

case with francium, a radioactive member

of the alkali metals family.



atom in a particular energy state. The six

alkali metals all have valence electrons in

an s1 orbital, which describes a more or less

spherical shape.






chemical characteristics; the alkali metals

occupy Group 1.





ticles. “Decay” in this sense does not mean

“rot”; instead, radioactive isotopes contin-

ue changing into other isotopes until they

become stable.

SALT: A compound formed, along with

water, by the reaction of an acid and a base.

Generally, salts are any compounds that

bring together a metal and a nonmetal.



atom, and which are involved in chemical

bonding. The alkali metals all have one

valence electron, in the s1 orbital.

KEY TERMS


Alkali Metals


essential to the functioning of muscles. In largerquantities, however, it can be dangerous, causinga state of permanent relaxation known as potas-sium inhibition.

Since plants depend on potassium forgrowth, it was only logical that potassium, in theform of potassium chloride, was eventuallyapplied as a fertilizer. This, at least, distinguishesit from its sister element: sodium, or sodiumchloride, which can kill plants if administered tothe soil in large enough quantities.

Another application of potassium is in thearea pioneered by the Chinese about 800 yearsago: the manufacture of fireworks and gunpow-der from potassium nitrate. Like ammoniumnitrate, made infamous by its use in the 1993World Trade Center bombing and the OklahomaCity bombing in 1995, potassium nitrate doublesas a fertilizer.

Rubidium, Cesium, and Francium

The three heaviest alkali metals are hardly house-hold names, though one of them, cesium, doeshave several applications in industry. Rubidiumand cesium, discovered in 1860 by Germanchemist R. W. Bunsen (1811-1899) and Germanphysicist Gustav Robert Kirchhoff (1824-1887),were the first elements ever found using a spec-troscope. Matter emits electromagnetic radiationalong various spectral lines, which can be record-ed using a spectroscope and then analyzed to dis-cern the particular “fingerprint” of the substancein question.

When Bunsen and Kirchhoff saw the bluishspectral lines emitted by one of the two elements,they named it cesium, after a Latin word mean-ing “sky blue.” Cesium, which is very rare,appears primarily in compounds such as pollu-cite. It is used today in photoelectric cells, mili-tary infrared lamps, radio tubes, and videoequipment. During the 1940s, American physi-cist Norman F. Ramsey, Jr. (1915-) built a highlyaccurate atomic clock based on the natural fre-quencies of cesium atoms.

Rubidium, by contrast, has far fewer applica-tions, and those are primarily in areas of scientif-ic research. On Earth it is found in pollucite, lep-idolite, and carnallite. It is considerably moreabundant than cesium, and vastly more so thanfrancium. Indeed, it is estimated that if all thefrancium in Earth’s crust were combined, itwould have a mass of about 25 grams.

Francium was discovered in 1939 by Frenchphysicist Marguerite Perey (1909-1975), studentof the famous French-Polish physicist andchemist Marie Curie (1867-1934). For about fourdecades, scientists had been searching for themysterious Element 87, and while studying thedecay products of an actinium isotope, actinium-227, Perey discovered that one out of 100 suchatoms decayed to form the undiscovered ele-ment. She named it francium, after her home-land. Though the discovery of francium solved amystery, the element has no known uses outsideof its applications in research.

WHERE TO LEARN MORE

“Alkali Metals” (Web site). <http://www.midlink.com/~jfromm/elements/alkali.htm> (May 24, 2001).

“Alkali Metals” ChemicalElements.com (Web site).<http://www.chemicalelements.com/groups/alkali.html> (May 24, 2001).

“Hydrogen and the Alkali Metals.” University of ColoradoDepartment of Physics (Web site). <http://www.colorado.edu/physics/2000/periodic_table/alkalimetals.html> (May 24, 2001).

Kerrod, Robin. Matter and Materials. Illustrated by TerryHadler. Tarrytown, N.Y.: Benchmark Books, 1996.




“Visual Elements: Group I—The Alkali Metals” (Website). <http://www.chemsoc.org/viselements/pages/data/intro_groupi_data.html> (May 24, 2001).



ALKALINE EARTH METALSAlkaline Earth Metals

CONCEPTThe six alkaline earth metals—beryllium, mag-nesium, calcium, strontium, barium, and radi-um—comprise Group 2 on the periodic table ofelements. This puts them beside the alkali metalsin Group 1, and as their names suggest, the twofamilies share a number of characteristics, mostnotably their high reactivity. Also, like the alkalimetals, or indeed any other family on the period-ic table, not all members of the alkali metal fam-ily are created equally in terms of their abun-dance on Earth or their usefulness to human life.Magnesium and calcium have a number of uses,ranging from building and other structuralapplications to dietary supplements. In fact, bothare significant components in the metabolism ofliving things—including the human body. Bari-um and beryllium have numerous specializedapplications in areas from jewelry to medicine,while strontium is primarily used in fireworks.Radium, on the other hand, is rarely used outsideof laboratories, in large part because its radioac-tive qualities pose a hazard to human life.

HOW IT WORKS

Defining a Family

The expression “families of elements” refers togroups of elements on the periodic table thatshare certain characteristics. These include (inaddition to the alkaline earth metals and thealkali metals) the transition metals, halogens,noble gases, lanthanides, and actinides. (All ofthese are covered in separate essays within thisbook.) In addition, there are several larger cate-gories with regard to shared traits that often cross

family lines; thus all elements are classified eitheras metals, metalloids, and nonmetals. (These arealso discussed in separate essays, which includereference to “orphans,” or elements that do notbelong to one of the families mentioned above.)

These groupings, both in terms of familyand the broader divisions, relate both to external,observable characteristics, as well as to behaviorson the part of electrons in the elements’ atomicstructures. For instance, metals, which comprisethe vast majority of elements on the periodictable, tend to be shiny, hard, and malleable (thatis, they can bend without breaking.) Many ofthem melt at fairly high temperatures, and virtu-ally all of them vaporize (become gases) at hightemperatures. Metals also form ionic bonds, thetightest form of chemical bonding.

ELECTRON CONFIGURATIONS

OF THE ALKALINE EARTH METALS.

Where families are concerned, there are certainobservable properties that led chemists in thepast to group the alkaline earth metals together.These properties will be discussed with regard tothe alkaline earth metals, but another pointshould be stressed in relation to the division ofelements into families. With the advances inunderstanding that followed the discovery of theelectron in 1897, along with the development ofquantum theory in the early twentieth century,chemists developed a more fundamental definition of family in terms of electron con-figuration.

As noted, the alkaline earth metal familyoccupies the second group, or column, in theperiodic table. All elements in a particular group,regardless of their apparent differences, have acommon pattern in the configuration of their


AlkalineEarth Metals


valence electrons—the electrons at the “outside”of the atom, involved in chemical bonding. (Bycontrast, the core electrons, which occupy lowerregions of energy within the atom, play no role inthe bonding of elements.)

All members of the alkaline earth metal fam-ily have a valence electron configuration of s2.This means that two electrons are involved inchemical bonding, and that these electrons movethrough an orbital, or range of probabilities,roughly corresponding to a sphere. The s orbitalpattern corresponds to the first of several sub-levels within a principal energy level.

Whatever the number of the principal ener-gy level which corresponds to the period, or row,

on the periodic table the atom has the samenumber of sublevels. Thus beryllium, on Period2, has two principal energy levels, and its valenceelectrons are in sublevel 2s2. At the other end ofthe group is radium, on Period 7. Though radi-um is far more complex than beryllium, withseven energy levels instead of two, nonetheless ithas the same valence electron configuration, onlyon a higher energy level: 7s2.

HELIUM AND THE ALKALINE

EARTH METALS. If one studies the valenceelectron configurations of elements on the peri-odic table, one notices an amazing symmetry andorder. All members of a group, though their prin-cipal energy levels differ, share characteristics in

DURING WORLD WAR II, MAGNESIUM WAS HEAVILY USED IN AIRCRAFT COMPONENTS. IN THIS 1941 PHOTO, WORK-ERS POUR MOLTEN MAGNESIUM INTO A CAST AT THE WRIGHT AERONAUTICAL CORPORATION. (Bettmann/Corbis. Reproduced

by permission.)


AlkalineEarth

Metals


their valence shell patterns. Furthermore, for theeight groups numbered in the North Americanversion of the periodic table, the group numbercorresponds to the number of valence electrons.

There is only one exception: helium, with avalence electron configuration of 1s2, is normallyplaced in Group 8 with the noble gases. Based onthat s2 configuration, it might seem logical toplace helium atop beryllium in the alkaline earthmetals family; but there are several reasons whythis is not done. First of all, helium is obviouslynot a metal. More importantly, helium behaves ina manner quite different from that of the alkalineearth metals.

Whereas helium, like the rest of the noblegases, is highly resistant to chemical reactionsand bonding, alkaline earth metals are known fortheir high reactivity—that is, a tendency forbonds between atoms or molecules to be made orbroken so that materials are transformed. (Asimilar relationship exists in Group 1, whichincludes hydrogen and the alkali metals. All havethe same valence configuration, but hydrogen isnever included as a member of the alkali metalsfamily.)

Characteristics of the Alkaline Earth Metals

Like the alkali metals, the alkaline earth metalshave the properties of a base, as opposed to anacid. The alkaline earth metals are shiny, andmost are white or silvery in color. Like their“cousins” in the alkali metal family, they glowwith characteristic colors when heated. Calciumglows orange, strontium a very bright red, andbarium an apple green. Physically they are soft,though not as soft as the alkali metals, many ofwhich can be cut with a knife.

Yet another similarity the alkaline earth met-als have with the alkali metals is the fact that fourof the them—magnesium, calcium, strontium,and barium—were either identified or isolated inthe first decade of the nineteenth century by Eng-lish chemist Sir Humphry Davy (1778-1829).Around the same time, Davy also isolated sodi-um and potassium from the alkali metal family.

REACTIVITY. The alkaline earth metalsare less reactive than the alkali metals, but like thealkali metals they are much more reactive thanmost elements. Again like their “cousins,” theyreact with water to produce hydrogen gas and the

metal hydroxide, though their reactions are lesspronounced than those of the alkali metals. Mag-nesium metal in its pure form is combustible,and when exposed to air, it burns with an intensewhite light, combining with the oxygen to pro-duce magnesium oxide. Likewise calcium, stron-tium, and barium react with oxygen to formoxides.

Due to their high levels of reactivity, thealkaline earth metals rarely appear by themselvesin nature; rather, they are typically found withother elements in compound form, often as car-bonates or sulfates. This, again, is another simi-larity with the alkali metals. But whereas thealkali metals tend to form 1+ cations (positivelycharged atoms), the alkaline earth metals form2+ cations—that is, cations with a positivecharge of 2.

BOILING AND MELTING

POINTS. One way that the alkaline earth met-als are distinguished from the alkali metals iswith regard to melting and boiling points—thosetemperatures, respectively, at which a solid metalturns into a liquid, and a liquid metal into avapor. For the alkali metals, the temperatures ofthe boiling and melting points decrease with anincrease in atomic number. The pattern is not soclear, however, for the alkaline earth metals.

The highest melting and boiling points arefor beryllium, which indeed has the lowest atom-ic number. It melts at 2,348.6°F (1,287°C), andboils at 4,789.8°F (2,471°C). These figures aremuch higher than for lithium, the alkali metal onthe same period as beryllium, which melts at356.9°F (180.5°C) and boils at 2,457°F (1,347°C).

Magnesium, the second alkali earth metal,melts at 1,202°F (650°C), and boils at 1,994°F(1,090°C)—significantly lower figures than forberyllium. However, the melting and boilingpoints are higher for calcium, third of the alka-line earth metals, with figures of 1,547.6°F(842°C) and 2,703.2°F (1,484°C) respectively.Melting and boiling temperatures steadilydecrease as energy levels rise through strontium,barium, and radium, yet these temperatures arenever lower than for magnesium.

ABUNDANCE. Of the alkaline earthmetals, calcium is the most abundant. It ranksfifth among elements in Earth’s crust, accountingfor 3.39% of the elemental mass. It is also fifthmost abundant in the human body, with a share




of 1.4%. Magnesium, which makes up 1.93% of

Earth’s crust, is the eighth most abundant ele-

ment on Earth. It ranks seventh in the human

body, accounting for 0.50% of the body’s mass.

Barium ranks seventeenth among elements

in Earth’s crust, though it accounts for only

0.04% of the elemental mass. Neither it nor the

other three alkali metals appear in the body in

significant quantities: indeed, barium and beryl-

lium are poisonous, and radium is so radioactive

that exposure to it can be extremely harmful.

Within Earth’s crust, strontium is present in

quantities of 360 parts per million (ppm), which

in fact is rather abundant compared to a number

of elements. In the ocean, its presence is about 8

ppm. By contrast, the abundance of beryllium in

Earth’s crust is measured in parts per billion

(ppb), and is estimated at 1,900 ppb. Vastly more

rare is radium, which accounts for just 0.6 parts

per trillion of Earth’s crust—a fact that made

its isolation by French-Polish physicist and

chemist Marie Curie (1867-1934) all the more

impressive.

MANY ALKALINE EARTH METALS ARE USED IN THE PRODUCTION OF FIREWORKS. (Bill Ross/Corbis. Reproduced by permission.)


AlkalineEarth

Metals


Beryllium

In the eighteenth century, French mineralogistRené Just-Haüy (1743-1822) had observed thatboth emeralds and the mineral beryl had similarproperties. French chemist Louis-NicolasVauquelin (1763-1829) in 1798 identified the ele-ment they had in common: beryllium (Be),which has an atomic number of 4 and an atomicmass of 9.01 amu. Some three decades passedbefore German chemist Friedrich Wöhler (1800-1882) and French chemist Antoine Bussy (1794-1882), working independently, succeeded in iso-lating the element.

Beryllium is found primarily in emeraldsand aquamarines, both precious stones that areforms of the beryllium alluminosilicate com-pound beryl. Though it is toxic to humans, beryl-lium nonetheless has an application in thehealth-care industry: because it lets throughmore x rays than does glass, beryllium is oftenused in x-ray tubes.

Metal alloys that contain about 2% berylli-um tend to be particularly strong, resistant towear, and stable at high temperatures. Copper-beryllium alloys, for instance, are applied in handtools for industries that use flammable solvents,since tools made of these alloys do not causesparks when struck against one another. Alloys ofberyllium and nickel are applied for specializedelectrical connections, as well as for high-tem-perature uses.

Magnesium

English botanist and physician Nehemiah Grew(1641-1712) in 1695 discovered magnesium sul-fate in the springs near the English town ofEpsom, Surrey. This compound, called “Epsomsalts” ever since, has long been noted for itsmedicinal value. Epsom salts are used for treatingeclampsia, a condition that causes seizures inpregnant women. The compound is also a pow-erful laxative, and is sometimes used to rid thebody of poisons—such as magnesium’s sister ele-ment, barium.

For some time, scientists confused the oxidecompound magnesia with lime or calcium car-bonate, which actually involves another alkalineearth metal. In 1754, Scottish chemist and physi-

cist Joseph Black (1728-1799) wrote “Experi-ments Upon Magnesia, Alba, Quick-Lime, andSome Other Alkaline Substances,” an importantwork in which he distinguished between magne-sia and lime. Davy in 1808 declared magnesia theoxide of a new element, which he dubbed mag-nesium, but some 20 years passed before Bussysucceeded in isolating the element.

Magnesium (Mg) has an atomic number of12, and an atomic mass of 24.31 amu. It is foundprimarily in minerals such as dolomite and mag-nesite, both of which are carbonates; and in car-nallite, a chloride. Magnesium silicates includeasbestos, soapstone or talc, and mica. Not allforms of asbestos contain magnesium, but thefact that many do only serves to show the waysthat chemical reactions can change the propertiesan element possesses in isolation.

AN IMPORTANT COMPONENT

OF HEALTH. Whereas magnesium is flam-mable, asbestos was once used in large quantitiesas a flame retardant. And whereas asbestos hasbeen largely removed from public buildingsthroughout the United States due to reports link-ing asbestos exposure with cancer, magnesium isan important component in the health of livingorganisms. It plays a critical role in chlorophyll,the green pigment in plants that captures energyfrom sunlight, and for this reason, it is also usedin fertilizers.

In the human body, magnesium ions(charged atoms) aid in the digestive process, andmany people take mineral supplements contain-ing magnesium, sometimes in combination withcalcium. There is also its use as a laxative, alreadymentioned. Epsom salts, as befits their base oralkaline quality, are exceedingly bitter—the kindof substance a person only ingests under condi-tions of the most dire necessity. On the otherhand, milk of magnesia is a laxative with a far lessunpleasant taste.

MAGNESIUM GOES TO WAR. It isa hallmark of magnesium’s chemical versatilitythat the same element, so important in preserv-ing life, has also been widely used in warfare. Justbefore World War I, Germany was a leading man-ufacturer of magnesium, thanks in large part to amethod of electrolysis developed by Germanchemist R. W. Bunsen (1811-1899). When theUnited States went to war against Germany,American companies began producing magne-sium in large quantities.




Bunsen had discovered that powdered mag-nesium burns with a brilliant white flame, and inthe war, magnesium was used in flares, tracerbullets, and incendiary bombs, which ignite andburn upon impact. The bright light produced byburning magnesium has also led to a number ofpeacetime applications—for instance, in fire-works, and for flashes used in photography.

Magnesium saw service in another worldwar. By the time Nazi tanks rolled into Poland in1939, the German military-industrial complexhad begun using the metal for building aircraftand other forms of military equipment. Americaonce again put its own war-production machineinto operation, dramatically increasing magne-sium output to a peak of nearly 184,000 tons(166,924,800 kg) in 1943.

STRUCTURAL APPLICATIONS.

Magnesium’s principal use in World War I wasfor its incendiary qualities, but in World War II itwas primarily used as a structural metal. It islightweight, but stronger per unit of mass thanany other common structural metal. As a metalfor building machines and other equipment,magnesium ranks in popularity only behind ironand aluminum (which is about 50% more densethan magnesium).

The automobile industry is one area of man-ufacturing particularly interested in magnesium’sstructural qualities. On both sides of the Atlantic,automakers are using or testing vehicle partsmade of alloys of magnesium and other metals,primarily aluminum. Magnesium is easily castinto complex structures, which could mean areduction in the number of parts needed forbuilding a car—and hence a streamlining of theassembly process.

Among the types of sports equipmentemploying magnesium alloys are baseball catch-ers’ masks, skis, racecars, and even horseshoes.Various brands of ladders, portable tools, elec-tronic equipment, binoculars, cameras, furni-ture, and luggage also use parts made of thislightweight, durable metal.

Calcium

Davy isolated calcium (Ca) by means of electrol-ysis in 1808. The element, whose name is derivedfrom the Latin calx, or “lime,” has an atomicnumber of 20, and an atomic mass of 40.08. Theprincipal sources of calcium are limestone and

dolomite, both of which are carbonates, as well asthe sulfate gypsum.

In the form of limestone and gypsum, calci-um has been used as a building material sinceancient times, and continues to find applicationin that area. Lime is combined with clay to makecement, and cement is combined with sand andwater to make mortar. In addition, when mixedwith sand, gravel, and water, cement makes con-crete. Marble—once used to build palaces andtoday applied primarily for decorative touches—also contains calcium.

The steel, glass, paper, and metallurgicalindustries use slaked lime (calcium hydroxide)and quicklime, or calcium oxide. It helps removeimpurities from steel, and pollutants fromsmokestacks, while calcium carbonate in paperprovides smoothness and opacity to the finishedproduct. When calcium carbide (CaC2) is addedto water, it produces the highly flammable gasacetylene (C2H2), used in welding torches. In var-ious compounds, calcium is used as a bleach; amaterial in the production of fertilizers; and as asubstitute for salt as a melting agent on icy roads.

The food, cosmetic, and pharmaceuticalindustries use calcium in antacids, toothpaste,chewing gum, and vitamins. To an even greaterextent than magnesium, calcium is important toliving things, and is present in leaves, bone, teeth,shells, and coral. In the human body, it helps inthe clotting of blood, the contraction of muscles,and the regulation of the heartbeat. Found ingreen vegetables and dairy products, calcium(along with calcium supplements) is recom-mended for the prevention of osteoporosis. Thelatter, a condition involving a loss of bone densi-ty, affects elderly women in particular, and caus-es bones to become brittle and break easily.

Strontium

Irish chemist and physician Adair Crawford(1748-1795) and Scottish chemist and surgeonWilliam Cumberland Cruikshank (1790-1800) in1790 discovered what Crawford called “a newspecies of earth” near Strontian in Scotland. Ayear later, English chemist Thomas Charles Hope(1766-1844) began studying the ore found byCrawford and Cruikshank, which they haddubbed strontia.

In reports produced during 1792 and 1793,Hope explained that strontia could be distin-guished from lime or calcium hydroxide on the



AlkalineEarth

Metals

one hand, and baryta or barium hydroxide onthe other, by virtue of its response to flame tests.Whereas calcium produced a red flame and barium a green one, strontia glowed a brilliantred easily distinguished from the darker red ofcalcium.

Once again, it was Davy who isolated thenew element, using electrolysis, in 1808. Subse-quently dubbed strontium (Sr), its atomic num-ber is 38, and its atomic mass 87.62. Silverywhite, it oxidizes rapidly in air, forming a paleyellow oxide crust on any freshly cut surface.

Though it has properties similar to those ofcalcium, the comparative rarity of strontium andthe expense involved in extracting it offer no eco-nomic incentives for using it in place of its muchmore abundant sister element. Nonetheless,strontium does have a few uses, primarilybecause of its brilliant crimson flame. Thereforeit is applied in the making of fireworks, signalflares, and tracer bullets—that is, rounds thatemit a light as they fly through the air.

One of the more controversial “applications”of strontium involved the radioactive isotopestrontium-90, a by-product of nuclear weaponstesting in the atmosphere from the late 1940s

onward. The isotope fell to earth in a fine pow-der, coated the grass, was ingested by cows, andeventually wound up in the milk they produced.Because of its similarities to calcium, the isotopebecame incorporated into the teeth and gums ofchildren who drank the milk, posing health con-cerns that helped bring an end to atmospherictesting in the early 1960s.

Barium

Aspects of barium’s history are similar to those ofother alkaline earth metals. During the eigh-teenth century, chemists were convinced thatbarium oxide and calcium oxide constituted thesame substance, but in 1774, Swedish chemistCarl Wilhelm Scheele (1742-1786) demonstratedthat barium oxide was a distinct compound.Davy isolated the element, as he did two otheralkaline earth metals, by means of electrolysis,in 1808.

Barium (Ba) has an atomic mass of 137.27and an atomic number of 56. It appears primari-ly in ores of barite, a sulfate, and witherite, a car-bonate. Barium sulfate is used as a white pigmentin paints, while barium carbonate is applied inthe production of optical glass, ceramics, glazed


IN THIS 1960 PHOTO, A MOTHER TESTS HER SON’S MILK FOR SIGNS OF RADIOACTIVITY, THE RESULT OF NUCLEAR

WEAPONS TESTING DURING THE 1950S THAT INVOLVED THE RADIOACTIVE ISOTOPE STRONTIUM-90. THE ISOTOPE

FELL TO EARTH IN A FINE POWDER, WHERE IT COATED THE GRASS, WAS INGESTED BY COWS, AND EVENTUALLY WOUND

UP IN THE MILK THEY PRODUCED. (Bettmann/Corbis. Reproduced by permission.)




pottery, and specialty glassware. One of its mostimportant uses is as a drill-bit lubricant—knownas a “mud” or slurry—for oil drilling. Like anumber of its sister elements, barium (in theform of barium nitrate) is used in fireworks andflares. Motor oil detergents for keeping enginesclean use barium oxide and barium hydroxide.

Beryllium is not the only alkaline earthmetal used in making x rays, nor is magnesiumthe only member of the family applied as a laxa-tive. Barium is used in enemas, and barium sul-fate is used to coat the inner lining of the intes-tines to allow a doctor to examine a patient’sdigestive system. (Though barium is poisonous,in the form of barium sulfate it is safe for inges-tion because the compound does not dissolve inwater or other bodily fluids.) Prior to receiving xrays, a patient may be instructed to drink achalky barium sulfate liquid, which absorbs agreat deal of the radiation emitted by the x-ray

machine. This adds contrast to the black-and-white x-ray photo, enabling the doctor to make amore informed diagnosis.

Radium

Today radium (Ra; atomic number 88; atomicmass 226 amu) has few uses outside of research;nonetheless, the story of its discovery by MarieCurie and her husband Pierre (1859-1906), aFrench physicist, is a compelling chapter not onlyin the history of chemistry, but of humanendeavor in general. Inspired by the discovery ofuranium’s radioactive properties by Frenchphysicist Henri Becquerel (1852-1908), MarieCurie became intrigued with the subject ofradioactivity, on which she wrote her doctoraldissertation. Setting out to find other radioactiveelements, she and Pierre refined a large quantityof pitchblende, an ore commonly found in ura-nium mines. Within a year, they had discovered



valence electron configurations of ns2. The

six alkaline earth metals, all of which are

highly reactive chemically, are beryllium,

magnesium, calcium, strontium, barium,

and radium.

ALKALI METALS: The elements in

Group 1 of the periodic table of elements,

with the exception of hydrogen. The alkali

metals all have one valence electron in the

s1 orbital, and are highly reactive.



All of the alkaline earth metals tend to

form 2+ cations (pronounced KAT-ie-

unz).











unstable—that is, radioactive. Such is the

case with the isotopes of radium, a

radioactive member of the alkaline earth

metals family.



atom in a particular energy state. The six

alkaline earth metals all have valence elec-

trons in an s2 orbital, which describes a

more or less spherical shape.



KEY TERMS


AlkalineEarth

Metals


the element polonium, but were convinced thatanother radioactive ingredient was present—though in much smaller amounts—in pitch-blende.

The Curies spent most of their savings topurchase a ton of ore, and began working toextract enough of the hypothesized Element 88for a usable sample—0.35 oz (1 g). Laboring vir-tually without ceasing for four years, theCuries—by then weary and in financial difficul-ties—finally produced the necessary quantity ofradium. Their fortunes were about to improve: in1903 they shared the Nobel Prize in physics withBecquerel, and in 1911, Marie received a secondNobel, this one in chemistry, for her discoveriesof polonium and radium. She is the only individ-ual in history to win Nobels in two different sci-entific categories.

Because the Curies failed to patent theirprocess, however, they received no profits from

the many “radium centers” that soon sprung up,touting the newly discovered element as a curefor cancer. In fact, as it turned out, the hazardsassociated with this highly radioactive substanceoutweighed any benefits. Thus radium, which atone point was used in luminous paint and onwatch dials, was phased out of use. Marie Curie’sdeath from leukemia in 1934 resulted from herprolonged exposure to radiation from radiumand other elements.

WHERE TO LEARN MORE

“Alkaline Earth Metals.” ChemicalElements.com (Website). <http://www.chemicalelements.com/groups/alkaline.html>> (May 25, 2001).

“The Alkaline Earth Metals” (Web site). <http://www.nidlink.com/~jfromm/elements/alkaline.htm> (May25, 2001).

Ebbing, Darrell D.; R. A. D. Wentworth; and James P.Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.

cipal energy levels in the atoms of the ele-

ments involved.









orbitals.






“rot”; instead, radioactive isotopes contin-

ue changing into other isotopes until they

become stable.




transformed.

SALT: Generally speaking, a compound

that brings together a metal and a non-

metal. More specifically, salts (along with

water) are the product of a reaction

between an acid and a base.

SHELL: A group of electrons within the

same principal energy level.

SUBLEVEL: A region within the princi-

pal energy level occupied by electrons in an

atom. Whatever the number n of the prin-

cipal energy level, there are n sublevels. At

each principal energy level, the first sublev-

el to be filled is the one corresponding to

the s orbital pattern—where the alkaline

earth metals all have their valence electrons.



atom, and which are involved in chemical

bonding.

KEY TERMS CONTINUED






Snedden, Robert. Materials. Des Plaines, IL: Heinemann

Library, 1999.

“Visual Elements: Group 1—The Alkaline Earth Metals”

(Web site). <http://www.chemsoc.org/viselements/

pages/data/intro_groupii_data.html> (May 25,

2001).




TRANSITION METALSTransition Metals

CONCEPTBy far the largest family of elements is the oneknown as the transition metals, sometimes calledtransition elements. These occupy the “dip” inthe periodic table between the “tall” sets ofcolumns or groups on either side. Consisting of10 columns and four rows or periods, the transi-tion metals are usually numbered at 40. With theinclusion of the two rows of transition metals inthe lanthanide and actinide series respectively,however, they account for 68 elements—consid-erably more than half of the periodic table. Thetransition metals include some of the most wide-ly known and commonly used elements, such asiron—which, along with fellow transition metalsnickel and cobalt, is one of only three elementsknown to produce a magnetic field. Likewise,zinc, copper, and mercury are household words,while cadmium, tungsten, chromium, man-ganese, and titanium are at least familiar. Othertransition metals, such as rhenium or hafnium,are virtually unknown to anyone who is not sci-entifically trained. The transition metals includethe very newest elements, created in laboratories,and some of the oldest, known since the earlydays of civilization. Among these is gold, which,along with platinum and silver, is one of severalprecious metals in this varied family.

HOW IT WORKS

Two Numbering Systems forthe Periodic Table

The periodic table of the elements, developed in1869 by Russian chemist Dmitri IvanovitchMendeleev (1834-1907), is discussed in severalplaces within this volume. Within the table, ele-

ments are arranged in columns, or groups, as wellas in periods, or rows.

As discussed in the Periodic Table of Ele-ments essay, two basic versions of the chart, dif-fering primarily in their method of numbergroups, are in use today. The North Americansystem numbers only eight groups—correspon-ding to the representative or main-group ele-ments—leaving the 10 columns representing thetransition metals unnumbered. On the otherhand, the IUPAC system, approved by the Inter-national Union of Pure and Applied Chemistry(IUPAC), numbers all 18 columns. Both versionsof the periodic table show seven periods.

In general, this book applies the NorthAmerican system, in part because it is the versionused in most American classrooms. Additionally,the North American system relates group num-ber to the number of valence electrons in theouter shell of the atom for the element represent-ed. Hence the number of valence electrons (asubject discussed below as it relates to the transi-tion metals) for Group 8 is also eight. Yet wherethe transition metals are concerned, the NorthAmerican system is less convenient than theIUPAC system.

ATTEMPTS TO ADJUST THE

NORTH AMERICAN SYSTEM. Inaddressing the transition metals, there is no easyway to apply the North American system’s corre-spondence between the number of valence elec-trons and the group number. Some versions ofthe North American system do, however, makean attempt. These equate the number of va-lence electrons in the transition metals to a group number, and distinguish these by adding aletter “B.”


TransitionMetals


Thus the noble gases are placed in Group 8Aor VIIIA because they are a main-group familywith eight valence electrons, whereas the transi-tion metal column beginning with iron (Group 8in the IUPAC system) is designated 8B or VIIIB.But this adjustment to the North American sys-tem only makes things more cumbersome,because Group VIIIB also includes two othercolumns—ones with nine and 10 valence elec-trons respectively.

VALUE OF THE IUPAC SYSTEM.

Obviously, then, the North American system—while it is useful in other ways, and as noted, isapplied elsewhere in this book—is not as practi-cal for discussing the transition metals. Ofcourse, one could simply dispense with the term“group” altogether, and simply refer to thecolumns on the periodic table as columns. To doso, however, is to treat the table simply as a chart,rather than as a marvelously ordered means oforganizing the elements.

It makes much more sense to apply theIUPAC system and to treat the transition metalsas belonging to groups 3 through 12. This is par-ticularly useful in this context, because thosegroup numbers do correspond to the numbers ofvalence electrons. Scandium, in group 3 (hence-forth all group numbers refer to the IUPAC ver-

sion), has three valence electrons; likewise zinc,in Group 12, has 12. Neither the IUPAC nor theNorth American system provide group numbersfor the lanthanides and actinides, known collec-tively as the inner transition metals.

THE TRANSITION METALS BY

IUPAC GROUP NUMBER. Here, then, isthe list of the transition metals by group number,as designated in the IUPAC system. Under eachgroup heading are listed the four elements in thatgroup, along with the atomic number, chemicalsymbol, and atomic mass figures for each, inatomic mass units.

Note that, because the fourth member ineach group is radioactive and sometimes existsfor only a few minutes or even seconds, mass fig-ures are given merely for the most stable isotope.In addition, the last elements in groups 8, 9, and10, as of 2001, had not received official names.For reasons that will be explained below, lan-thanum and actinium, though included here, arediscussed in other essays.

Group 1

• 21. Scandium (Sc): 44.9559

• 39. Yttrium (Y): 88.9059

• 57. Lanthanum (La): 138.9055

• 89. Actinium (Ac): 227.0278

SEVERAL PRECIOUS METALS, INCLUDING GOLD, ARE AMONG THE TRANSITION METALS. (Charles O’Rear/Corbis. Reproduced

by permission.)


TransitionMetals

Group 2

• 22. Titanium (Ti): 47.9

• 40. Zirconium (Zr): 91.22

• 72. Hafnium (Hf): 178.49

• 104. Rutherfordium (Rf): 261

Group 3

• 23. Vanadium (V): 50.9415

• 41. Niobium (Nb): 92.9064

• 73. Tantalum (Ta): 180.9479

• 105. Dubnium (Db): 262

Group 4

• 24. Chromium (Cr): 51.996

• 42. Molybdenum (Mo): 95.95

• 74. Tungsten (W): 183.85

• 106. Seaborgium (Sg): 263

Group 5

• 25. Manganese (Mn): 54.938

• 43. Technetium (Tc): 98

• 75. Rhenium (Re): 186.207

• 107. Bohrium (Bh): 262

Group 6

• 26. Iron (Fe): 55.847

• 44. Ruthenium (Ru): 101.07

• 76. Osmium (Os): 190.2

• 108. Hassium (Hs): 265

Group 7

• 27. Cobalt (Co): 58.9332

• 45. Rhodium (Rh): 102.9055

• 77. Iridium (Ir): 192.22

• 109. Meitnerium (Mt): 266

Group 8

• 28. Nickel (Ni): 58.7

• 46. Palladium (Pd): 106.4

• 78. Platinum (Pt): 195.09

• 110. Ununnilium (Uun): 271

Group 9

• 29. Copper (Cu): 63.546

• 47. Silver (Ag): 107.868

• 79. Gold (Au): 196.9665

• 111. Unununium (Uuu): 272

Group 10

• 30. Zinc (Zn): 65.38

• 48. Cadmium (Cd): 112.41

• 80. Mercury (Hg): 200.59

• 112. Ununbium (Uub): 277


Electron Configurations andthe Periodic Table

We can now begin to explain why transition met-

als are distinguished from other elements, and

why the inner transition metals are further sepa-

rated from the transition metals. This distinction

relates not to period (row), but to group (col-

umn)—which, in turn, is defined by the configu-

ration of valence electrons, or the electrons that

are involved in chemical bonding. Valence elec-

trons occupy the highest energy level, or shell,

of the atom—which might be thought of as

the orbit farthest from the nucleus. However, the

term “orbit” is misleading when applied to

the ways that an electron moves.

Electrons do not move around the nucleus

of an atom in regular orbits, like planets around

the Sun; rather, their paths can only be loosely

defined in terms of orbitals, a pattern of proba-

bilities indicating the regions that an electron

may occupy. The shape or pattern of orbitals is

determined by the principal energy level of the

AN IRON ORE MINE IN SOUTH AFRICA. IRON IS THE

FOURTH MOST ABUNDANT ELEMENT ON EARTH,ACCOUNTING FOR 4.71% OF THE ELEMENTAL MASS IN

THE PLANET’S CRUST. (Charles O’Rear/Corbis. Reproduced by

permission.)


TransitionMetals


atom, which indicates a distance that an electronmay move away from the nucleus.

Principal energy level defines period: ele-ments in Period 4, for instance, all have theirvalence electrons at principal energy level 4. Thehigher the number, the further the electron isfrom the nucleus, and hence, the greater theenergy in the atom. Each principal energy level isdivided into sublevels corresponding to thenumber n of the principal energy level: thus,principal energy level 4 has four sublevels, prin-cipal energy level 5 has five, and so on.

ORBITAL PATTERNS. The four basictypes of orbital patterns are designated as s, p, d,and f. The s shape might be described as spheri-cal, which means that in an s orbital, the totalelectron cloud will probably end up being moreor less like a sphere.

The p shape is like a figure eight around thenucleus; the d like two figure eights meeting atthe nucleus; and the f orbital pattern is so com-plex that most basic chemistry textbooks do noteven attempt to explain it. In any case, theseorbital patterns do not indicate the exact path ofan electron. Think of it, instead, in this way: ifyou could take millions of photographs of theelectron during a period of a few seconds, the

resulting blur of images in a p orbital, forinstance, would somewhat describe the shape ofa figure eight.

Since the highest energy levels of the transi-tion metals begin at principal energy level 4, wewill dispense with a discussion of the first threeenergy levels. The reader is encouraged to consultthe Electrons essay, as well as the essay on Fami-lies of Elements, for a more detailed discussion ofthe ways in which these energy levels are filled forelements on periods 1 through 3.

Orbital Filling and the Periodic Table

If all elements behaved as they “should,” the pat-tern of orbital filling would be as follows. In agiven principal energy level, first the two slotsavailable to electrons on the s sublevel are filled,then the six slots on the p sublevel, then the 10slots on the d sublevel. The f sublevel, which isfourth to be filled, comes into play only at prin-cipal energy level 4, and it would be filled beforeelements began adding valence electrons to the ssublevel on the next principal energy level.

That might be the way things “should” hap-pen, but it is not the way that they do happen.The list that follows shows the pattern by which

THE AMERICAN FIVE-CENT COIN, CALLED A “NICKEL,” IS ACTUALLY AN ALLOY OF NICKEL AND COPPER. (Layne

Kennedy/Corbis. Reproduced by permission.)


TransitionMetals

orbitals are filled from sublevel 3p onward. Fol-lowing each sublevel, in parentheses, is the num-ber of “slots” available for electrons in that shell.Note that in several places, the pattern of fillingdefies the ideal order described in the precedingparagraph.

Orbital Filling by Principal Energy Level andSublevel from 3p Onward:

• 3p (6)

• 4s (2)

• 3d (10)

• 4p (6)

• 5s (2)

• 4d (10)

• 5p (6)

• 6s (2)

• 4f (14)

• 5d (10)

• 6p (6)

• 7s (2)

• 5f (14)

• 6d (10)

The 44 representative elements follow a reg-ular pattern of orbital filling through the first 18elements. By atomic number 18, argon, 3p hasbeen filled, and at that point element 19 (potas-sium) would be expected to begin filling row 3d.However, instead it begins filling 4s, to which cal-cium (20) adds the second and last valence elec-tron. It is here that we come to the transitionmetals, distinguished by the fact that they fill thed orbitals.

Note that none of the representative ele-ments up to this point, or indeed any that follow,fill the d orbital. (The d orbital could not comeinto play before Period 3 anyway, because at leastthree sublevels—corresponding to principalenergy level 3—are required in order for there tobe a d orbital.) In any case, when the representa-tive elements fill the p orbitals of a given energylevel, they skip the d orbital and go on to the nextprincipal energy level, filling the s orbitals. Fillingof the d orbital on the preceding energy level onlyoccurs with the transition metals: in other words,on period 4, transition metals fill the 3d sublevel,and so on as the period numbers increase.

Distinguishing the TransitionMetals

Given the fact that it is actually the representativeelements that skip the d sublevels, and the transi-

tion metals that go back and fill them, one mightwonder if the names “representative” and “transi-tion” (implying an interruption) should bereversed. But it is the transition metals that arethe exception to the rule, for two reasons. First ofall, as we have seen, they are the only elementsthat fill the d orbitals.

Secondly, they are the only elements whoseouter-shell electrons are on two different prin-cipal energy levels. The orbital filling patterns of transition metals can be identified thus:ns(n-1)d, where n is the number of the period onwhich the element is located. For instance, zinc,on period 4, has an outer shell of 4s23d10. In otherwords, it has two electrons on principal energylevel 4, sublevel 4s—as it “should.” But it also has10 electrons on principal energy level 3, sublevel3d. In effect, then, the transition metals “go back”and fill in the d orbitals of the preceding period.

There are further complications in the pat-terns of orbital filling for the transition metals,which will only be discussed in passing here as afurther explanation as to why these elements arenot “representative.” Most of them have two svalence electrons, but many have one, and palla-dium has zero. Nor do they add their d electronsin regular patterns. Moving across period 4, forinstance, the number of electrons in the d orbital“should” increase in increments of 1, from 1 to10. Instead the pattern goes like this: 1, 2, 3, 5, 5,6, 7, 8, 10, 10.

LANTHANIDES AND ACTI-

NIDES. The lanthanides and actinides are setapart even further from the transition metals. Inmost versions of the periodic table, lanthanum(57) is followed by hafnium (72) in the transitionmetals section of the chart. Similarly, actinium(89) is followed by rutherfordium (104). The“missing” metals—lanthanides and actinidesrespectively—are listed at the bottom of thechart. There are reasons for this, as well as for thenames of these groups.

After the 6s orbital fills with the representa-tive element barium (56), lanthanum does whata transition metal does—it begins filling the 5dorbital. But after lanthanum, something strangehappens: cerium (58) quits filling 5d, and movesto fill the 4f orbital. The filling of that orbitalcontinues throughout the entire lanthanideseries, all the way to lutetium (71). Thus lan-thanides can be defined as those metals that fillthe 4f orbital; however, because lanthanum



TransitionMetals

exhibits similar properties, it is usually includedwith the lanthanides.

A similar pattern occurs for the actinides.The 7s orbital fills with radium (88), after whichactinium (89) begins filling the 6d orbital. Nextcomes thorium, first of the actinides, whichbegins the filling of the 5f orbital. This is com-pleted with element 103, lawrencium. Actinidescan thus be defined as those metals that fill the 5forbital; but again, because actinium exhibits sim-ilar properties, it is usually included with theactinides.


Survey of the TransitionMetals

The fact that the transition elements are all met-als means that they are lustrous or shiny inappearance, and malleable, meaning that theycan be molded into different shapes withoutbreaking. They are excellent conductors of heatand electricity, and tend to form positive ions bylosing electrons.

Generally speaking, metals are hard, thougha few of the transition metals—as well as mem-bers of other metal families—are so soft they canbe cut with a knife. Like almost all metals, theytend to have fairly high melting points, andextremely high boiling points.

Many of the transition metals, particularlythose on periods 4, 5, and 6, form useful alloys—mixtures containing more than one metal—withone another, and with other elements. Because oftheir differences in electron configuration, how-ever, they do not always combine in the sameways, even within an element. Iron, for instance,sometimes releases two electrons in chemicalbonding, and at other times three.

ABUNDANCE OF THE TRANSI-

TION METALS. Iron is the fourth mostabundant element on Earth, accounting for4.71% of the elemental mass in the planet’s crust.Titanium ranks 10th, with 0.58%, and man-ganese 13th, with 0.09%. Several other transitionmetals are comparatively abundant: even gold ismuch more abundant than many other elementson the periodic table. However, given the factthat only 18 elements account for 99.51% of

Earth’s crust, the percentages for elements out-side of the top 18 tend to very small.

In the human body, iron is the 12th mostabundant element, constituting 0.004% of thebody’s mass. Zinc follows it, at 13th place,accounting for 0.003%. Again, these percentagesmay not seem particularly high, but in view ofthe fact that three elements—oxygen, carbon,and hydrogen—account for 93% of human ele-mental body mass, there is not much room forthe other 10 most common elements in the body.Transition metals such as copper are present intrace quantities within the body as well.

DIVIDING THE TRANSITION

METALS INTO GROUPS. There is noeasy way to group the transition metals, thoughcertain of these elements are traditionally catego-rized together. These do not constitute “families”as such, but they do provide useful ways to breakdown the otherwise rather daunting 40-elementlineup of the transition metals.

In two cases, there is at least a relationbetween group number on the periodic table andthe categories loosely assigned to a collection oftransition metals. Thus the “coinage metals”—copper, silver, and gold—all occupy Group 9 onthe periodic table. These have traditionally beenassociated with one another because their resist-ance to oxidation, combined with their mal-leability and beauty, has made them useful mate-rials for fashioning coins.

Likewise the members of the “zinc group”—zinc, cadmium, and mercury—occupy Group 10on the periodic table. These, too, have often beenassociated as a miniature unit due to commonproperties. Members of the “platinum group”—platinum, iridium, osmium, palladium, rhodi-um, and ruthenium—occupy a rectangle on thetable, corresponding to periods 5 and 6, andgroups 6 through 8. What actually makes them a“group,” however, is the fact that they tend toappear together in nature.

Iron, nickel, and cobalt, found alongside oneanother on Period 4, may be grouped togetherbecause they are all magnetic to some degree oranother. This is far from the only notable charac-teristic about such metals, but provides a conven-ient means of further dividing the transitionmetals into smaller sections.

To the left of iron on the periodic table is arectangle corresponding to periods 4 through 6,groups 4 through 7. These 11 elements—titani-



TransitionMetals

um, zirconium, hafnium, vanadium, niobium,tantalum, chromium, molybdenum, tungsten,manganese, and rhenium—are referred to hereas “alloy metals.” This is not a traditional desig-nation, but it is nonetheless useful for describingthese metals, most of which form importantalloys with iron and other elements.

One element was left out of the “rectangle”described in the preceding paragraph. This istechnetium, which apparently does not occur innature. It is lumped in with a final category, “rareand artificial elements.”

It should be stressed that there is nothinghard and fast about these categories. The “alloymetals” are not the only ones that form alloys;nickel is used in coins, though it is not called acoinage metal; and platinum could be listed withgold and silver as “precious metals.” Nonetheless,the categories used here seem to provide themost workable means of approaching the manytransition metals.

Coinage Metals

GOLD. Gold almost needs no introduc-tion: virtually everyone knows of its value, andhistory is full of stories about people who killedor died for this precious metal. Part of its valuesprings from its rarity in comparison to, say iron:gold is present on Earth’s crust at a level of about5 parts per billion (ppb). Yet as noted earlier, it ismore abundant than some metals. Furthermore,due to the fact that it is highly unreactive (reac-tivity refers to the tendency for bonds betweenatoms or molecules to be made or broken in sucha way that materials are transformed), it tends tobe easily separated from other elements.

This helps to explain the fact that gold maywell have been the first element ever discovered.No ancient metallurgist needed a laboratory inwhich to separate gold; indeed, because it sooften keeps to itself, it is called a “noble” metal—meaning, in this context, “set apart.” Anothercharacteristic of gold that made it valuable wasits great malleability. In fact, gold is the mostmalleable of all metals: A single troy ounce (31.1g) can be hammered into a sheet just 0.00025 in(0.00064 cm) thick, covering 68 ft2 (6.3 m2).

Gold is one of the few metals that is not sil-ver, gray, or white, and its beautifully distinctivecolor caught the eyes of metalsmiths and royaltyfrom the beginning of civilization. Records fromIndia dating back to 5000 B.C. suggest a familiar-

ity with gold, and jewelry found in Egyptiantombs indicates the use of sophisticated tech-niques among the goldsmiths of Egypt as early as2600 B.C. Likewise the Bible mentions gold inseveral passages. The Romans called it aurum(“shining dawn”), which explains its chemicalsymbol, Au.

Early chemistry had its roots among themedieval alchemists, who sought in vain to turnplain metals into gold. Likewise the history ofEuropean conquest in the New World is filledwith bitter tales of conquistadors, lured by gold,who wiped out entire nations of native peoples.

At one time, gold was used in coins, andnations gauged the value of their currency interms of the gold reserves they possessed. Fewcoins today—with special exceptions, such as theSouth African Krugerrand—are gold. Likewisethe United States, for instance, went off the goldstandard (that is, ceased to tie its currency to goldreserves) in the 1930s; nonetheless, the federalgovernment maintains a vast supply of gold bul-lion at Fort Knox in Kentucky.

Gold is as popular as ever for jewelry andother decorative objects, of course, but for themost part, it is too soft to have many other com-mercial purposes. One of the few applications forgold, a good conductor of electricity, is in someelectronic components. Also, the radioactivegold-198 isotope is sometimes implanted in tis-sues as a means of treating forms of cancer.

SILVER. Like gold, silver has been a partof human life from earliest history. Usually it isconsidered less valuable, though some societieshave actually placed a higher value on silverbecause it is harder and more durable than gold.In the seventh century B.C., the Lydian civiliza-tion of Asia Minor (now Turkey) created the firstcoins using silver, and in the sixth century B.C.,the Chinese began making silver coins. Succeed-ing dynasties in China continued to mint thesecoins, round with square holes in them, until theearly twentieth century.

The Romans called silver argentum, andtherefore today its chemical symbol is Ag. Its usesare much more varied than those of gold, bothbecause of its durability and the fact that it is lessexpensive. Alloyed with copper, which addsstrength to it, it makes sterling silver, used incoins, silverware, and jewelry. Silver nitrate com-pounds are used in silver plating, applied in mir-



TransitionMetals

rors and tableware. (Most mirrors today, howev-er, use aluminum.)

Like gold, silver was once widely used forcoinage, and silver coins are still more commonthan their gold counterparts. In 1963, however,the United States withdrew its silver certifica-tions, paper currency exchangeable for silver.Today, silver coins are generally issued in specialsituations, such as to commemorate an event.

A large portion of the world’s silver supply isused by photographers for developing pictures.In addition, because it is an excellent conductorof heat and electricity, silver has applications inthe electronics industry; however, its expense hasled many manufacturers to use copper or alu-minum instead. Silver is also present, along withzinc and cadmium, in cadmium batteries. Likegold, though to a much lesser extent, it is still animportant jewelry-making component.

COPPER. Most people think of penniesas containing copper, but in fact the penny is theonly American coin that contains no copperalloys. Because the amount of copper necessaryto make a penny today costs more than $0.01, apenny is actually made of zinc with a thin coppercoating. Yet copper has long been a commonlyused coinage metal, and long before that,humans used it for other purposes.

Seven thousand years ago, the peoples of theTigris-Euphrates river valleys, in what is nowIraq, were mining and using copper, and latercivilizations combined copper with zinc to makebronze. Indeed, the history of prehistoric andancient humans’ technological development isoften divided according to the tools they made,the latter two of which came from transitionmetals: the Stone Age, the Bronze Age (c. 3300-1200 B.C.), and the Iron Age.

Thus, by the time the Romans dubbed cop-per cuprum (from whence its chemical symbol,Cu, comes), copper in various forms had longbeen in use. Today copper is purified by a form ofelectrolysis, in which impurities such as iron,nickel, arsenic, and zinc are removed. Other“impurities” often present in copper ore includegold, silver, and platinum, and the separation ofthese from the copper pays for the large amountsof electricity used in the electrolytic process.

Like the other coinage metals, copper is anextremely efficient conductor of heat and elec-tricity, and because it is much less expensive thanthe other two, pure copper is widely used for

electrical wiring. Because of its ability to conductheat, copper is also applied in materials used formaking heaters, as well as for cookware. Due tothe high conductivity of copper, a heated copperpan has a uniform temperature, but copper potsmust be coated with tin because too much cop-per in food is toxic.

Copper is also like its two close relatives inthat it resists corrosion, and this makes it idealfor plumbing. Its use in making coins resultedfrom its anti-corrosive qualities, combined withits beauty: like gold, copper has a distinctivecolor. This aesthetic quality led to the use of cop-per in decorative applications as well: many oldbuildings used copper roofs, and the Statue ofLiberty is covered in 300 thick copper plates.

Why, then, is the famous statue not copper-colored? Because copper does eventually corrodewhen exposed to air for long periods of time.Over time, it develops a thin layer of black cop-per oxide, and as the years pass, carbon dioxide inthe air leads to the formation of copper carbon-ate, which imparts a greenish color.

The human body is about 0.0004% copper,though as noted, larger quantities of copper canbe toxic. Copper is found in foods such as shell-fish, nuts, raisins, and dried beans. Whereashuman blood has hemoglobin, a molecule withan iron atom at the center, the blood of lob-sters and other large crustaceans contains hemocyanin, in which copper performs a similarfunction.

The Zinc Group

ZINC. Together with copper, zincappeared in another alloy that, like bronze,helped define the ancient world: brass. (The lat-ter is mentioned in the Bible, for instance in theBook of Daniel, when King Nebuchadnezzardreams of a statue containing brass and othersubstances, symbolizing various empires.) Usedat least from the first millennium B.C. onward,brass appeared in coins and ornaments through-out Asia Minor. Though it is said that the Chi-nese purified zinc in about A.D. 1000, the Swissalchemist Paracelsus (1493-1541) is usually cred-ited with first describing zinc as a metal.

Bluish-white, with a lustrous sheen, zinc isfound primarily in the ore sulfide sphalerite. Thelargest natural deposits of zinc are in Australiaand the United States, and after mining, themetal is subjected to a purification and reduction



TransitionMetals

process involving carbon. Zinc is used in galva-nized steel, developed in the eighteenth centuryby Italian physicist Luigi Galvani (1737-1798).

Just as a penny is not really copper but zinc,“tin” roofs are usually made of galvanized steel.Highly resistant to corrosion, galvanized steelfinds application in everything from industrialequipment to garbage cans. Zinc oxide is appliedin textiles, batteries, paints, and rubber products,while luminous zinc sulfide appears in televisionscreens, movie screens, clock dials, and fluores-cent light bulbs.

Zinc phosphide is used as a rodent poison.Like several other transition metals, zinc is a partof many living things, yet it can be toxic in largequantities or specific compounds. For a humanbeing, inhaling zinc oxide causes involuntaryshaking. On the other hand, humans and manyanimals require zinc in their diets for the diges-tion of proteins. Furthermore, it is believed thatzinc contributes to the healing of wounds and tothe storage of insulin in the pancreas.

CADMIUM. In 1817, German chemistFriedrich Strohmeyer (1776-1835) was workingas an inspector of pharmacies for the Germanstate of Hanover. While making his rounds, hediscovered that one pharmacy had a sample ofzinc carbonate labeled as zinc oxide, and whileinspecting the chemical in his laboratory, he dis-covered something unusual. If indeed it werezinc carbonate, it should turn into zinc oxidewhen heated, and since both compounds werewhite, there should be no difference in color.Instead, the mysterious compound turned a yel-lowish-orange.

Strohmeyer continued to analyze the sam-ple, and eventually realized that he had discov-ered a new element, which he named after the oldGreek term for zinc carbonate, kadmeia. Indeed,cadmium typically appears in nature along withzinc or zinc compounds. Silvery white and lus-trous or shiny, cadmium is soft enough to be cutwith a knife, but chemically it behaves much likezinc: hence the idea of a “zinc group.”

Today cadmium is used in batteries, and forelectroplating of other metals to protect themagainst corrosion. Because the cost of cadmiumis high due to the difficulty of separating it fromzinc, cadmium electroplating is applied only inspecialized situations. Cadmium also appears inthe control rods of nuclear power plants, where

its ready absorption of neutrons aids in control-ling the rate at which nuclear fission occurs.

Cadmium is highly toxic, and is believed tobe the cause behind the outbreak of itai-itai(“ouch-ouch”) disease in Japan in 1955. Peopleingested rice oil contaminated with cadmium,and experienced a number of painful side effectsassociated with cadmium poisoning: nausea,vomiting, choking, diarrhea, abdominal pain,headaches, and difficulty breathing.

MERCURY. One of only two elements—along with bromine—that appears in liquid format room temperature, mercury is both toxic andhighly useful. The Romans called it hydragyrum(“liquid silver”), from whence comes its chemicalsymbol, Hg. Today, however, it is known by thename of the Romans’ god Mercury, the nimbleand speedy messenger of the gods. Mercurycomes primarily from a red ore called cinnabar,and since it often appears in shiny globules thatform outcroppings from the cinnabar, it was rel-atively easy to discover.

Several things are distinctive about mercury,including its bright silvery color. But nothing dis-tinguishes it as much as its physical properties—not only its liquidity, but the fact that it rolls rap-idly, like the fleet-footed god after which it isnamed. Its surface tension (the quality that caus-es it to bead) is six times greater than that ofwater, and for this reason, mercury never wetsthe surfaces with which it comes in contact.

Mercury, of course, is widely used in thermometers, an application for which it isextremely well-suited. In particular, it expands ata uniform rate when heated, and thus a mercurythermometer (unlike earlier instruments, whichused water, wine, or alcohol) can be easily cali-brated. (Note that due to the toxicity of the element, mercury thermometers in schools arebeing replaced by other types of thermometers.)At temperatures close to absolute zero, mercuryloses its resistance to the flow of electric current,and therefore it presents a promising area ofresearch with regard to superconductivity.

Even with elements present in the humanbody, such as zinc, there is a danger of toxicitywith exposure to large quantities. Mercury, onthe other hand, does not occur naturally in thehuman body, and is therefore extremely danger-ous. Because it is difficult for the body to elimi-nate, even small quantities can accumulate andexert their poisonous effects, resulting in disor-



TransitionMetals

ders of the nervous system. In fact, the phrase“mad as a hatter” refers to the symptoms of mer-cury poisoning suffered by hatmakers of thenineteenth century, who used a mercury com-pound in making hats of beaver and fur.

Magnetic Metals

IRON. In its purest form, iron is relativelysoft and slightly magnetic, but when hardened, itbecomes much more so. As with several of theelements discovered long ago, iron has a chemi-cal symbol (Fe) reflecting an ancient name, theLatin ferrum. But long before the Romans’ ances-tors arrived in Italy, the Hittites of Asia Minorwere purifying iron ore by heating it with char-coal over a hot flame.

The Hittites jealously guarded their secret,which gave them a technological advantage overenemies such as the Egyptians. But when Hittitecivilization crumbled in the wake of an invasionby the mysterious “Sea Peoples” in about 1200B.C., knowledge of ore smelting spread through-out the ancient world. Thus began the Iron Age,in which ancient technology matured greatly. (Itshould be noted that the Bantu peoples, in whatis now Nigeria and neighboring countries, devel-oped their own ironmaking techniques a fewcenturies later, apparently without benefit ofcontact with any civilization exposed to Hittitetechniques.)

Ever since ancient times, iron has been avital part of human existence, and with thegrowth of industrialization during the eighteenthcentury, demand for iron only increased. Butheating iron ore with charcoal required the cut-ting of trees, and by the latter part of that centu-ry, England’s forests had been so badly denudedthat British ironmakers sought a new means ofsmelting the ore. It was then that British inventorAbraham Darby (1678?-1717) developed amethod for making coke from soft coal, whichwas abundant throughout England. Adoption ofDarby’s technique sped up the rate of industrial-ization in England, and thus advanced the entireIndustrial Revolution soon to sweep the Westernworld.

Symbolic of industrialism was iron in itsmany forms: pig iron, ore smelted in a coke-burning blast furnace; cast iron, a variety of mix-tures containing carbon and/or silicon; wroughtiron, which contains small amounts of variousother elements, such as nickel, cobalt, copper,

chromium, and molybdenum; and most of allsteel. Steel is a mixture of iron with manganeseand chromium, purified with a blast of hot air inthe Bessemer converter, named after Englishinventor Henry Bessemer (1813-1898). Moderntechniques of steelmaking improved on Besse-mer’s design by using pure oxygen rather thanmerely hot air.

The ways in which iron is used are almosttoo obvious (and too numerous) to mention. Ifiron and steel suddenly ceased to exist, therecould be no skyscrapers, no wide-span bridges,no ocean liners or trains or heavy machinery orautomobile frames. Furthermore, alloys of steelwith other transition metals, such as tungstenand niobium, possess exceptionally greatstrength, and find application in everything fromhand tools to nuclear reactors. Then, of course,there are magnets and electromagnets, which canonly be made of iron and/or one of the othermagnetic elements, cobalt and nickel.

In the human body, iron is a key part ofhemoglobin, the molecule in blood that trans-ports oxygen from the lungs to the cells. If a per-son fails to get sufficient quantities of iron—present in foods such as red meat and spinach—the result is anemia, characterized by a loss ofskin color, weakness, fainting, and heart palpita-tions. Plants, too, need iron, and without theappropriate amounts are likely to lose their color,weaken, and die.

COBALT. Isolated in about 1735 bySwedish chemist Georg Brandt (1694-1768),cobalt was the first metal discovered since prehis-toric, or at least ancient, times. The name comesfrom Kobald, German for “underground gnome,”and this reflects much about the early history ofcobalt. In legend, the Kobalden were mischievoussprites who caused trouble for miners, and in reallife, ores containing the element that came to beknown as cobalt likewise caused trouble to menworking in mines. Not only did these ores con-tain arsenic, which made miners ill, but becausecobalt had no apparent value, it only interferedwith their work of extracting other minerals.

Yet cobalt had been in use by artisans longbefore Brandt’s isolated the element. The color ofcertain cobalt compounds is a brilliant, shockingblue, and this made it popular for the coloring ofpottery, glass, and tile. The element, which makesup less than 0.002% of Earth’s crust, is foundtoday primarily in ores extracted from mines in



TransitionMetals

Canada, Zaire, and Morocco. One of the mostimportant uses of cobalt is in a highly magneticalloy known as alnico, which also contains iron,nickel, and aluminum. Combined with tungstenand chromium, cobalt makes stellite, a very hardalloy used in drill bits. Cobalt is also applied in jetengines and turbines.

NICKEL. Moderately magnetic in its pureform, nickel had an early history much like thatof cobalt. English workers mining copper wereoften dismayed to find a metal that looked likecopper, but was not, and they called it “Old Nick’scopper”—meaning that it was a trick played onthem by Old Nick, or the devil. The Germansgave it a similar name: Kupfernickel, or “imp cop-per.”

Though nickel was not identified as a sepa-rate metal by Swedish mineralogist Axel FredrikCronstedt (1722-1765) until the eighteenth cen-tury, alloys of copper, silver, and nickel had beenused as coins even in ancient Egypt. Today, nick-el is applied, not surprisingly, in the Americanfive-cent piece—that is, the “nickel”—made froman alloy of nickel and copper. Its anti-corrosivenature also provides a number of other applica-tions for nickel: alloyed with steel, for instance, itmakes a protective layer for other metals.

The Platinum Group

PLATINUM. First identified by an Italianphysician visiting the New World in the mid-six-teenth century, platinum—now recognized as aprecious metal—was once considered a nuisancein the same way that nickel and cadmium were.Miners, annoyed with the fact that it got in theway when they were looking for gold, called itplatina, or “little silver.” One of the reasons whyplatinum did not immediately catch the world’sfancy is because it is difficult to extract, and typ-ically appears with the other metals of the “plat-inum group”: iridium, osmium, palladium,rhodium, and ruthenium.

Only in 1803 did English physician andchemist William Hyde Wollaston (1766-1828)develop a means of extracting platinum, andwhen he did, he discovered that the metal couldbe hammered into all kinds of shapes. Platinumproved such a success that it made Wollastonfinancially independent, and he retired from hismedical practice at age 34 to pursue scientificresearch. Today, platinum is used in everythingfrom thermometers to parts for rocket engines,

both of which take advantage of its ability towithstand high temperatures.

IRIDIUM AND OSMIUM. In thesame year that Wollaston isolated platinum, threeFrench chemists applied a method of extractingit with a mixture of nitric and hydrochloric acids.The process left a black powder that they wereconvinced was a new element, and in 1804 Eng-lish chemist Smithson Tennant (1761-1815) gaveit the name iridium. Iris was the Greek goddess ofthe rainbow, and the name reflected the ele-ment’s brilliant, multi-colored sheen. As withplatinum, iridium is used today in a number ofways, particularly for equipment such as lasercrystals that must withstand high tempera-tures. (In addition, one particular platinum-iridium bar provides the standard measure of akilogram.)

Tennant discovered a second element in1804, also from the residue left over from the acidprocess for extracting platinum. This one had adistinctive smell when heated, so he named itosmium after osme, Greek for “odor.” In 1898,Austrian chemist Karl Auer, Baron von Welsbach(1858-1929), developed a light bulb using osmi-um as a filament, the material that is heated.Though osmium proved too expensive for com-mercial use, Auer’s creation paved the way for theuse of another transition metal, tungsten, inmaking long-lasting filaments. Osmium, which isvery hard and resistant to wear, is also used inelectrical devices, fountain-pen tips, and phono-graph needles.

PALLADIUM, RHODIUM, AND

RUTHENIUM. By the time Tennant isolatedosmium, Wollaston had found yet another ele-ment that emerged from the refining of platinumwith nitric and hydrochloric acids. This was pal-ladium, named after a recently discovered aster-oid, Pallas. Soft and malleable, palladium resiststarnishing, which makes it valuable to the jewel-ry industry. Combined with yellow gold, it makesthe alloy known as white gold. Because palladiumhas the unusual ability of absorbing up to 900times its volume in hydrogen, it provides a usefulmeans of extracting that element.

Also in 1805, Wollaston discovered rhodi-um, which he named because it had a slightly redcolor. The density of rhodium is lower, and themelting point higher, than for most of the plat-inum group elements, and it is often used as analloy to harden platinum. As with many elements



TransitionMetals

in this group, it has a number of high-tempera-ture applications. When electroplated, it is highlyreflective, and this makes it useful as a materialfor optical instruments.

The element that came to be known asruthenium had been detected several timesbefore 1844, when Russian chemist Carl ErnestClaus (1796-1864) finally identified it as a dis-tinct element. Thus it was the only platinumgroup element in whose isolation neither Ten-nant nor Wollaston played a leading role. Since ithad been found in Russia, the element was calledruthenium, in honor of that country’s ancientname, Ruthenia.

Alloy Metals

The elements described here as “alloy metals” willbe treated briefly, but the amount of space givento them here does not reflect their importance tocivilization. Several of these are relatively abun-dant, and many—titanium, tungsten, man-ganese, chromium and others—are vital parts ofdaily life. However, due to space considerations,the treatment of these metals that follows is aregrettably abbreviated one.

As suggested by the informal name given tothem here, they are often applied in alloys,typically with other transition metals such asiron, or with other “alloy metals.” They are alsoused in a variety of compounds for a wide vari-ety of applications.

Titanium, Zirconium, andHafnium

English clergyman and amateur scientist WilliamGregor (1761-1817) in 1791 noticed what hecalled “a new metallic substance” in a mineralcalled menchanite, which he found in the Men-achan region of Cornwall, England. Four yearslater, German chemist Martin Heinrich Klaproth(1743-1817) studied menchanite and found thatit contained a new element, which he named tita-nium after the Titans of Greek mythology. Onlyin 1910, however, was the element isolated. Dueto its high strength-to-weight ratio, titanium isapplied today in alloys used for making airplanesand spacecraft. It is also used in white paints andin a variety of other combinations, mostly inalloys containing other transition metals, as wellas in compounds with nonmetals such as oxygen.

Six years before he identified titanium, in1789, Klaproth had found another element,which he named zirconium after the mineral zir-con in which it was found. Zircon had been inuse for many centuries, and another zirconium-containing mineral, jacinth, is mentioned in theBible. As with titanium, there was a long periodfrom identification to isolation, which did notoccur until 1914. Zirconium alloys are used todayin a number of applications, including supercon-ducting magnets and as a construction metal innuclear power plants. Compounds including zir-conium also have a variety of uses, ranging fromfurnace linings to poison ivy lotion.

Because it is almost always found with zirco-nium, hafnium—named after the ancient nameof Copenhagen, Denmark—proved extremelydifficult to isolate. Only in 1923 was it isolated,using x-ray analysis, and today it is applied pri-marily in light bulb filaments and electrodes.Sometimes it is accidentally “applied,” when zir-conium is actually the desired metal, because thetwo are so difficult to separate.

VANADIUM, TANTALUM, AND

NIOBIUM. Vanadium, named after the Norsegoddess Vanadis or Freyja, was first identified in1801, but only isolated in 1867. It is often appliedin alloys with steel, and sometimes used forbonding steel and titanium.

Tantalum and niobium are likewise namedafter figures from mythology—in this case,Greek. These two elements appear together sooften that their identification as separate sub-stances was a long process, involving numerousscientists and heated debates. For this reason,alloys involving tantalum—valued for its highmelting point—usually include niobium as well.

CHROMIUM, TUNGSTEN, AND

MANGANESE, AND OTHER MET-

ALS. Because it displays a wide variety of col-ors, chromium was named after the Greek wordfor “color.” It is used in chrome and tinted glass,and as a strengthening agent for stainless steel.Tungsten likewise strengthens steel, in large partbecause it has the highest melting point of anymetal: 6,170°F (3,410°C). Its name comes from aSwedish word meaning “heavy stone,” and itschemical symbol (W) refers to the ore namedwolframite, in which it was first discovered. Itshigh resistance to heat gives tungsten alloysapplications in areas ranging from light bulb fil-aments to furnaces to missiles.



TransitionMetals




which does not fill the 5f orbital—exhibits




ALLOY: A mixture containing more than

one metal.











lanthanides and actinides, both of which

fill the f orbitals. For this reason, they are

usually treated separately.











IUPAC SYSTEM: A version of the peri-

odic table of elements, authorized by the

International Union of Pure and Applied

Chemistry (IUPAC), which numbers all

groups on the table from 1-18. Thus in the

IUPAC system, in use primarily outside of

North America, both the representative or

main-group elements and the transition

metals are numbered.


that fill the 4f orbital. Because lan-

thanum—which does not fill the 4f








group number. Hence the transition metals

are usually not numbered in this system.

Some North American charts, however, do

provide group numbers for transition met-

als by including an “A” after the group

numbers for representative or main-group

elements, and “B” after those of transition

metals.


indicating the regions that may be occu-

pied by an electron. The higher the princi-

pal energy level, the more complex the pat-

tern of orbitals.






element. Elements are arranged in groups

or columns, as well as in rows or periods,

according to specific aspects of their

valence electron configurations.

PERIODS: Rows on the periodic table of



in the atoms of the elements involved. Ele-

ments in the transition metal family occu-

py periods 4, 5, 6, or 7.

KEY TERMS


TransitionMetals


Likewise manganese, first identified in themid-eighteenth century, is extraordinarilystrong. Its major use is in steelmaking. In addi-tion, compounds such as manganese oxide areused in fertilizers. Other alloy metals include rhe-nium (identified only in 1925) and molybde-num. These two are often applied in alloystogether, and with tungsten, for a variety ofindustrial purposes.

Rare and Artificial Elements

RARE EARTH-LIKE ELEMENTS.

Though they are not part of the lanthanide

series, scandium and yttrium share many proper-

ties with those elements, once known as the “rare

earths.” (In this context, “rarity” refers not so

much to scarcity as to difficulty of ex-

traction.) Furthermore, like most of the lan-









orbitals. Elements in the transition metal

family have principal energy levels of 4, 5,

6, or 7.









becomes stable.




transformed.

REPRESENTATIVE OR MAIN-GROUP

ELEMENTS: The 44 elements in Groups

1 through 8 on the periodic table of ele-

ments in the North American system, for

which the number of valence electrons

equals the group number. (The only excep-

tion is helium.) In the IUPAC system, these

elements are assigned group numbers 1, 2,

and 13 through 18. (In these last six

columns on the IUPAC chart, there is no

relation between the number of valence

electrons and group number.)






cipal energy level, there are n sublevels. At

each principal energy level, the first sublev-

el to be filled is the one corresponding to

the s orbital pattern, followed by the p and

then the d pattern.


elements, which are not assigned a group

number in the version of the periodic table

of elements known as the North American

system. These are the only elements that fill

the dorbital. In addition, the transition

metals—unlike the representative or main-

group elements—have their valence elec-

trons on two different principal energy lev-

els. Though the lanthanides and actinides

are considered inner transition metals,

they are usually treated separately.





KEY TERMS CONTINUED


TransitionMetals

thanides, these two elements were first discov-

ered in Scandinavia.

Their origin is reflected in the name given by

Swedish chemist Lars Nilson (1840-1899) to

scandium, which he discovered in 1876. As for

yttrium, it came from the complex mineral

known as ytterite, from whence emerged many of

the lanthanides. Yttrium is used in alloys to

impart strength, and has a wide if specialized

array of applications. Scandium has no impor-

tant uses.

ARTIFICIAL ELEMENTS. Tech-

netium, with an atomic number of 43, is unusu-

al in that it is one of the few elements with an

atomic number less than 92 that does not occur

in nature. This is reflected in its name, from the

Greek technetos, meaning “artificial.” The strange

thing is that technetium, produced in a laborato-

ry in 1936, does occur naturally in the universe—

but it seems to be found only in the spectral lines

from older stars, and not in those of younger

stars such as our Sun.

The other elements in the transition metals

family are referred to as “transuranium ele-

ments,” meaning that they have atomic numbers

higher than that of uranium (92). These have all

been created artificially; all are highly radioac-

tive; few survive for more than a few minutes (atmost); and few have applications in ordinary life.

WHERE TO LEARN MORE

The Copper Page (Web site). <http://www.copper.org>(May 26, 2001).

Gold Institute (Web site). <http://www.goldinstitute.org>(May 26, 2001).




“The Pictorial Periodic Table” (Web site). <http://chemlab.pc.maricopa.edu/periodic/periodic.html>(May 22, 2001).


“Transition Metals” ChemicalElements.com (Web site).<http://www.chemicalelements.com/groups/transition.html> (May 26, 2001).

“The Transition Metals.”University of Colorado Depart-ment of Physics (Web site). <http://www.colorado.edu/physics/2000/periodic_table/transition_elements.html> (May 26, 2001).


Zincworld (Web site). <http://www.zincworld.org> (May22, 2001).




ACTINIDESActinides

CONCEPTThe actinides (sometimes called actinoids) occu-py the “bottom line” of the periodic table—a rowof elements normally separated from the others,placed at the foot of the chart along with the lan-thanides. Both of these families exhibit unusualatomic characteristics, properties that set themapart from the normal sequence on the periodictable. But there is more that distinguishes theactinides, a group of 14 elements along with thetransition metal actinium. Only four of themoccur in nature, while the other 10 have beenproduced in laboratories. These 10 are classified,along with the nine elements to the right ofactinium on Period 7 of the periodic table, astransuranium (beyond uranium) elements. Fewof these elements have important applications indaily life; on the other hand, some of the lower-number transuranium elements do have special-ized uses. Likewise several of the naturally occur-ring actinides are used in areas ranging frommedical imaging to powering spacecraft. Thenthere is uranium, “star” of the actinide series: forcenturies it seemed virtually useless; then, in amatter of years, it became the most talked-aboutelement on Earth.

HOW IT WORKS

The Transition Metals

Why are actinides and lanthanides set apart fromthe periodic table? This can best be explained byreference to the transition metals and their char-acteristics. Actinides and lanthanides are referredto as inner transition metals, because, althoughthey belong to this larger family, they are usuallyconsidered separately—rather like grown chil-

dren who have married and started families oftheir own.

The qualities that distinguish the transitionmetals from the representative or main-groupelements on the periodic table are explained indepth within the Transition Metals essay. Thereader is encouraged to consult that essay, as wellas the one on Families of Elements, which furtherplaces the transition metals within the largercontext of the periodic table. Here these specificswill be discussed only briefly.

ORBITAL PATTERNS OF THE

TRANSITION METALS. The transitionmetals are distinguished by their configuration ofvalence electrons, or the outer-shell electronsinvolved in chemical bonding. Together with thecore electrons, which are at lower energy levels,valence electrons move in areas of probabilityreferred to as orbitals. The pattern of orbitals isdetermined by the principal energy level of theatom, which indicates a distance that an electronmay move away from the nucleus.

Each principal energy level is divided intosublevels corresponding to the number n of theprincipal energy level. The actinides, whichwould be on Period 7 if they were included onthe periodic table with the other transition met-als, have seven principal energy levels. (Note thatperiod number and principal energy level num-ber are the same.) In the seventh principal ener-gy level, there are seven possible sublevels.

The higher the energy level, the larger thenumber of possible orbital patterns, and themore complex the patterns. Orbital patternsloosely define the overall shape of the electroncloud, but this does not necessarily define thepaths along which the electrons move. Rather, it


Actinides


means that if you could take millions of photo-graphs of the electron during a period of a fewseconds, the resulting blur of images woulddescribe more or less the shape of a specifiedorbital.

The four basic types of orbital patterns arediscussed in the Transition Metals essay, and willnot be presented in any detail here. It is impor-tant only to know that, unlike the representativeelements, transition metals fill the sublevel corre-sponding to the d orbitals. In addition, they arethe only elements that have valence electrons ontwo different principal energy levels.

LANTHANIDES AND ACTI-

NIDES. The lanthanides and actinides are fur-ther set apart even from the transition metals,due to the fact that these elements also fill thehighly complex f orbitals. Thus these two familiesare listed by themselves. In most versions of theperiodic table, lanthanum (57) is followed byhafnium (72) in the transition metals section ofthe chart; similarly, actinium (89) is followed byrutherfordium (104). The “missing” metals—lanthanides and actinides, respectively—areshown at the bottom of the chart.

The lanthanides can be defined as thosemetals that fill the 4f orbital. However, becauselanthanum (which does not fill the 4f orbital)exhibits similar properties, it is usually includedwith the lanthanides. Likewise the actinides canbe defined as those metals that fill the 5f orbital;but again, because actinium exhibits similarproperties, it is usually included with theactinides.

Isotopes

One of the distinguishing factors in the actinidefamily is its great number of radioactive isotopes.Two atoms may have the same number of pro-tons, and thus be of the same element, yet differin their number of neutrons—neutrally chargedpatterns alongside the protons at the nucleus.Such atoms are called isotopes, atoms of the sameelement having different masses.

Isotopes are represented symbolically in oneof several ways. For instance, there is this format:�ma�S, where S is the chemical symbol of the ele-

ment, a is the atomic number (the number ofprotons in its nucleus), and m the mass num-ber—the sum of protons and neutrons. For the

isotope known as uranium-238, for instance, thisis shown as �

29328

�U.

Because the atomic number of any elementis established, however, isotopes are usually rep-resented simply with the mass number thus: 238U.They may also be designated with a subscriptnotation indicating the number of neutrons, sothat this information can be obtained at a glancewithout it being necessary to do the arithmetic.For the uranium isotope shown here, this is writ-ten as �

29328

�U146.

Radioactivity

The term radioactivity describes a phenomenonwhereby certain materials are subject to a form ofdecay brought about by the emission of high-energy particles, or radiation.

Types of particles emitted in radiationinclude:

• Alpha particles, or helium nuclei;

• Beta particles—either an electron or a sub-atomic particle called a positron;

• Gamma rays or other very high-energy elec-tromagnetic waves.

Isotopes are either stable or unstable, withthe unstable variety, known as radioisotopes,being subject to radioactive decay. In this con-text, “decay” does not mean “rot”; rather, aradioisotope decays by turning into another iso-tope. By continuing to emit particles, the isotopeof one element may even turn into the isotope ofanother element.

Eventually the radioisotope becomes a stableisotope, one that is not subject to radioactivedecay. This is a process that may take seconds,minutes, hours, days, years—and sometimes mil-lions or even billions of years. The rate of decayis gauged by the half-life of a radioisotope sam-ple: in other words, the amount of time it takesfor half the nuclei (plural of nucleus) in the sam-ple to become stable.

Actinides decay by a process that begins withwhat is known as K-capture, in which an electronof a radioactive atom is captured by the nucleusand taken into it. This is followed by the splitting,or fission, of the atom’s nucleus. This fission pro-duces enormous amounts of energy, as well asthe release of two or more neutrons, which mayin turn bring about further K-capture. This iscalled a chain reaction.


Actinides



The First Three NaturallyOccurring Actinides

In the discussion of the actinides that follows,atomic number and chemical symbol will followthe first mention of an element. Atomic mass fig-ures are available on any periodic table, and thesewill not be mentioned in most cases. The atomicmass figures for actinide elements are very high,as fits their high atomic number, but for most of these, figures are usually for the most stableisotope, which may exist for only a matter ofseconds.

Though it gives its name to the group as awhole, actinium (Ac, 89) is not a particularly sig-nificant element. Discovered in 1902 by Germanchemist Friedrich Otto Giesel (1852-1927), it isfound in uranium ores. Actinium is 150 timesmore radioactive than radium, a highly radioac-tive alkaline earth metal isolated around thesame time by French-Polish physicist andchemist Marie Curie (1867-1934) and her hus-band Pierre (1859-1906).

THORIUM. More significant than actini-um is thorium (Th, 90), first detected in 1815 by

the renowned Swedish chemist Jons Berzelius(1779-1848). Berzelius promptly named the ele-ment after the Norse god Thor, but eventuallyconcluded that what he had believed to be a newelement was actually the compound yttriumphosphate. In 1829, however, he examinedanother mineral and indeed found the elementhe believed he had discovered 14 years earlier.

It 1898, Marie Curie and an English chemistnamed Gerhard Schmidt, working independent-ly, announced that thorium was radioactive.Today it is believed that the enormous amountsof energy released by the radioactive decay ofsubterranean thorium and uranium plays a sig-nificant part in Earth’s high internal tempera-ture. The energy stored in the planet’s thoriumreserves may well be greater than all the energyavailable from conventional fossil and nuclearfuels combined.

Thorium appears on Earth in an abundanceof 15 parts per million (ppm), many timesgreater than the abundance of uranium. With itshigh energy levels, thorium has enormous poten-tial as a nuclear fuel. When struck by neutrons,thorium-232 converts to uranium-233, one ofthe few known fissionable isotopes—that is, iso-topes that can be split to start nuclear reactions.

GLENN T. SEABORG HOLDS A SAMPLE CONTAINING ARTIFICIAL ELEMENTS BETWEEN 94 AND 102 ON THE PERIOD-IC TABLE. (Roger Ressmeyer/Corbis. Reproduced by permission.)


ActinidesIt is perhaps ironic that this element, with itspotential for use in some of the most high-techapplications imaginable, is widely applied in avery low-tech fashion. In portable gas lanternsfor camping and other situations without electricpower, the mantle often contains oxides of thori-um and cerium, which, when heated, emit a bril-liant white light. Thorium is also used in themanufacture of high-quality glass, and as a cata-lyst in various industrial processes.

PROTACTINIUM. Russian chemistDmitri Ivanovitch Mendeleev (1834-1907),father of the periodic table, used the table’sarrangement of elements as a means of predict-ing the discovery of new substances: wherever hefound a “hole” in the table, Mendeleev could saywith assurance that a new element would even-tually be found to fill it. In 1871, Mendeleev pre-dicted the discovery of “eka-tantalum,” an ele-ment that filled the space below the transitionmetal tantalum. (At this point in history, justtwo years after Mendeleev created the periodictable, the lanthanides and actinides had not beenseparated from the rest of the elements on thechart.)

Forty years after Mendeleev foretold its exis-tence, two German chemists found what theythought might be Element 91. It had a half-life ofonly 1.175 minutes, and, for this reason, theynamed it “brevium.” Then in 1918, Austrianphysicist Lise Meitner (1878-1968)—who, alongwith Curie and French physicist MargueritePerey (1909-1975) was one of several womeninvolved in the discovery of radioactive ele-ments—was working with German chemist OttoHahn (1879-1968) when the two discoveredanother isotope of Element 91. This one had amuch, much longer half-life: 3.25 • 104 years, orabout five times as long as the entire span ofhuman civilization.

Originally named protoactinium, the nameof Element 91 was changed to protactinium (Pa),whose longest-lived isotope has an atomic massof 231. Shiny and malleable, protactinium has amelting point of 2,861.6°F (1,572°C). It is highlytoxic, and so rare that no commercial uses havebeen found for it. Indeed, protactinium couldonly be produced from the decay products ofuranium and radium, and thus it is one of thefew elements with an atomic number less than 93that cannot be said to occur in nature.

Uranium

URANIUM’S EARLY HISTORY.

Chemistry books, in fact, differ as to the numberof naturally occurring elements. Some say 88,which is the most correct figure, because protac-tinium, along with technetium (43) and two oth-ers, cannot be said to appear naturally on Earth.Other books say 92, a less accurate figure thatnonetheless reflects an indisputable fact: aboveuranium (U, 92) on the periodic table, there areno elements that generally occur in nature.(However, a few do occur as radioactive by-prod-ucts of uranium.)

But uranium is much more than just the lasttruly natural element, though for about a centu-ry it apparently had no greater importance.When German chemist Martin HeinrichKlaproth (1743-1817) discovered it in 1789, henamed it after another recent discovery: the plan-et Uranus. During the next 107 years, uraniumhad a very quiet existence, befitting a rather dull-looking material.

Though it is silvery white when freshly cut,uranium soon develops a thin coating of blackuranium oxide, which turns it a flat gray. Yetglassmakers did at least manage to find a use forit—as a coating for decorative glass, to which itimparted a hazy, fluorescent yellowish green hue.Little did they know that they were using one ofthe most potentially dangerous substances onEarth.

THE DESTRUCTIVE POWER OF

URANIUM. In 1895, German physicist Wil-helm Röntgen (1845-1923) noticed that photo-graphic plates held near a Crookes tube—adevice for analyzing electromagnetic radia-tion—became fogged. He dubbed the rays thathad caused this x rays. A year later, in 1896,French physicist Henri Becquerel (1852-1908)left some photographic plates in a drawer with asample of uranium, and discovered that the ura-nium likewise caused a fogging of the photo-graphic plates. This meant that uranium wasradioactive.

With the development of nuclear fission byHahn and German chemist Fritz Strassman in1938, uranium suddenly became all-importantbecause of its ability to undergo nuclear fission,accompanied by the release of huge amounts ofenergy. During World War II, in what was knownas the Manhattan Project, a team of scientists inLos Alamos, New Mexico, developed the atomic



Actinides

bomb. The first of the two atomic bombsdropped on Japan in 1945 contained uranium,while the second contained the transuraniumelement plutonium.

OTHER USES FOR URANIUM.

Though nuclear weapons have fortunately not

been used against human beings since 1945, ura-

nium has remained an important component of


GLOVED HANDS HOLD A GRAY LUMP OF URANIUM. THIS MATERIAL HAS BEEN REMOLDED AFTER HAVING BEEN

REMOVED FROM A TITAN II MISSILE, PART OF THE DISARMAMENT AFTER THE END OF THE COLD WAR. (Martin Mariet-

ta; Roger Ressmeyer/Corbis. Reproduced by permission.)


Actinidesnuclear energy—both in the development ofbombs and in the peaceful application of nuclearpower. It has other uses as well, due to the factthat it is extremely dense.

Indeed, uranium has a density close to thatof gold and platinum, but is much cheaper,because it is more abundant on Earth. In addi-tion, various isotopes of uranium are a by-prod-uct of nuclear power, which separates these iso-topes from the highly fissionable 235U isotope.Thus, quantities of uranium are available for usein situations where a great deal of mass isrequired in a small space: in counterweights foraircraft control systems, for instance, or as ballastfor missile reentry vehicles.

Because 238U has a very long half-life—4.47 •109 years, or approximately the age of Earth—it isused to estimate the age of rocks and other geo-logical features. Uranium-238 is the “parent” of aseries of “daughter” isotopes that geologists findin uranium ores. Uranium-235 also produces“daughter” isotopes, including isotopes of radi-um, radon, and other radioactive series. Eventu-ally, uranium isotopes turn into lead, but this cantake a very long time: even 235U, which lasts for amuch shorter period than 238U, has a half-life ofabout 700 million years.

The radiation associated with various iso-topes of uranium, as well as other radioactivematerials, is extremely harmful. It can cause allmanner of diseases and birth defects, and ispotentially fatal. The tiny amounts of radiationproduced by uranium in old pieces of decorativeglass is probably not enough to cause any realharm, but the radioactive fallout from Hiroshimaresulted in birth defects among the Japanesepopulation during the late 1940s.

Transuranium Elements

Transuranium elements are those elements withatomic numbers higher than that of uranium.None of these occur in nature, except as isotopesthat develop in trace amounts in uranium ore.The first such element was neptunium (Np, 93),created in 1940 by American physicist EdwinMattison McMillan (1907-1991) and Americanphysical chemist Philip Hauge Abelson.

The development of the cyclotron by Amer-ican physicist Ernest Lawrence (1901-1958) atthe University of California at Berkeley in the1930s made possible the artificial creation of newelements. A cyclotron speeds up protons or ions

(charged atoms) and shoots them at atoms ofuranium or other elements with the aim ofadding positive charges to the nucleus. In the firsttwo decades after the use of the cyclotron to cre-ate neptunium, scientists were able to developeight more elements, all the way up to mendele-vium (Md, 101), named in honor of the manwho created the periodic table.

Most of these efforts occurred at Berkeleyunder the leadership of American nuclearchemist Glenn T. Seaborg (1912-1999) andAmerican physicist Albert Ghiorso. The pace ofdevelopment in transuranium elements slowedafter about 1955, however, primarily due to theneed for ever more powerful cyclotrons and ion-accelerating machines. In addition to Berkeley,there are two other centers for studying thesehigh-energy elements: the Joint Institute forNuclear Research in Dubna, Russia, and theGesellschaft für Schwerionenforschung (GSI) inDarmstadt, Germany.

The Transuranium Actinides

PLUTONIUM. Just as neptunium hadbeen named for the next planet beyond Uranus,a second transuranium element, discovered bySeaborg and two colleagues in 1940, was namedafter Pluto. Among the isotopes of plutonium(Pu, 94) is plutonium-239, one of the few fis-sionable isotopes other than uranium-233 anduranium-235. For that reason, it was applied inthe second bomb dropped on Japan.

In addition to its application in nuclearweapons, plutonium is used in nuclear powerreactors, and in thermoelectric generators, whichconvert the heat energy it releases into electricity.Plutonium is also used as a power source in arti-ficial heart pacemakers. Huge amounts of the ele-ment are produced each year as a by-product ofnuclear power reactors.

BERKELEY DISCOVERIES OF

THE 1950S. Because the lanthanide ele-ment above it on the periodic table was namedeuropium after Europe, americium (Am, 95) wasnamed after America. Discovered by Seaborg,Ghiorso, and two others in 1944, it was first pro-duced in a nuclear reaction involving pluto-nium-239. Americium radiation is used in meas-uring the thickness of glass during production;in addition, the isotope americium-241 is used as an ionization source in smoke detectors, and



Actinides


ACTINIDES: Those elements that fill the

5f orbital. Because actinium—which does

not fill the 5f orbital—exhibits characteris-

tics similar to those of the actinides, it is

usually considered part of the actinides

family. Of the other 14 actinides, usually

shown at the bottom of the periodic table,

only the first three occur in nature.










amount. For a sample of radioisotopes, the

half-life is the amount of time it takes for

half of the nuclei to become stable iso-

topes. Half-life can be a few seconds; on the

other hand, for uranium-238, it is a matter

of several billion years.


lanthanides and actinides, which are

unique in that they fill the f orbitals.

For this reason, they are usually treated

separately.






unstable. The unstable type, known as



that fill the 4f orbital.


and neutrons in the atom’s nucleus. The

designation 238U, for uranium-238, means

that this particular isotope of uranium has

a mass number of 238. Since uranium has

an atomic number of 92, this means that

uranium-238 has 146 neutrons in its

nucleus.




protons.





ORBITAL: A region of probabilities





KEY TERMS

in portable devices for taking gamma-ray

photographs.

Above curium (Cm, 96) on the periodic

table is the lanthanide gadolinium, named after

Finnish chemist Johan Gadolin (1760-1852).

Therefore, the discoverers of Element 96 also

decided to name it after a person, Marie Curie. As

with some of the other relatively low-number

transuranium elements, this one is not entirely

artificial: its most stable isotope, some geologists

believe, may have been present in rocks many

millions of years ago, but these isotopes have

long since decayed. Because curium generates

great amounts of energy as it decays, it is used for

providing compact sources of power in remote

locations on Earth and in space vehicles.


Actinides










orbitals. Elements in the transition metal

family have principal energy levels of 4, 5,

6, or 7.








beings.


phenomenon whereby certain isotopes,

known as radioisotopes, are subject to a






becomes stable.

RADIOISOTOPE: An isotope subject to

the decay associated with radioactivity. A

radioisotope is thus an unstable isotope.




cipal energy level, there are n sublevels.

Actinides are distinguished by the fact that

their valence electrons are in a sublevel

corresponding to the 5f orbital.


elements (counting lanthanum and actini-

um), which are the only elements that fill

the d orbital. In addition, the transition

metals have their valence electrons on two

different principal energy levels. Though

the lanthanides and actinides are consid-

ered inner transition metals, they are usu-

ally considered separately.

TRANSURANIUM ELEMENTS: Ele-

ments with an atomic number higher than

that of uranium (92). These have all have

been produced artificially. The transurani-

um elements include 11 actinides, as well

as 9 transition metals.





KEY TERMS CONTINUED

When Seaborg, Ghiorso, and others createdElement 97, berkelium (Bk), they again took thenaming of lanthanides as their cue. Just as ter-bium, directly above it on the periodic table, hadbeen named for the Swedish town of Ytterby,where so many lanthanides were discovered, theynamed the new element after the American citywhere so many transuranium elements had beendeveloped. Berkelium has no known applications

outside of research. The Berkeley team likewise

named californium (Cf, 98) after the state where

Berkeley is located. Researchers today are study-

ing the use of californium radiation for treat-

ment of tumors involved in various forms of

cancer.

THE REMAINING TRANSURANI-

UM ACTINIDES. The remaining transura-


Actinides nium actinides were all named after famous peo-ple: einsteinium (Es, 99) for Albert Einstein(1879-1955); fermium (Fm, 100) after Italian-American physicist Enrico Fermi (1901-1954);mendelevium after Mendeleev; nobelium (No,102) after Swedish inventor and philanthropistAlfred Nobel (1833-1896); and lawrencium (Lr,103) after Ernest Lawrence.

Both einsteinium and fermium were by-products of nuclear testing at Bikini Atoll in thesouth Pacific in 1952. For this reason, their exis-tence was kept a secret for two years. Neither ele-ment has a known application. The same is trueof mendelevium, produced by Seaborg, Ghiorso,and others with a cyclotron at Berkeley in 1955,as well as the other two transuranium actinides.

Beyond the TransuraniumActinides

As noted earlier, there are nine additionaltransuranium elements, which properly belongto the transition metals. The first of these isrutherfordium, discovered in 1964 by the Dubnateam and named after Ernest Rutherford (1871-1937), the British physicist who discovered thenucleus. The Dubna team named dubnium, dis-covered in 1967, after their city, just as berkeliumhad been named after the Berkeley team’s city.Both groups developed versions of Element 106in 1974, and both agreed to name it seaborgiumafter Seaborg, but this resulted in a controversythat was not settled for some time.

The name of bohrium, created at Dubna in1976, honors Danish physicist Niels Bohr (1885-

1962), who developed much of the model ofelectron energy levels discussed earlier in thisessay. Hassium, produced at the GSI in 1984, isnamed for the German state of Hess. Two yearsearlier, the GSI team also created the last namedelement on the periodic table, meitnerium (109),named after Meitner. Beyond meitnerium arethree elements, as yet unnamed, created at theGSI in the mid-1990s.

WHERE TO LEARN MORE

Cooper, Dan. Enrico Fermi and the Revolutions in Mod-ern Physics. New York: Oxford University Press, 1999.

“Exploring the Table of Isotopes” (Web site). <http://ie.lbl.gov/education/isotopes.htm> (May 15, 2001).

Kidd, J. S. and Renee A. Kidd. Quarks and Sparks: TheStory of Nuclear Power. New York: Facts on File, 1999.


“A Periodic Table of the Elements” Los Alamos NationalLaboratory (Web site). <http://pearl1.lanl.gov/period-ic/ (May 22, 2001).

“The Pictorial Periodic Table” (Web site). <http://chem-lab.pc.maricopa.edu/periodic/periodic.html> (May22, 2001).

Sherrow, Victoria. The Making of the Atom Bomb. SanDiego, CA: Lucent Books, 2000.

“Some Physics of Uranium.” The Uranium Institute (Website). <http://www.uilondon.org/education_resources/physics_of_uranium/in dex1.htm> (May27, 2001).


Uranium Information Center (Web site). <http://www.uic.com.au/> (May 27, 2001).




LANTHANIDESLanthanides

CONCEPTAlong the bottom of the periodic table of ele-ments, separated from the main body of thechart, are two rows, the first of which representsthe lanthanides. Composed of lanthanum andthe 14 elements of the lanthanide series, the lan-thanides were once called the “rare earth” metals.In fact, they are not particularly rare: many ofthem appear in as much abundance as morefamiliar elements such as mercury. They are,however, difficult to extract, a characteristic thatdefines them as much as their silvery color;sometimes high levels of reactivity; and sensitiv-ity to contamination. Though some lanthanideshave limited uses, members of this group arefound in everything from cigarette lighters to TVscreens, and from colored glass to control rods innuclear reactors.

HOW IT WORKS

Defining the Lanthanides

The lanthanide series consists of the 14 elements,with atomic numbers 58 through 71, that followlanthanum on the periodic table of elements.These 14, along with the actinides—atomicnumbers 90 through 103—are set aside from theperiodic table due to similarities in propertiesthat define each group.

Specifically, the lanthanides and actinidesare the only elements that fill the f-orbitals. Thelanthanides and actinides are actually “branches”of the larger family known as transition metals.The latter appear in groups 3 through 12 on theIUPAC version of the periodic table, though

they are not numbered on the North Americanversion.

The lanthanide series is usually combinedwith lanthanum, which has an atomic number of57, under the general heading of lanthanides. Astheir name indicates, members of the lanthanideseries share certain characteristics with lan-thanum; hence the collective term “lanthanides.”These 15 elements, along with their chemicalsymbols, are:

• Lanthanum (La)

• Cerium (Ce)

• Praseodymium (Pr)

• Neodymium (Nd)

• Promethium (Pm)

• Samarium (Sm)

• Europium (Eu)

• Gadolinium (Gd)

• Terbium (Tb)

• Dysprosium (Dy)

• Holmium (Ho)

• Erbium (Er)

• Thulium (Tm)

• Ytterbium (Yb)

• Lutetium (Lu)

Most of these are discussed individually inthis essay.

PROPERTIES OF LANTHA-

NIDES. Bright and silvery in appearance,many of the lanthanides—though they are met-als—are so soft they can be cut with a knife. Lan-thanum, cerium, praseodymium, neodymium,and europium are highly reactive. When exposedto oxygen, they form an oxide coating. (An oxideis a compound formed by metal with an oxygen.)To prevent this result, which tarnishes the


Lanthanides


metal, these five lanthanides are kept stored inmineral oil.

The reactive tendencies of the other lan-thanides vary: for instance, gadolinium andlutetium do not oxidize until they have beenexposed to air for a very long time. Nonetheless,lanthanides tend to be rather “temperamental” asa class. If contaminated with other metals, suchas calcium, they corrode easily, and if contami-nated with nonmetals, such as nitrogen or oxy-gen, they become brittle. Contamination alsoalters their boiling points, which range from1,506.2°F (819°C) for ytterbium to 3,025.4°F(1,663°C) for lutetium.

Lanthanides react rapidly with hot water, ormore slowly with cold water, to form hydrogengas. As noted earlier, they also are quite reactivewith oxygen, and they experience combustionreadily in air. When a lanthanide reacts withanother element to form a compound, it usuallyloses three of its outer electrons to form what arecalled tripositive ions, or atoms with an electriccharge of +3. This is the most stable ion for lan-thanides, which sometimes develop less stable +2or +4 ions. Lanthanides tend to form ionic com-pounds, or compounds containing either posi-tive or negative ions, with other substances—inparticular, fluorine.

Are They Really “Rare”?

Though they were once known as the rare earthmetals, lanthanides were so termed because, aswe shall see, they are difficult to extract fromcompounds containing other substances—including other lanthanides. As for rarity, thescarcest of the lanthanides, thulium, is moreabundant than either arsenic or mercury, andcertainly no one thinks of those as rare sub-stances. In terms of parts per million (ppm),thulium has a presence in Earth’s crust equivalentto 0.2 ppm. The most plentiful of the lan-thanides, cerium, has an abundance of 46 ppm,greater than that of tin.

If, on the other hand, rarity is understoodnot in terms of scarcity, but with regard to diffi-culty in obtaining an element in its pure form,then indeed the lanthanides are rare. Becausetheir properties are so similar, and because theyare inclined to congregate in the same sub-stances, the original isolation and identificationof the lanthanides was an arduous task that tookwell over a century. The progress followed a com-mon pattern.

First, a chemist identified a new lanthanide;then a few years later, another scientist camealong and extracted another lanthanide from the

MISCH METAL, WHICH INCLUDES CERIUM, IS OFTEN USED IN CIGARETTE LIGHTERS. (Massimo Listri/Corbis. Reproduced by

permission.)


Lanthanidessample that the first chemist had believed to be asingle element. In this way, the lanthanidesemerged over time, each from the one before it,rather like Russian matryoshka or “nesting” dolls.

EXTRACTING LANTHANIDES.

Though most of the lanthanides were first isolat-ed in Scandinavia, today they are found in con-siderably warmer latitudes: Brazil, India, Aus-tralia, South Africa, and the United States. Theprincipal source of lanthanides is monazite, aheavy, dark sand from which about 50% of thelanthanide mass available to science and industryhas been extracted.

In order to separate lanthanides from otherelements, they are actually combined with othersubstances—substances having a low solubility,or tendency to dissolve. Oxalates and fluoridesare low-solubility substances favored for thispurpose. Once they are separated from non-lan-thanide elements, ion exchange is used to sepa-rate one lanthanide element from another.

There is a pronounced decrease in the radiiof lanthanide atoms as they increase in atomicnumber: in other words, the higher the atomicnumber, the smaller the radius. This decrease,known as the lanthanide contraction, aids in theprocess of separation by ion exchange. The lan-thanides are mixed in an ionic solution, thenpassed down a long column containing a resin.Various lanthanide ions bond more or less tight-ly, depending on their relative size, with the resin.

After this step, the lanthanides are washedout of the ion exchange column and into varioussolutions. One by one, they become fully separat-ed, and are then mixed with acid and heated toform an oxide. The oxide is then converted to afluoride or chloride, which can then be reducedto metallic form with the aid of calcium.


The Historical Approach

In studying the lanthanides, one can simplymove along the periodic table, from lanthanumall the way to lutetium. However, in light of thedifficulties involved in extracting the lan-thanides, one from another, an approach alonghistorical lines aids in understanding the uniqueplace each lanthanide occupies in the overallfamily.


The terms “lanthanide series” or even “lan-

thanides” did not emerge for some time—in

other words, scientists did not immediately know

that they were dealing with a whole group of

metals. As is often the case with scientific discov-

ery, the isolation of lanthanides followed an

irregular pattern, and they did not emerge in


Cerium was in fact discovered long before

lanthanum itself, in the latter half of the eigh-

teenth century. There followed, a few decades

later, the discovery of a mineral called ytterite,

named after the town of Ytterby, Sweden, near

which it was found in 1787. During the next cen-

tury, most of the remaining lanthanides were

extracted from ytterite, and the man most

responsible for this was Swedish chemist Carl

Gustav Mosander (1797-1858).

Because Mosander had more to do with the

identification of the lanthanides than any one

individual, the middle portion of this historical

overview is devoted to his findings. The recogni-

tion and isolation of lanthanides did not stop

with Mosander, however; therefore another

SAMARIUM IS USED IN NUCLEAR POWER PLANT CON-TROL RODS, SUCH AS THE ONE SHOWN HERE. (Tim

Wright/Corbis. Reproduced by permission.)


Lanthanides group of minerals is discussed in the context ofthe latter period of lanthanide discovery.

Early Lanthanides

CERIUM. In 1751, Swedish chemist AxelCrönstedt (1722-1765) described what hethought was a new form of tungsten, which hehad found at the Bastnäs Mine near Riddarhyt-tan, Sweden. Later, German chemist MartinHeinrich Klaproth (1743-1817) and Swedishchemist Wilhelm Hisinger (1766-1852) inde-pendently analyzed the material Crönstedt haddiscovered, and both concluded that this must bea new element. It was named cerium in honor ofCeres, an asteroid between Mars and Jupiter dis-covered in 1801. Not until 1875 was cerium actu-ally extracted from an ore.

Among the applications for cerium is analloy called misch metal, prepared by fusing thechlorides of cerium, lanthanum, neodymium,and praseodymium. The resulting alloy ignites ator below room temperature, and is often used asthe “flint” in a cigarette lighter, because it sparkswhen friction from a metal wheel is applied.

Cerium is also used in jet engine parts, as acatalyst in making ammonia, and as an anti-knock agent in gasoline—that is, a chemical thatreduces the “knocking” sounds sometimes pro-duced in an engine by inferior grades of fuel. Incerium (IV) oxide, or CeO2, it is used to extractthe color from formerly colored glass, and is alsoapplied in enamel and ceramic coatings.

GADOLINIUM. In 1794, seven yearsafter the discovery of ytterite, Finnish chemistJohan Gadolin (1760-1852) concluded thatytterite contained a new element, which was laternamed gadolinite in his honor. A very similarname would be applied to an element extractedfrom ytterite, and the years between Gadolin’sdiscovery and the identification of this elementspanned the period of the most fruitful activityin lanthanide identification.

During the next century, all the other lan-thanides were discovered within the compositionof gadolinite; then, in 1880, Swiss chemist Jean-Charles Galissard de Marignac (1817-1894)found yet another element hiding in it. Frenchchemist Paul Emile Lecoq de Boisbaudran (1838-1912) rediscovered the same element six yearslater, and proposed that it be called gadolinium.

Silvery in color, but with a sometimes yel-lowish cast, gadolinium has a high tendency tooxidize in dry air. Because it is highly efficient forcapturing neutrons, it could be useful in nuclearpower reactors. However, two of its seven iso-topes are in such low abundance that it has hadlittle nuclear application. Used in phosphors forcolor television sets, among other things,gadolinium shows some promise for ultra high-tech applications: at very low temperatures itbecomes highly magnetic, and may function as asuperconductor.

Mosander’s Lanthanides

LANTHANUM. Between 1839 and 1848,Mosander was consumed with extracting variouslanthanides from ytterite, which by then hadcome to be known as gadolinite. When he firstsucceeded in extracting an element, he named itlanthana, meaning “hidden.” The material, even-tually referred to as lanthanum, was not preparedin pure form until 1923.

Like a number of other lanthanides, lan-thanum is very soft—so soft it can be cut with aknife—and silvery-white in color. Among themost reactive of the lanthanides, it decomposesrapidly in hot water, but more slowly in coldwater. Lanthanum also reacts readily with oxy-gen, and corrodes quickly in moist air.

As with cerium, lanthanum is used in mischmetal. Because lanthanum compounds bringabout special optical qualities in glass, it alsoused for the manufacture of specialized lenses. Inaddition, compounds of lanthanum with fluo-rine or oxygen are used in making carbon-arclamps for the motion picture industry.

SAMARIUM. While analyzing an oxideformed from lanthanide in 1841, Mosanderdecided that he had a new element on his hands,which he called didymium. Four decades later,Boisbaudran took another look at didynium, andconcluded that it was not an element; rather, itcontained an element, which he named samari-um after the mineral samarskite, in which it isfound. Still later, Marignac was studyingsamarskite when he discovered what came to beknown as gadolinium. But the story did not endthere: even later, in 1901, French chemist Eugène-Anatole Demarçay (1852-1903) found yet anoth-er element, europium, in samarskite.

Samarium is applied today in nuclear powerplant control rods, in carbon-arc lamps, and in



Lanthanidesoptical masers and lasers. In alloys with cobalt, itis used in manufacturing the most permanentelectromagnets available. Samarium is also uti-lized in the manufacture of optical glass, and as acatalyst in the production of ethyl alcohol.

ERBIUM AND TERBIUM. To returnto Mosander, he was examining ytterite in 1843when he identified three different “earths,” all ofwhich he also named after Ytterby: yttria, erbia,and terbia. Erbium was the first to be extracted.A pure sample of its oxide was prepared in 1905by French chemist Georges Urbain (1872-1938)and American chemist Charles James (1880-1928), but the pure metal itself was only extract-ed in 1934.

Soft and malleable, with a lustrous silverycolor, erbium produces salts (which are usuallycombinations of a metal with a nonmetal) thatare pink and rose, making it useful as a tintingagent. One of its oxides is utilized, for instance, totint glass and porcelain with a pinkish cast. It isalso applied, to a limited extent, in the nuclearpower industry.

Mosander also identified another element,terbium, in ytterite in 1839, and Marignac isolat-ed it in a purer form nearly half a century later, in1886. To repeat a common theme, it is silvery-gray and soft enough to be cut with a knife.When hit by an electron beam, a compound con-taining terbium emits a greenish color, and thusit is used as a phosphor in color television sets.

Later Isolation of Lanthanides

YTTERBIUM, HOLMIUM, AND

THULIUM. For many years after Mosander,there was little progress in the discovery of lan-thanides, and when it came, it was in the form ofa third element, named after the town where somany of the lanthanides were discovered. In1878, while analyzing what Mosander had callederbia, Marignac realized that it contained one orpossibly two elements.

A year later, Swedish chemist Lars FrederikNilson (1840-1899) concluded that it did indeedcontain two elements, which were named ytter-bium and scandium. (Scandium, with an atomicnumber of 21, is not part of the lanthanideseries.) Urbain is sometimes credited for discov-ering ytterbium: in 1907, he showed that thematerials Nilson had studied were actually a mix-ture of two oxides. In any case, Urbain said that


ALLOY: A mixture of two or more

metals.








a net electrical charge.

LANTHANIDE CONTRACTION: A

progressive decrease in the radius of lan-

thanide atoms as they increase in atomic

number.

LANTHANIDE SERIES: A group of 14

elements, with atomic numbers 58 through

71, that follow lanthanum on the periodic

table of elements.

LANTHANIDES: The lanthanide se-

ries, along with lanthanum.

OXIDE: A compound formed by the

chemical bonding of a metal with oxygen.


chart showing the elements arranged in

order of atomic number, grouping them

according to common characteristics.

RARE EARTH METALS: An old name

for the lanthanides, reflecting the difficulty

of separating them from compounds

containing other lanthanides or other

substances.

TRANSITION METALS: Groups 3

through 12 on the IUPAC or European ver-

sion of the periodic table of elements. The

lanthanides and actinides, which appear at

the bottom of the periodic table, are

“branches” of this family.

KEY TERMS


Lanthanides the credit should be given to Marignac, who isthe most important figure in the history of lan-thanides other than Mosander. As for ytterbium,it is highly malleable, like other lanthanides, butdoes not have any significant applications inindustry.

Swedish chemist Per Teodor Cleve (1840-1905) found in 1879 that erbia contained twomore elements, which he named holmium and thulium. Thulium refers to the ancient namefor Scandinavia, Thule. Rarest of all the lan-thanides, thulium is highly malleable—and alsohighly expensive. Hence it has few commercialapplications.

DYSPROSIUM. Named for the Greekword dysprositos, or “hard to get at,” dysprosiumwas discovered by Boisbaudran. Separatingytterite in 1886, he found gallium (atomic num-ber 31—not a lanthanide); samarium (discussedabove); and dysprosium. Yet again, a mineralextracted from ytterite had been named after apreviously discovered element, and, yet again, itturned out to contain several elements. The sub-stance in question this time was holmium, which,as Boisbaudran discovered, was actually a com-plex mixture of terbium, erbium, holmium, andthe element he had identified as dysprosium. Apure sample was not obtained until 1950.

Because dysprosium has a high affinity forneutrons, it is sometimes used in control rods fornuclear reactors, “soaking up” neutrons rather asa sponge soaks up water. Soft, with a lustrous silver color like other lanthanides, dysprosium is also applied in lasers, but otherwise it has few uses.

EUROPIUM AND LUTETIUM.

Whereas many other lanthanides are named forregions in northern Europe, the name for europi-um refers to the European continent as a whole,and that of lutetium is a reference to the oldRoman name for Paris. As mentioned earlier,

Demarçay found europium in samarskite, a dis-covery he made in 1901. Actually, Boisbaudranhad noticed what appeared to be a new elementabout a decade previously, but he did not pursueit, and thus the credit goes to his countryman.

Most reactive of the lanthanides, europiumresponds both to cold water and to air. In addi-tion, it is capable of catching fire spontaneously.Among the most efficient elements for the capture of neutrons, it is applied in the controlsystems of nuclear reactors. In addition, its com-pounds are utilized in the manufacture of phos-phors for TV sets: one such compound, forinstance, emits a reddish glow. Yet anothereuropium compound is added to the glue onpostage stamps, making possible the electronicscanning of stamps.

Urbain, who discovered lutetium, named itafter his hometown. James also identified a formof the lanthanide, but did not announce his dis-covery until much later. Except for some uses at acatalyst in the production of petroleum, lutetiumhas few industrial applications.

WHERE TO LEARN MORE

Cotton, Simon. Lanthanides and Actinides. New York:Oxford University Press, 1991.

Heiserman, David L. Exploring Chemical Elements andTheir Compounds. Blue Ridge Summit, PA: TabBooks, 1992.

“Luminescent Lanthanides” (Web site). <http://orgwww.chem.uva.nl/lanthanides/> (May 16, 2001).


Oxlade, Chris. Metal. Chicago: Heinemann Library,2001.


Whyman, Kathryn. Metals and Alloys. Illustrated byLouise Nevett and Simon Bishop. New York: Glouces-ter Press, 1988.



211

SC IENCE OF EVERYDAY TH INGSr e a l - l i f e CHEMISTRY

NONMETALS ANDMETALLO IDS

NONMETALS ANDMETALLO IDS

NONMETALS

METALLO IDS

HALOGENS

NOBLE GASES

CARBON

HYDROGEN



NONMETALSNonmetals

CONCEPTNonmetals, as their name implies, are elementsthat display properties quite different from thoseof metals. Generally, they are poor conductors ofheat and electricity, and they are not ductile: inother words, they cannot be easily reshaped.Included in this broad grouping are the six noblegases, the five halogens, and eight “orphan” ele-ments. Two of these eight—hydrogen and car-bon—are so important that separate essays aredevoted to them. Two more, addressed in thisessay, are absolutely essential to human life: oxy-gen and nitrogen. Hydrogen and helium, a non-metal of the noble gas family, together accountfor about 99% of the mass of the universe, whileEarth and the human body are composed prima-rily of oxygen, with important components ofcarbon, nitrogen, and hydrogen. Indeed, much oflife—human, animal, and plant—can besummed up with these four elements, whichtogether make Earth different from any otherknown planet. Among the other “orphan”nonmetals are phosphorus and sulfur, the “brimstone” of the Bible, as well as boron andselenium.

HOW IT WORKS

Defining the Nonmetals

The majority of elements on the periodic tableare metals: solids (along with one liquid, mercu-ry) which are lustrous or shiny in appearance.Metals are ductile, or malleable, meaning thatthey can be molded into different shapes withoutbreaking. They are excellent conductors of heatand electricity, and tend to form positive ions by

losing electrons. The vast majority of elements—

87 in all—are metals, and these occupy the left,

center, and part of the right-hand side of the

periodic table.

With the exception of hydrogen, placed in

Group 1 above the alkali metals, the nonmetals

fill a triangle-shaped space in the upper right

corner of the periodic table. All gaseous elements

are nonmetals, a group that also includes some

solids, as well as bromine, the only element other

than mercury that is liquid at room temperature.

In general, the nonmetals are characterized by

properties opposite those of metals. They are

poor conductors of heat and electricity (though

carbon, a good electrical conductor, is an excep-

tion), and they tend to form negative ions by

gaining electrons. They are not particularly mal-

leable, and whereas most metals are shiny, non-

metals may be dull-colored, black (in the case of

carbon in some forms or allotropes), or invisible,

as is the case with most of the gaseous nonmetals.

Grouping the Nonmetals

Whereas the metals can be broken down into five

families, along with seven “orphans”—elements

that do not fit into a family grouping—the non-

metals are arranged in two families, as well as

eight “orphans.” Hydrogen, because of its great

abundance in the universe and its importance to

chemical studies, is addressed in a separate essay

within this book. The same is true of carbon:

though not nearly as abundant as hydrogen, car-

bon is a common element of all living things, and

therefore it is discussed in the essay on Carbons

and Organic Chemistry.

set_vol1_sec6 9/13/01 11:50 AM

Nonmetals


The other six “orphan” nonmetals, whichwill be examined in detail later in this essay, arelisted below by atomic number:

• 5. Boron (B)

• 7. Nitrogen (N)

• 8. Oxygen (O)

• 15. Phosphorus (P)

• 16. Sulfur (S)

• 34. Selenium (Se)

THE NOBLE GASES. The noblegases, discussed in detail within an essay de-voted to that subject, are listed below by atomicnumber:

• 2. Helium (He)

• 10. Neon (Ne)

• 18. Argon (Ar)

• 36. Krypton (Kr)

• 54. Xenon (Xe)

• 86. Radon (Rn)

Occupying Group 8 of the North Americanperiodic table, the noble gases—with the excep-tion of helium—have valence electron configura-tions of ns2np6. This means that they have twovalence electrons (the electrons involved inchemical bonding) on the orbital designated as s,and six more on the p orbital. As for n, this des-

ignates the energy level, a number that corre-sponds to the period or row that an elementoccupies on the periodic table.

Most elements bond according to what isknown as the octet rule, forming a shell com-posed of eight electrons. Since the noble gasesalready have eight electrons on their shell, theytend not to bond with other elements: hence thename “noble,” which in this context means “setapart.” Helium, on the other hand, has a valenceshell of s2; however, it too is characterized by alack of reactivity, and therefore is included in thenoble gas family. Noble gases are all monatomic,meaning that they exist as individual atoms,rather than in molecules. To put it another way,their molecules are single atoms. (By contrast,atoms of oxygen, for instance, usually combine toform a diatomic molecule, designated O2.)

Due to their apparent lack of reactivity, thenoble gases—also known as the rare gases—wereonce called the inert (inactive) gases. Indeed,helium, neon, and argon have not been found tocombine with other elements to form com-pounds. However, in 1962, English chemist NeilBartlett (1932-) succeeded in preparing a com-pound of xenon with platinum and fluorine(XePtF6), thus overturning the idea that the

LIQUID NITROGEN ACCOUNTS FOR ABOUT ONE-THIRD OF ALL COMMERCIAL USES OF NITROGEN. HERE, A MEDICAL

WORKER PUTS A SAMPLE OF BONE MARROW INTO A TANK OF LIQUID NITROGEN FOR PRESERVATION. (Leif Skoogfors/

Corbis. Reproduced by permission.)

set_vol1_sec6 9/13/01 11:50 AM

Nonmetals


noble gases were entirely “inert.” Since that time,numerous compounds of xenon with other ele-ments, most notably oxygen and fluorine, havebeen developed. Fluorine has also been used toform simple compounds with krypton andradon.

THE HALOGENS. Next to the noblegases, in Group 7, are the halogens, a family dis-cussed in a separate essay. These are listed below,along with atomic number and chemical symbol.Note that astatine is sometimes included with themetalloids, elements that display characteristicsboth of metals and nonmetals.

• 9. Fluorine (F)

• 17. Chlorine (Cl)

• 35. Bromine (Br)

• 53. Iodine (I)

• 85. Astatine (At)

The halogens all have valence electron con-figurations of ns2np5: in other words, at anyenergy level n, they have two valence electrons inthe s orbital, and 5 in the p orbital. In terms ofphase of matter, the halogens are the most variedfamily on the periodic table. Fluorine and chlo-rine are gases, iodine is a solid, and bromine (asnoted earlier) is one of only two elements exist-ing at room temperature as a liquid. As for asta-

tine, it is a solid too, but so highly radioactive it ishard to determine much about its properties.(The only other nonmetal that has a large num-ber of radioactive isotopes is the noble gas radon,considered highly dangerous.)

Though they are “next door” to the noblegases, the halogens could not be more different.Whereas their neighbors to the right are the leastreactive elements on the periodic table, the halo-gens are the most reactive. Indeed, fluorine hasthe highest possible value of electronegativity,the relative ability of an atom to attract valenceelectrons. It is therefore one of the only elementsthat will bond with a noble gas.

All of the halogens tend to form salts, com-pounds—formed, along with water, from thereaction of an acid and base—that bring togeth-er a metal and a nonmetal. Due to this tendency,the first of the family to be isolated—chlorine, in1811—was originally named “halogen,” a combi-nation of the Greek words halos, or salt, and gen-nan, “to form or generate.” In their pure form,halogens are diatomic, and in contact with otherelements, they form ionic bonds, which are thestrongest form of chemical bond. In the processof bonding, they become negatively charged ions,or anions.

AT AN “OXYGEN BAR,” YOU CAN RELAX BY BREATHING OXYGEN. HERE, A WOMAN IS BREATHING FLAVORED OXYGEN

AT A BAR IN JAPAN. (AFP/Corbis. Reproduced by permission.)

set_vol1_sec6 9/13/01 11:50 AM

NonmetalsAbundance of the Nonmetals

IN THE UNIVERSE. As is statedmany times throughout this book, humans maybe created equal, but the elements are not.Though 88 elements exist in nature, this certain-ly does not mean that each occupies a 1/88 shareof the total pie. Just two elements—the non-metals hydrogen and helium, which occupy thefirst two positions on the periodic table—account for 99.9% of the matter in the entire uni-verse, yet the percentage of matter they make upon Earth is very small.

The reason for this disparity is that, whereasour own planet (as far as we know) is unique insustaining life—requiring oxygen, carbon, nitro-gen, and other elements—the vast majority ofthe universe is made up of stars, composed pri-marily of hydrogen and helium.

ON EARTH AND IN THE AT-

MOSPHERE. Nonmetals account for a largeportion of Earth’s total known elemental mass—that is, the composition of the planet’s crust,waters, and atmosphere. Listed below are figuresranking nonmetals within the overall picture ofelements on Earth. The numbers following eachelement’s name indicate the percentage of eachwithin the planet’s total known elemental mass.

• 1. Oxygen (49.2%)

• 9. Hydrogen (0.87%)

• 11. Chlorine (0.19%)

• 12. Phosphorus (0.11%)

• 14. Carbon (0.08%)

• 15. Sulfur (0.06%)

• 17. Nitrogen (0.03%)

• 18. Fluorine (0.03%)

Some of the figures above may seem rathersmall, but in fact only 18 elements account for allbut 0.49% of the planet’s known elemental mass,the remainder being composed of numerousother elements in small quantities. In Earth’satmosphere, the composition is all nonmetallic,as one might expect:

• 1. Nitrogen (78%)

• 2. Oxygen (21%)

• 3. Argon (0.93%)

• 4. Various trace gases, including watervapor, carbon dioxide, and ozone or O3

(0.07%)

IN THE HUMAN BODY. As notedearlier, carbon is a component in all living things,and it constitutes the second-most abundant ele-

ment in the human body. Together with oxygenand hydrogen, it accounts for 93% of the body’smass. Listed below are the most abundant non-metals in the human body, by ranking.

• 1. Oxygen (65%)

• 2. Carbon (18%)

• 3. Hydrogen (10%)

• 4. Nitrogen (3%)

• 6. Phosphorus (1.0%)

• 9. Sulfur (0.26%)

• 11. Chlorine (0.14%)


Boron

The first “orphan” nonmetal, by atomic number,is boron, named after the Arabic word buraq orthe Persian burah. (It is thus unusual in beingone of the only elements whose name is notbased on a word from a European language, or—for the later elements—the name of a person orplace.) It was discovered in 1808 by Englishchemist Sir Humphry Davy (1778-1829), a manresponsible for identifying or isolating numerouselements; French physicist and chemist JosephGay-Lussac (1778-1850), known in part for thegas law named after him; and French chemistLouis Jacques Thénard (1777-1857). These scien-tists used the reaction of boric acid (H3BO3) withpotassium to isolate the element.

Sometimes classified as a metalloid becauseit is a semiconductor of electricity, boron isapplied in filaments used in fiber optics research.Few of its other applications, however, havemuch to do with electrical conductivity. Boric orboracic acid is used as a mild antiseptic, and inNorth America, it is applied for the control ofcockroaches, silverfish, and other pests. A com-pound known as borax is a water softener inwashing powders, while other compounds areused to produce enamels for the coating ofrefrigerators and other appliances. Compoundsinvolving boron are also present in pyrotechnicflares, because they emit a distinctive green color,and in the igniters of rockets.

Nitrogen

Though most people think of air as consistingprimarily of oxygen, in fact—as noted above—the greater part of the air we breathe is made up


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Nonmetalsof nitrogen. Scottish chemist David Rutherford(1749-1819) is usually given credit for discover-ing the element: in 1772, he identified nitrogen asthe element that remained when oxygen wasremoved from air. Several other scientists aroundthe same time made a similar discovery.

Because of its heavy presence in air, nitrogenis obtained primarily by cooling air to tempera-tures below the boiling points of its major com-ponents. Nitrogen boils (that is, turns into a gas)at a lower temperature than oxygen: -320.44°F (-195.8°C), as opposed to -297.4°F (-183°C).When air is cooled to -328°F (-200°C) and thenallowed to warm slowly, the nitrogen boils first,and therefore evaporates first. The nitrogen gas iscaptured, cooled, and liquefied once more.

Nitrogen can also be obtained from com-pounds such as potassium nitrate or saltpeter,found primarily in India; or sodium nitrate(Chilean saltpeter), which comes from the desertregions of Chile. To isolate nitrogen, variousprocesses are undertaken in a laboratory—forinstance, heating barium azide or sodium azide,both of which contain nitrogen.

REACTIONS WITH OTHER EL-

EMENTS. Existing as diatomic molecules,nitrogen forms very strong triple bonds, and as aresult tends to be fairly unreactive at low temper-atures. Thus, for instance, when a substanceburns, it reacts with the oxygen in the air, but notwith the nitrogen. At high temperatures, howev-er, nitrogen combines with other elements, react-ing with metals to form nitrides; with hydrogento form ammonia; with O2 to form nitrites; andwith O3 to form nitrates.

Nitrogen and oxygen, in particular, react athigh temperatures to form numerous com-pounds: nitric oxide (NO), nitrous oxide (N2O),nitrogen dioxide (NO2), dinitrogen trioxide(N2O3), and dinitrogen pentoxide (N2O5). Inreaction with halogens, nitrogen forms unstable,explosive compounds such as nitrogen trifluo-ride (NF3), nitrogen trichloride (NCl3), andnitrogen triiodide (NI3). One thing is for sure:never mix ammonia with bleach, which involvesthe halogen chlorine. Together, these two pro-duce chloramines, gases with poisonous vapors.

USES FOR NITROGEN. One ofthe most striking scenes in a memorable film,“Terminator 2: Judgment Day” (1991), occursnear the end of the movie, when the villainous T-1000 robot steps into a pool of liquid nitrogen

and cracks to pieces. As noted, nitrogen must beat extremely low temperatures to assume liquidform. Liquid nitrogen, which accounts for aboutone-third of all commercial uses of the element,is applied for quick-freezing foods, and for pre-serving foods in transit. Liquid nitrogen alsomakes it possible to process materials, such assome forms of rubber, that are too pliable formachining at room temperature. These materialsare first cooled in liquid nitrogen, and thenbecome more rigid.

In processing iron or steel, which formsundesirable oxides if exposed to oxygen, a blan-ket of nitrogen is applied to prevent this reaction.The same principle is applied in making comput-er chips and even in processing foods, since theseitems too are detrimentally affected by oxidation.Because it is far less combustible than air (mag-nesium is one of the few elements that burnsnitrogen in combustion), nitrogen is also used toclean tanks that have carried petroleum or othercombustible materials.

As noted, nitrogen combines with hydrogento form ammonia, used in fertilizers and cleaningmaterials. Ammonium nitrate, applied primarilyas a fertilizer, is also a dangerous explosive, asshown with horrifying effect in the bombing ofthe Alfred P. Murrah Federal Building in Okla-homa City on April 19, 1995—a tragedy thattook 168 lives. Nor is ammonium nitrate the onlynitrogen-based explosive. Nitric acid is used inmaking trinitrotoluene (TNT), nitroglycerin,and dynamite, as well as gunpowder and smoke-less powder.

NITROGEN, THE ENVIRON-

MENT, AND HEALTH. The nitrogencycle is the process whereby nitrogen passes fromthe atmosphere into living things, and ultimatelyback into the atmosphere. Both through theaction of lightning in the sky, and of bacteria inthe soil, nitrogen is converted to nitrates andnitrites—compounds of nitrogen and oxygen.These are then absorbed by plants to form plantproteins, which convert to animal proteins in thebodies of animals who eat the plants. When ananimal dies, the proteins are returned to the soil,and denitrifying bacteria break down these com-pounds, returning elemental nitrogen to theatmosphere.

Not all nitrogen in the atmosphere is health-ful, however. Oxides of nitrogen, formed in thehigh temperatures of internal combustion


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Nonmetals engines, pass into the air as nitric oxide. Thiscompound reacts readily with oxygen in the airto form nitrogen dioxide, a toxic reddish-browngas that adds to the tan color of smog over majorcities.

Another health concern is posed by sodiumnitrate and sodium nitrite, added to bacon,sausage, hot dogs, ham, bologna and other foodproducts to inhibit the growth of harmfulmicroorganisms. Many researchers believe thatnitrites impair the ability of a young child’s bloodto carry oxygen. Furthermore, nitrites often com-bine with amines, a form of organic compound,to create a variety of toxins known asnitrosoamines. Because of concerns about thesedangers, scientists and health activists have calledfor a ban on the use of nitrites and nitrates asfood additives.

Oxygen

Discovered independently by Swedish chemistCarl W. Scheele (1742-1786) and English chemistJoseph Priestley (1733-1804) in the period 1773-1774, oxygen was named by a third scientist,French chemist Antoine Lavoisier (1743-1794).Believing (incorrectly) that all acids containedthe newly discovered element, Lavoisier called itoxygen, which comes from a French word mean-ing “acid-former.”

Like many elements, oxygen has been in usesince the beginning of time. But this is quite dif-ferent from saying, for instance, that iron hasbeen in use since the early stages of human histo-ry. A person can live without iron (except for thenecessary quantities in the human body), andcan survive for weeks or even months withoutfood. One can live for a few days without water(the most famous and plentiful of all oxygen-containing compounds); but one cannot survivefor more than a few minutes without the oxygenin air.

Oxygen appears in three allotropes (differ-ent versions of the same element, distinguishedby molecular structure): monatomic oxygen (O);diatomic oxygen or O2; and triatomic oxygen(O3), better known as ozone. The diatomic formdominates the natural world, but in the upperatmosphere, ozone forms a protective layer thatkeeps the Sun’s harmful ultraviolet radiationfrom reaching Earth. Concerns that chlorofluo-rocarbons (CFCs) may be depleting the ozonelayer by converting these triatomic molecules to

O2 has led to a reduction in the output of CFCsby industrialized nations.

OBTAINING OXYGEN. Oxygen, ofcourse, is literally “in the air,” mixed with largerquantities of nitrogen. Higher up in the atmos-phere, it occurs as a free element. Through elec-trolysis,, it can be obtained from water; however,this process is prohibitively expensive for mostcommercial applications.

Oxygen-containing compounds are alsosources of oxygen for commercial use, but gener-ally oxygen is obtained by the fractional distilla-tion of liquid air, described above with regard tonitrogen. After the nitrogen has been separated,argon and neon (which also have lower boilingpoints than oxygen) also boil off, leaving behindan impure form of oxygen. This is further puri-fied by a process of cooling, liquefying, and evap-oration, which eliminates traces of noble gasessuch as krypton and xenon.

Many millions of years ago, when Earth wasfirst formed, there was no oxygen on the planet.The growth of oxygen on Earth coincided withthe development of organisms that, as theyevolved, increasingly needed oxygen. The presentconcentration of oxygen in the atmosphere,oceans, and the rocks of Earth’s crust was reachedabout 580 million years ago, and is sustainedtoday by biological activity. When plants under-go photosynthesis, carbon dioxide and waterreact in the presence of chlorophyll to producecarbohydrates and oxygen.

OXIDES AND OTHER COM-

POUNDS. Though the bond in diatomicoxygen is strong, once it is broken, monatomicoxygen reacts readily with other elements to forma seemingly limitless range of compounds:oxides, silicates, carbonates, phosphates, sulfates,and other more complex substances.

The process known as oxidation results inthe formation of numerous oxides. Sometimesoxygen and another element form several oxides,as for example in the case of nitrogen, whose fiveoxides are listed above. Water is an oxide; so tooare carbon dioxide and carbon monoxide. Whenanimals and plants die, the organic materials thatmake them up react with oxygen in the air,resulting in a complex form of oxidation knownas decay—or, in common language, “rotting.”

Oxygen, reacting with compounds such ashydrocarbons, produces carbon dioxide andwater vapor at high temperatures. If the oxida-


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Nonmetalstion process is extremely rapid, and takes place at high temperatures, it is usually identified ascombustion. In addition, oxygen reacts with ironand other metals to form oxides. Many ofthese oxides, commonly known as rust, are undesirable.

Every year, millions upon millions of dollarsare spent on painting metal structures, or forother precautions to protect against the forma-tion of metallic oxides. On the other hand,metallic oxides may be produced deliberately forapplications in materials such as mortar color, toenhance the appearance of a brick building.

USES OF OXYGEN. Aside from theobvious application of oxygen for breathing,there are four major fields that make use of thiselement: medicine, metallurgy, rocketry, and thefield of chemistry concerned with chemical syn-thesis. The medical application is closest to howwe normally use oxygen in our daily lives. In oxy-gen therapy, a patient having difficulty breathingis given doses of pure, or nearly pure, oxygen.This is used during surgical procedures, and totreat patients who have had heart attacks, as wellas those suffering from various infectious or res-piratory diseases.

The use of oxygen in metallurgy involvesrefining coke, which is almost pure carbon, tomake carbon monoxide. Carbon monoxide, inturn, reduces iron oxides to pure metallic iron.Oxygen is also used in blast furnaces to convertpig iron to steel by removing excess carbon, sili-con, and metallic impurities. In addition, oxygenis applied in torches for welding and cutting. Inthe form of liquefied oxygen, or LOX, oxygen isused in rockets and missiles. The space shuttle,for instance, carries a huge internal tank contain-ing oxygen and hydrogen, which, when theyreact, give the vehicle enormous thrust.

In chemical synthesis—the preparation ofcompounds (especially organic ones) from easilyavailable chemicals—commercial chemists useoxygen, for instance, to loosen the bonds inhydrocarbons. If this is done too quickly, itresults in combustion; but at a controlled rate,the chemical synthesis of hydrocarbons can gen-erate products such as acetylene, ethylene, andpropylene.

Oxygen can be used to produce syntheticfuels, as well as for water purification and sewagetreatment. Airplanes carry oxygen supplies incase of depressurization at high altitudes; in

addition, divers carry tanks in which oxygen ismixed with helium, rather than nitrogen, to prevent the dangerous condition known as “the bends.”

As early as the 1960s, smog-ridden citiessuch as Tokyo and Mexico City were equippedwith coin-operated oxygen booths—a sort of“phone booth for the lungs.” After inserting theappropriate amount of money, a person receiveda dose of oxygen for inhaling. This idea, spawnedby necessity, is the likely inspiration for a ratherbizarre fad that took hold in the trendier cities ofNorth America during the mid-1990s: oxygenbars. Popular in Los Angeles, New York, andToronto, these are establishments in whichpatrons pay up to a dollar per minute to inhalepure or flavored oxygen. Enthusiasts have toutedthe health benefits of this practice, but somephysicians have warned of oxygen toxicity andother dangers.

The Other “Orphan” Nonmetals

PHOSPHORUS. Some elements, suchas iron or gold, were known from ancient or evenprehistoric times—meaning that the identity ofthe discoverer is unknown. Phosphorus was thefirst element whose discoverer is known: Germanchemist Hennig Brand (c. 1630-c. 1692), whoidentified it in 1674.

Highly reactive with oxygen, phosphorus isused in the production of safety matches, smokebombs, and other incendiary devices. It is alsoimportant in fertilizers, and in various industrialapplications. Phosphorus forms a number ofimportant compounds, most notably phos-phates, on which animals and plants depend.

Phosphorus pollution, created by the use ofhousehold detergents containing phosphates,raised environmental concerns in the 1960s and1970s. It was feared that high phosphate levels inrivers and creeks would lead to runaway, detri-mental growth of plants and algae near bodies ofwater, a condition known as eutrophication.These concerns led to a ban on the use of phos-phates in detergents.

SULFUR. On its own, sulfur has nosmell, but in combination with other elements, itoften acquires a foul odor, which has given it anunpleasant reputation. The element’s smell, com-bined with its combustibility, led to the associa-


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Nonmetals


ALLOTROPES: Different versions of

the same element, distinguished by molec-

ular structure.



two atoms. This is in contrast to monatom-

ic elements.

ELECTROLYSIS: The use of an elec-

tric current to cause a chemical reaction.


table of elements, with valence electron

configurations of ns2np5. In contrast to the

noble gases, the halogens are known for

high levels of reactivity.

ION: An atom or group of atoms that

has lost or gained one or more electrons,

and thus has a net electric charge.






unstable. The unstable type, known as


METALS: Elements that are lustrous or

shiny in appearance; malleable, meaning

that they can be molded into different

shapes without breaking; and excellent

conductors of heat and electricity. Metals,

which constitute the vast majority of all

elements, tend to form positive ions by los-

ing electrons.

MONATOMIC: A term describing an

element that exists as single atoms. This in

contrast to diatomic elements.

NOBLE GASES: Group 8 of the peri-

odic table of elements, all of whom (with

the exception of helium) have valence elec-

tron configurations of ns2np6. The noble

gases are noted for their extreme lack of

reactivity—in other words, they tend not

to react to, or bond with, other elements.

NONMETALS: Elements that have a

dull appearance; are not malleable; are

poor conductors of heat and electricity;

and tend to gain electrons to form negative

ions. They are thus opposite of metals in

most regards, as befits their name. In addi-

tion to hydrogen, in Group 1 of the peri-

odic table, the other 18 nonmetals occupy

the upper right-hand side of the chart.

They include the noble gases, halogens,

and seven “orphan” elements: boron, car-

bon, nitrogen, oxygen, phosphorus, sulfur,

and selenium.














becomes stable.




transformed.







KEY TERMS

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Nonmetalstion of “brimstone”—the ancient name for sul-fur—with the fires of hell.

Because it is not usually combined withother elements in nature, the discovery of sulfurwas relatively easy. Pure, or nearly pure, sulfur ismined on the Gulf Coast of the United States, aswell as in Poland and Sicily. Sulfur compoundsalso appear in a number of ores, such as gypsum(calcium sulfate), or magnesium sulfate, betterknown as Epsom salts.

Applications of the aforementioned com-pounds are discussed in the Alkaline Earth Met-als essay. In addition, sulfates are used as agricul-tural insecticides, and for killing algae in watersupplies. Potassium aluminum sulfate, a gelati-nous solid that sinks to the bottom whendropped into water, is also used in water purifi-cation. Sulfur is sometimes applied in pure formas a fungicide, or in matches, fireworks, and gun-powder. More often, however, it is found in com-pounds such as the sulfates or the sulfides—including sulfuric acid and the evil-smelling gasknown as hydrogen sulfide.

Rotten eggs and intestinal gas are two exam-ples of hydrogen sulfide, which, though poison-ous, usually poses little danger, because the smellkeeps people away. Another sulfur compound ismercaptan, an ingredient in the skunk’s distinc-tive aroma. Tiny quantities of mercaptan areadded to natural gas (which has no odor) so thatdangerous gas leaks can be detected by smell.

SELENIUM. When Swedish chemistJons Berzelius (1779-1848) first discovered sele-nium in 1817, in deposits at the bottom of a tankin a sulfuric acid factory, he thought it was tel-lurium, a metalloid discovered in 1800. A fewmonths later, he reconsidered the evidence, andrealized he had found a new element. Becausetellurium, which lies just below selenium on theperiodic table, had been named for the Earth (tel-lus in Latin), he named his new discovery afterthe Greek word for the Moon, selene.

Found primarily in impurities from sulfideores, selenium is often obtained commercially as

a by-product of the refining of copper by elec-trolysis. It occurs in at least three allotropicforms, variously black and red in color. Plantsand animals, including humans, need smallamounts of selenium to survive, but larger quan-tities can be toxic. This was demonstrated in thelate 1970s, when waterfowl in the area of Kester-son Reservoir in northern California began turn-ing up with birth defects. The cause was latertraced to the dumping of selenium from agricul-tural wastes and industrial plants.

Because selenium is photovoltaic (able toconvert light directly into electricity) and photo-conductive (meaning that its resistance to theflow of electric current decreases in the presenceof light), it has applications in photocells, expo-sure meters, and solar cells. It is also used for theconversion of alternating current to direct cur-rent, and is applied as a semiconductor in elec-tronic and solid-state appliances. Photocopiersuse selenium in toners, and compounds contain-ing selenium are used to tint glass red, orange, orpink.

WHERE TO LEARN MORE

Beatty, Richard. Phosphorus. New York: BenchmarkBooks, 2001.

Blashfield, Jean F. Nitrogen. Austin, TX: Raintree Steck-Vaughn, 1999.

Fitzgerald, Karen. The Story of Oxygen. New York: F.Watts, 1996.

“Group Trends: Selected Nonmetals” (Web site).<http://wine1.sb.fsu.edu/chm1045/notes/Periodic/NMTrends/Period08.htm> (May 29, 2001).


“Nonmetals” Rutgers University Department of Chemistry(Web site). <http://chemistry.rutgers.edu/genchem/nonmet.html> (May 29, 2001).

“Nonmetals, Semimetals, and Their Compounds” (Website). <http://ull.chemistry.uakron.edu/genchem/21/title.html> (May 29, 2001).



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ME TALLO IDSMetalloids

CONCEPTThe term “metalloid” may sound like a referenceto a heavy-metal music fan, but in fact itdescribes a small collection of elements on theright-hand side of the periodic table. Forming adiagonal between boron and astatine, which liesfour rows down and four columns to the right ofboron, the metalloids are six elements that dis-play qualities of both metals and nonmetals.(Some classifications include boron and astatineas well, but in this book, they are treated as anonmetal and a halogen respectively.) Of thesesix—silicon, germanium, arsenic, antimony, tel-lurium, and polonium—only a few are house-hold names. People know that arsenic is poison,and they have some general sense that silicon isimportant in surgical implants. But most peopledo not know that silicon, also the principal mate-rial in sand and glass, is the second-most plenti-ful element on Earth. Without silicon, fromwhich computer chips are made, our computer-based society simply could not exist.

HOW IT WORKS

Families and “Orphans”

Most elements fit into some sort of “family”grouping: the alkali metals, the alkaline earthmetals, transition metals, lanthanides, actinides,halogens, and noble gases. These seven families,five metallic and two nonmetallic, account for 91of 112 elements on the periodic table.

The “orphan” elements, or those not readilyclassifiable within a family, occupy groups 3through 6 on the North American version of the

periodic table, which only numbers the “tall”columns.

Twenty elements, occupying a rectangle thatstretches across four groups or columns, and fiveperiods or rows, are “orphans.” (This is in addi-tion to hydrogen, on Period 1, Group 1.) Theseare best classified, not by family, but by cross-family characteristics that unite large groups ofelements. These cross-family groupings can belikened to nations: on one level, a person identi-fies with his or her family (“I’m a Smith”); but onanother level, a person has a national identity(“I’m an American.”)

Likewise gold, for instance, is both a mem-ber of the transition metals family, and of thelarger metals grouping. The “orphan” elements,because they have no family, are best identifiedwith characteristics that unite a large body of ele-ments. (For this reason, the “orphan” metals arediscussed in the Metals essay, and the “orphan”nonmetals in the Nonmetals essay.)

Between Metals and Nonmetals

Of the twenty “orphan” elements in groups 3through 6, seven are metals and seven nonmetals.Between them runs a diagonal, comprising thesix metalloids. (As noted earlier, boron is some-times considered a metalloid, but in this book, itis discussed in the Nonmetals essay, while asta-tine—also sometimes grouped with the metal-loids—is treated in the Halogens essay.) What isa metalloid? The best way to answer that ques-tion is by evaluating the differences between ametal and a nonmetal.

On the periodic table, metals fill the left,center, and part of the right-hand side of the

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Metalloids


chart, as well as the two rows—corresponding tothe lanthanides and actinides—placed separatelyat the bottom. Thus it should not come as a sur-prise that most elements (87, in fact) are metals.Metals are lustrous or shiny in appearance, andductile, meaning that they can be molded intodifferent shapes without breaking. They areexcellent conductors of heat and electricity, andtend to form positive ions by losing electrons.

Nonmetals, which occupy the upper right-hand side of the periodic table, include the noblegases, halogens, and the seven “orphan” elementsalluded to above. With the addition of hydrogen(which, as noted, is covered in a separate essay),they comprise 19 elements. Whereas all metals

are solids, except for mercury—a liquid at roomtemperature—the nonmetals are a collection ofgases, solids, and one other liquid, bromine. (Notsurprisingly, all gaseous elements are nonmetals.)As their name suggests, nonmetals are oppositeto metals in most regards: dull in appearance,they are not particularly ductile or malleable.With the exception of carbon, they are poor con-ductors of heat and electricity, and tend to gainelectrons to form negative ions.

Survey of the Metalloids

The metalloids can thus be defined as those ele-ments which exhibit characteristics of both met-als and nonmetals. They are all solids, but not

WITHOUT SILICON, FROM WHICH COMPUTER CHIPS ARE MADE, OUR COMPUTER-BASED SOCIETY SIMPLY COULD NOT

EXIST. HERE, A WORKER HOLDS A SILICON WAFER WITH INTEGRATED CIRCUITS. (Charles O’Rear/Corbis. Reproduced by

permission.)

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Metalloids wise known from ancient times, and bothattracted the attention of alchemists—medievalmystics, in many regards the forerunners ofchemists, who believed it was possible to changeplain metals into gold.

Traces of the alchemical mindset remained apart of European science as late as the mid-eigh-teenth century, a tendency exemplified by the factthat tellurium, discovered around that time, wasbelieved to be “unripe gold.” On the other hand,the discovery of germanium more than a centurylater—following predictions of its existencebased on the periodic table—reflected a muchgreater degree of scientific rigor. Finally, the iso-lation of polonium, known for its radioactivity,was evidence of an even more advanced stage ofdevelopment in chemistry and science as awhole.


Silicon

Swedish chemist Jons Berzelius (1779-1848) dis-covered silicon in 1823, but humans had beenusing the element—in the form of compoundsknown as silicates—for thousands of years.Indeed, silicon may have been one of the first ele-ments formed, many millions of years before lifeappeared on this planet. Geologists believe thatthe Earth was once composed primarily ofmolten iron, oxygen, silicon, and aluminum, stillthe predominant elements in the planet’s crust.Because iron has a greater atomic mass, it settledtoward the center, while the more lightweightelements rose to the surface.

Silicon is found in everything from the Sunand other stars, as well as meteorites, to plantsand animal bones. Many hundreds of mineralson Earth contain silicon and oxygen in variousforms: for instance, silicon and oxygen form sili-ca (SiO2), commonly known as sand. This sand ismixed with lime and soda (sodium carbonate), aswell other substances, to make glass. The purestvarieties of silica form quartz, known for its clearcrystals, while impure quartz forms crystals ofsemiprecious gems such as amethyst, opal,and agate.

COMPOUNDS: SILICATES AND

SILICONES. Silicon is directly below car-bon on the periodic table, and the elements in


lustrous, and conduct heat and electricity mod-erately well.

The six metalloids treated here are listedbelow, with the atomic number and chemicalsymbol of each.

• 14. Silicon (Si)

• 32. Germanium (Ge)

• 33. Arsenic (As)

• 51. Antimony (Sb)

• 52. Tellurium (Te)

• 84. Polonium (Po)

Silicon is the second most abundant elementon Earth, comprising 25.7% of the planet’sknown elemental mass. Together with oxygen, itaccounts for nearly three-quarters of the knowntotal. (The planet’s core is probably composedlargely of iron, with significant deposits of urani-um and other metals between the crust and thecore, but geologists can only make educatedguesses as to the elemental composition beneaththe crust.)

THE METALLOIDS IN HISTO-

RY. By far the most important of the metal-loids, silicon has been in use as a compound forcenturies, but was only isolated in the early nine-teenth century. Arsenic and antimony were like-

POLONIUM WAS FIRST ISOLATED BY MARIE AND PIERRE

CURIE. MARIE CURIE NAMED THE ELEMENT IN HONOR

OF HER HOMELAND, POLAND.

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Metalloidsthat group (which includes germanium, seleni-um, and lead) are sometimes called the “carbonfamily.” Like carbon, silicon forms a huge array ofcompounds, because it has four electrons toshare in chemical bonding, and thus can formlong strings of atoms. Because silicon atoms aremuch larger than carbon atoms, however, not asmany other atoms can get close to silicon. Thusthe number of possible compounds involving sil-icon—though still quite impressive—is less thanthe many millions of carbon-based compounds

Often silicon is combined with oxygen inlong chains that use the oxygen atoms as separa-tors. Because oxygen has a valence of two, mean-ing that it can bond to two other atoms at a time,it bonds easily between two other silicon atoms.The silicon, meanwhile, has two other electronsto use in bonding, so it typically attaches to twomore, non-bridging, oxygen atoms. Compoundsof this type are known as silicates, and most ofthe rocks and clay on Earth—with the exceptionof lime, which is calcium oxide—are silicatesmade up of combinations of silicon, oxygen, andvarious metals such as aluminum, iron, sodium,and potassium.

Silicones are another variety of compoundinvolving silicon atoms strung together by bridg-ing oxygen atoms. Instead of attaching to twoother, non-bridging oxygen atoms, as in a sili-cate, the silicon atoms in a silicone attach toorganic groups—that is, molecules containingcarbon. Because silicones often resemble organiccompounds such as oils, greases, and rubber, sil-icone oils are frequently used in place of organicpetroleum as a lubricant because they can with-stand greater variations in temperature. Siliconerubbers are used in everything from bouncingballs to space vehicles. And because the body tol-erates the introduction of silicone implants bet-ter than it does organic ones, silicones are used insurgical implants as well.

THE MANY OTHER USES OF

SILICON. Silicones are also present in elec-trical insulators, rust preventives, fabric soften-ers, hair sprays, hand creams, furniture and auto-mobile polishes, paints, adhesives (includingbathtub sealers), and chewing gum. Yet even thislist does not exhaust the many applications ofelemental silicon, as opposed to silicone.

Due to its semi-metallic qualities, silicon isused as a semiconductor of electricity. Computerchips are tiny slices of ultra-pure silicon, etched

with as many as half a million microscopic andintricately connected electronic circuits. Thesechips manipulate voltages using binary codes, forwhich 1 means “voltage on” and 0 means “voltageoff.” By means of these pulses, silicon chips per-form multitudes of calculations in seconds—cal-culations that would take humans hours ormonths or even years.

A porous form of silica, known as silica gel,absorbs water vapor from the air, and is oftenpacked alongside moisture-sensitive productssuch as electronics components, in order to keepthem dry. Silicon carbine, an extremely hardcrystalline material manufactured by fusing sandwith coke (almost pure carbon) at high tempera-tures, has applications as an abrasive.

Arsenic

Silicon is unquestionably the “star” of the metal-loids, but several other elements in this broadgrouping have been known since ancient times.Among these is arsenic, well-known from many adetective novel. Highly poisonous, it is often afixture of such stories: in a typical plot, a greedynephew drops some arsenic into the tea of anelderly aunt who has left money to him in herwill, and it is up to the detective/hero/heroine touncover the details behind this dastardly scheme.The use of arsenic as a prop in such tales hadalready become something of a cliché by 1941,when “Arsenic and Old Lace,” a play by JosephKesselring that satirized an old-fashioned “who-dunit,” made its debut on Broadway.

Known since ancient times, arsenic is namedafter the Greek word arsenikon, meaning “yelloworpiment.” (Orpiment is arsenic trisulfide, orAs2S3.) But the Greeks were not the only peoplesin antiquity who knew of arsenic and its toxicproperties: the Egyptians and Chinese were like-wise familiar with it. Alchemists took an interestin arsenic, and one of them, Albertus Magnus(1193-1280)—a German physician who con-tributed significantly to the rebirth of interest inscience during the late Middle Ages—probablyisolated the element from arsenic trisulfide inabout 1250.

Arsenic is more than just a prop in a detec-tive story. (And in fact it is not a very effectivepoison for a murderer who hopes to get awaywith his crime: traces of arsenic remain in thebody of a murder victim—even in the person’shair—for years after death.) Today, arsenic is


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Metalloids used for purposes such as bronzing, and forhardening and improving the sphericity (round-ness) of lead shot in shotgun shells.

In solid-state devices such as transistors,arsenic or antimony is used for the purposes of“doping.” This may sound like another version ofthe whodunit-style poisoning described above,but actually doping involves the alteration ofchemical properties by deliberately introducingimpurities. Through the doping of silicon, itselectrical conductivity is improved.

Antimony

In ancient times, antimony was known by theLatin name stibium; hence its chemical symbol ofSb. The word, which can be deciphered as “ametal that does not occur by itself,” has rootsmuch older than the Roman civilization, howev-er. Many centuries before the rise of the Romans,the biblical prophet Jeremiah railed againstwomen’s use of “stibic stone” to paint their faces.In fact, stibic stone is antimony (III) sulfide, orSb2S3, a black mineral which has long been usedas an eyeliner.

During the Middle Ages, alchemists studiedantimony and its compounds, but the first scien-tist to distinguish between the compounds andthe element itself was French chemist Jean-Bap-tiste Buquet, in 1771. During the early modernera, antimony was used in making bells, tools,and type for printing presses, because its addi-tion to an alloy strengthened the other metals.Today, antimony is added to lead in storage bat-teries to make the lead harder and stronger. It isstill used in metal type, as well as bullets andcable sheathing.

Antimony was a component in early vari-eties of matches, but by 1855, the less volatilephosphorus had replaced it. In altered form,however, antimony still appears in some match-es. Oxides and sulfides of antimony are alsoapplied in paints, ceramics, fireworks, pharma-ceuticals, and other products. These are also uti-lized for dyeing cloth and fireproofing materials.In addition, as noted above with regard toarsenic, antimony is applied for the doping of sil-icon and other materials in semiconductors, par-ticularly those that use infrared detectors.

Tellurium

When Hungarian chemist Joseph Ramacsaházyfirst described tellurium in the mid-eighteenth

century, he did so in alchemical terms, referringto it as “unripe gold.” This reflected the belief, stilllingering from the Middle Ages, that metals“grow” to higher states. At least part of his confu-sion is understandable, since the element didappear in minerals that also contained gold.

A few years later, Baron Franz Joseph Müllervon Reichenstein, an Austrian mining inspectorin Transylvania, analyzed the substance, which heconcluded was bismuth sulfide. Later he changedhis mind, then subjected the material to a seriesof tests over a period of many years, until finallyhe sent a sample to the distinguished Germanchemist, Martin Heinrich Klaproth (1743-1817).Klaproth concluded that it was indeed a new ele-ment, and suggested the Latin word tellus, for“earth,” as a name.

Grayish-white and lustrous, tellurium lookslike a metal, but it is much more brittle than mostmetallic elements. Furthermore, it is a semicon-ductor, which further separates it from the typically conductive metals. Its semiconductiveproperties have led to applications in buildingelectronic components. Added to copper andstainless steel, tellurium improves the machin-ability of those metals; in addition, telluriumcompounds are used for the coloring of glass,porcelain, enamel, and ceramics.

Germanium and Polonium

Germanium. Both named after European coun-tries, germanium and polonium were identifiedmuch later than the other metalloids. In 1871,two years after he created his famous periodictable, Russian chemist Dmitri IvanovitchMendeleev (1834-1907) predicted the existenceof an element he called “eka-silicon.” This turnedout to be germanium, which sits directly beneathsilicon on the periodic table, and which was dis-covered in 1886 by German chemist ClemensAlexander Winkler (1838-1904).

Germanium is principally used as a dopingagent in making transistor elements. It is alsoused as a phosphor in fluorescent lamps. In addition, germanium and germanium oxide,transparent to the infrared portion of the elec-tromagnetic spectrum, are present in infraredspectroscopes and infrared detectors.

POLONIUM. While the applications ofgermanium are limited compared to those of, say,silicon, polonium is useful primarily as a source


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Metalloids


of neutrons in atomic laboratories—hardlysomething with which a person comes in contactduring the course of an ordinary day. Yet the iso-lation of polonium by Polish-French physicistand chemist Marie Curie (1867-1934), alongwith her husband Pierre (1859-1906), a Frenchphysicist, was a prelude to one of the greatest sto-ries in the history of chemistry.

Setting out to find radioactive elementsother than uranium, the Curies refined a largequantity of pitchblende, an ore commonly found

in uranium mines. Within a year, they had dis-covered polonium, which Marie named after herhomeland. Polonium is about 100 times asradioactive as uranium, but the Curies were cer-tain that yet another element lay in the pitch-blende, and they devoted the next four years tothe exhausting process of isolating it. The story ofradium, its isolation, and the great price MarieCurie paid for working with radioactive materi-als is discussed in the essay on Alkaline EarthMetals.











METALLOIDS: Elements which exhibit

characteristics of both metals and non-

metals. Metalloids are all solids, but are not

lustrous or shiny, and they conduct heat

and electricity moderately well. The six

metalloids occupy a diagonal region

between the metals and nonmetals on the

right side of the periodic table. Sometimes

boron and astatine are included with the

metalloids, but in this volume, they are

treated as an “orphan” nonmetal and a

halogen, respectively.

METALS: Elements that are lustrous, or

shiny in appearance, and malleable, mean-

ing that they can be molded into different

shapes without breaking. They are excel-

lent conductors of heat and electricity,

and tend to form positive ions by losing

electrons.

NONMETALS: Elements that have a

dull appearance; are not malleable; are

poor conductors of heat and electricity;

and tend to gain electrons to form negative

ions. They are thus the opposite of metals

in most regards, as befits their name.





group number—that is, the two “tall”

columns to the left of the transition metals,

as well as the six “tall” columns to the right.

“ORPHAN”: An element that does not

belong to any clearly defined family of ele-

ments. The metalloids are all “orphans.”









becomes stable.

KEY TERMS

set_vol1_sec6 9/13/01 11:50 AM

MetalloidsWHERE TO LEARN MORE

“The Arsenic Website Project.” Harvard University Depart-ment of Physics (Web site). <http://phys4.harvard.edu/~wilson/arsenic_project_main.html> (May 30,2001).


“Metalloids.” ChemicalElements.com (Web site).<http://www.chemicalelements.com/groups/metal-loids.html> (May 30, 2001).

“Metalloids” (Web site). <http://www.wealthhealthand-wisdom.com/metalloids.asp> (May 30, 2001).

Pflaum, Rosalynd. Marie Curie and Her Daughter Irène.Minneapolis, MN: Lerner Publications, 1993.

Thomas, Jens. Silicon. New York: Benchmark Books,2001.


set_vol1_sec6 9/13/01 11:50 AM


HALOGENSHalogens

CONCEPTTable salt, bleach, fluoride in toothpaste, chlorinein swimming pools—what do all of these have incommon? Add halogen lamps to the list, and theanswer becomes more clear: all involve one ormore of the halogens, which form Group 7 of theperiodic table of elements. Known collectively bya term derived from a Greek word meaning “salt-producing,” the halogen family consists of fiveelements: fluorine, chlorine, bromine, iodine,and astatine. The first four of these are widelyused, often in combination; the last, on the otherhand, is a highly radioactive and extremely raresubstance. The applications of halogens are manyand varied, including some that are dangerous,controversial, and deadly.

HOW IT WORKS

The Halogens on the Periodic Table

As noted, the halogens form Group 7 of the peri-odic table of elements. They are listed below,along with chemical symbol and atomic number:

• Fluorine (F) 9

• Chlorine (Cl): 17

• Bromine (Br): 35

• Iodine (I): 53

• Astatine (At): 85

On the periodic table, as displayed in chem-istry labs around the world, the number ofcolumns and rows does not vary, since these con-figurations are the result of specific and interre-lated properties among the elements. There arealways 18 columns; however, the way in whichthese are labeled differs somewhat from place to

place. Many chemists outside the United Statesrefer to these as 18 different groups of elements;however, within the United States, a somewhatdifferent system is used.

In many American versions of the chart,there are only eight groups, sometimes designat-ed with Roman numerals. The 40 transition met-als in the center are not designated by groupnumber, nor are the lanthanides and actinides,which are set apart at the bottom of the periodictable. The remaining eight columns are the onlyones assigned group numbers. In many ways, thisis less useful than the system of 18 group num-bers; however, it does have one advantage.

ELECTRON CONFIGURATIONS

AND BONDING. In the eight-group sys-tem, group number designates the number ofvalence electrons. The valence electrons, whichoccupy the highest energy levels of an atom, arethe electrons that bond one element to another.These are often referred to as the “outer shell” ofan atom, though the actual structure is muchmore complex. In any case, electron configura-tion is one of the ways halogens can be defined:all have seven valence electrons.

Because the rows in the periodic table indi-cate increasing energy levels, energy levels rise asone moves up the list of halogens. Fluorine, onrow 2, has a valence-shell configuration of 2s22p5;while that of chlorine is 3s23p5. Note that onlythe energy level changes, but not the electronconfiguration at the highest energy level. Thesame goes for bromine (4s24p5), iodine (4s24p5),and astatine (5s25p5).

All members of the halogen family have thesame valence-shell electron configurations, andthus tend to bond in much the same way. As we

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Halogens


shall see, they are inclined to form bonds morereadily than most other substances, and indeedfluorine is the most reactive of all elements.

Thus it is ironic that they are “next door” tothe Group 8 noble gases, the least reactive amongthe elements. The reason for this, as discussed inthe Chemical Bonding essay, is that most ele-ments bond in such a way that they develop avalence shell of eight electrons; the noble gasesare already there, so they do not bond, except insome cases—and then principally with fluorine.

Characteristics of the Halogens

In terms of the phase of matter in which they arenormally found, the halogens are a varied group.Fluorine and chlorine are gases, iodine is a solid,and bromine is one of only two elements thatexists at room temperature as a liquid. As forastatine, it is a solid too, but so highly radioactivethat it is hard to know much about its properties.

Despite these differences, the halogens havemuch in common, and not just with regard totheir seven valence electrons. Indeed, they wereidentified as a group possessing similar charac-teristics long before chemists had any way ofknowing about electrons, let alone electron con-

figurations. One of the first things scientistsnoticed about these five elements is the fact thatthey tend to form salts. In everyday terminology,“salt” refers only to a few variations on the samething—table salt, sea salt, and the like. In chem-istry, however, the meaning is much broader: asalt is defined as the result of bonding betweenan acid and a base.

Many salts are formed by the bonding of ametal and a nonmetal. The halogens are all non-metals, and tend to form salts with metals, as inthe example of sodium chloride (NaCl), a bondbetween chlorine, a halogen, and the metal sodi-um. The result, of course, is what people com-monly call “salt.” Due to its tendency to formsalts, the first of the halogens to be isolated—chlorine, in 1811—was originally named “halo-gen.” This is a combination of the Greek wordshalos, or salt, and gennan, “to form or generate.”

In their pure form, halogens are diatomic,meaning that they exist as molecules with twoatoms: F2, Cl2, and so on. When bonding withmetals, they form ionic bonds, which are thestrongest form of chemical bond. In the process,halogens become negatively charged ions, oranions. These are represented by the symbols F-,Cl-, Br-, and I-, as well as the names fluoride,chloride, bromide, and iodide. All of the halogens

SALT, LIKE THE MOUND SHOWN HERE, IS A SAFE AND COMMON SUBSTANCE THAT CONTAINS CHLORINE. (Charles

O’Rear/Corbis. Reproduced by permission.)

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Halogensare highly reactive, and will combine directlywith almost all elements.

Due to this high level of reactivity, the halo-gens are almost never found in pure form; rather,they have to be extracted. Extraction of halogensis doubly problematic, because they are danger-ous. Exposure to large quantities can be harmfulor fatal, and for this reason halogens have beenused as poisons to deter unwanted plants andinsects—and, in one of the most horrifyingchapters of twentieth century military history, asa weapon in World War I.


Chlorine

Chlorine is a highly poisonous gas, greenish-yellow in color, with a sharp smell that induceschoking in humans. Yet, it can combine withother elements to form compounds safe forhuman consumption. Most notable among thesecompounds is salt, which has been used as a foodpreservative since at least 3000 B.C.

Salt, of course, occurs in nature. By contrast,the first chlorine compound made by humanswas probably hydrochloric acid, created by dis-solving hydrogen chloride gas in water. The firstscientist to work with hydrochloric acid was Per-sian physician and alchemist Rhazes (ar-Razi;c. 864-c. 935), one of the most outstanding scien-tific minds of the medieval period. Alchemists,who in some ways were the precursors of truechemists, believed that base metals such as ironcould be turned into gold. Of course this is notpossible, but alchemists in about 1200 did at least succeed in dissolving gold using a mix-ture of hydrochloric and nitric acids known asaqua regia.

The first modern scientist to work withchlorine was Swedish chemist Carl W. Scheele(1742-1786), who also discovered a number ofother elements and compounds, including bari-um, manganese, oxygen, ammonia, and glycerin.However, Scheele, who isolated it in 1774,thought that chlorine was a compound; only in1811 did English chemist Sir Humphry Davy(1778-1829) identify it as an element. Anotherchemist had suggested the name “halogen” forthe alleged compound, but Davy suggested that it


be called chlorine instead, after the Greek wordchloros, which indicates a sickly yellow color.

USES OF CHLORINE. The dangersinvolved with chlorine have made it an effectivesubstance to use against stains, plants, animals—and even human beings. Chlorine gas is highlyirritating to the mucous membranes of the nose,mouth, and lungs, and it can be detected in air ata concentration of only 3 parts per million(ppm).

The concentrations of chlorine used againsttroops on both sides in World War I (beginningin 1915) was, of course, much higher. Thanks tothe use of chlorine gas and other antipersonnelagents, one of the most chilling images to emergefrom that conflict was of soldiers succumbing topoisonous gas. Yet just as it is harmful tohumans, chlorine can be harmful to microbes,thus preserving human life. As early as 1801, ithad been used in solutions as a disinfectant; in1831, its use in hospitals made it effective as aweapon against a cholera epidemic that sweptacross Europe.

Another well-known use of chlorine is as ableaching agent. Until 1785, when chlorine was

THIS WORKER IS REMOVING FREON FROM AN OLD

REFRIGERATOR. CFCS LIKE FREON MAY BE RESPONSI-BLE FOR THE DEPLETION OF THE OZONE LAYER. (James A.

Sugar/Corbis. Reproduced by permission.)

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Halogens first put to use as a bleach, the only way to getstains and unwanted colors out of textiles orpaper was to expose them to sunlight, not alwaysan effective method. By contrast, chlorine, stillused as a bleach today, can be highly effective—agood reason not to use regular old-fashionedbleach on anything other than white clothing.(Since the 1980s, makers of bleaches have devel-oped all-color versions to brighten and take outstains from clothing of other colors.)

Calcium hydrocholoride (CaOCl), both ableaching powder and a disinfectant used inswimming pools, combines both the disinfectantand bleaching properties of chlorine. This andthe others discussed here are just some of many,many compounds formed with the highly reac-tive element chlorine. Particularly notable—andcontroversial—are compounds involving chlo-rine and carbon.

CHLORINE AND ORGANIC

COMPOUNDS. Chlorine bonds well withorganic substances, or those containing carbon.In a number of instances, chlorine becomes partof an organic polymer such as PVC (polyvinylchloride), used for making synthetic pipe. Chlo-rine polymers are also applied in making syn-thetic rubber, or neoprene. Due to its resistanceto heat, oxidation, and oils, neoprene is used in anumber of automobile parts.

The bonding of chlorine with substancescontaining carbon has become increasingly con-troversial because of concerns over health andthe environment, and in some cases chlorine-carbon compounds have been outlawed. Suchwas the fate of DDT, a pesticide soluble in fatsand oils rather than in water. When it was dis-covered that DDT was carcinogenic, or cancer-causing, in humans and animals, its use in theUnited States was outlawed.

Other, less well-known, chlorine-relatedinsecticides have likewise been banned due totheir potential for harm to human life and theenvironment. Among these are chlorine-contain-ing materials once used for dry cleaning. Alsonotable is the role of chlorine in chlorofluorocar-bons (CFCs), which have been used in refriger-ants such as Freon, and in propellants for aerosolsprays. CFCs tend to evaporate easily, and con-cerns over their effect on Earth’s atmospherehave led to the phasing out of their use.

Fluorine

Fluorine has the distinction of being the mostreactive of all the elements, with the highest elec-tronegativity value on the periodic table. Becauseof this, it proved extremely difficult to isolate.Davy first identified it as an element, but was poi-soned while trying unsuccessfully to decomposehydrogen fluoride. Two other chemists were alsolater poisoned in similar attempts, and one ofthem died as a result.

French chemist Edmond Fremy (1814-1894)very nearly succeeded in isolating fluorine, andthough he failed to do so, he inspired his studentHenri Moissan (1852-1907) to continue the proj-ect. One of the problems involved in isolatingthis highly reactive element was the fact that ittends to “attack” any container in which it isplaced: most metals, for instance, will burst intoflames in the presence of fluorine. Like the othersbefore him, Moissan set about to isolate fluorinefrom hydrogen fluoride by means of electroly-sis—the use of an electric current to cause achemical reaction—but in doing so, he used aplatinum-iridium alloy that resisted attacks byfluorine. In 1906, he received the Nobel Prize forhis work, and his technique is still used today inmodified form.

PROPERTIES AND USES OF

FLUORINE. A pale green gas of low density,fluorine can combine with all elements exceptsome of the noble gases. Even water will burn inthe presence of this highly reactive substance.Fluorine is also highly toxic, and can cause severeburns on contact, yet it also exists in harmlesscompounds, primarily in the mineral known asfluorspar, or calcium fluoride. The latter gives offa fluorescent light (fluorescence is the term for atype of light not accompanied by heat), and flu-orine was named for the mineral that is one of itsprincipal “hosts”.

Beginning in the 1600s, hydrofluoric acidwas used for etching glass, and is still used forthat purpose today in the manufacture of prod-ucts such as light bulbs. The oil industry uses it asa catalyst—a substance that speeds along a chem-ical reaction—to increase the octane number ingasoline. Fluorine is also used in a polymer com-monly known as Teflon, which provides a non-stick surface for frying pans and other cooking-related products.

Just as chlorine saw service in World War I,fluorine was enlisted in World War II to create a


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Halogensweapon far more terrifying than poison gas: theatomic bomb. Scientists working on the Manhat-tan Project, the United States’ effort to developthe bombs dropped on Japan in 1945, neededlarge quantities of the uranium-235 isotope. Thisthey obtained in large part by diffusion of thecompound uranium hexafluoride, which consistsof molecules containing one uranium atom andsix fluorine anions.

FLUORIDATION OF WATER.

Long before World War II, health officials in theUnited States noticed that communities havinghigh concentration of fluoride in their drinkingwater tended to suffer a much lower incidence oftooth decay. In some areas the concentration offluoride in the water supply was high enoughthat it stained people’s teeth; still, at the turn ofthe century—an era when dental hygiene as weknow it today was still in its infancy—the pre-vention of tooth decay was an attractiveprospect. Perhaps, officials surmised, it would bepossible to introduce smaller concentrations offluoride into community drinking water, with aresulting improvement in overall dental health.

After World War II, a number of municipal-ities around the United States undertook thefluoridation of their water supplies, using con-centrations as low as 1 ppm. Within a few years,fluoridation became a hotly debated topic, withproponents pointing to the potential health ben-efits and opponents arguing from the standpointof issues not directly involved in science. It wasan invasion of personal liberty, they said, for gov-ernments to force citizens to drink water whichhad been supplemented with a foreign substance.

During the 1950s, in fact, fluoridationbecame associated in some circles with Commu-nism—just another manifestation of a govern-ment trying to control its citizens. In later years,ironically, antifluoridation efforts became associ-ated with groups on the political left rather thanthe right. By then, the argument no longerrevolved around the issue of government power;instead the concern was for the health risksinvolved in introducing a substance lethal inlarge doses.

Fluoride had meanwhile gained applicationin toothpastes. Colgate took the lead, introducing“stannous fluoride” in 1955. Three years later, thecompany launched a memorable advertisingcampaign with commercials in which a little girlshowed her mother a “report card” from the den-

tist and announced “Look, Ma! No cavities!”Within a few years, virtually all brands of tooth-paste used fluoride; however, the use of fluoridein drinking water remained controversial.

As late as 1993, in fact, the issue of fluorida-tion remained heated enough to spawn a studyby the U.S. National Research Council. Thecouncil found some improvement in dentalhealth, but not as large as had been claimed byearly proponents of fluoridation. Furthermore,this improvement could be explained by refer-ence to a number of other factors, including fluoride in toothpastes and a generally height-ened awareness of dental health among the U.S.populace.

C H LO R O F LU O R O CA R B O N S .

Another controversial application of fluorine isits use, along with chlorine and carbon, in chlo-rofluorocarbons. As noted above, CFCs havebeen used in refrigerants and propellants; anoth-er application is as a blowing agent forpolyurethane foam. This continued for severaldecades, but in the 1980s, environmentalistsbecame concerned over depletion of the ozonelayer high in Earth’s atmosphere.

Unlike ordinary oxygen (O2), ozone or O3 iscapable of absorbing ultraviolet radiation fromthe Sun, which would otherwise be harmful tohuman life. It is believed that CFCs catalyze theconversion of ozone to oxygen, and that this mayexplain the “ozone hole,” which is particularlynoticeable over the Antarctic in September andOctober.

As a result, a number of countries signed anagreement in 1996 to eliminate the manufactureof halocarbons, or substances containing halo-gens and carbon. Manufacturers in countries thatsigned this agreement, known as the MontrealProtocol, have developed CFC substitutes, mostnotably hydrochlorofluorocarbons (HCFCs),CFC-like compounds also containing hydrogenatoms.

The ozone-layer question is far from settled,however. Critics argue that in fact the depletionof the ozone layer over Antarctica is a naturaloccurrence, which may explain why it onlyoccurs at certain times of year. This may alsoexplain why it happens primarily in Antarctica,far from any place where humans have beenusing CFCs. (Ozone depletion is far less signifi-cant in the Arctic, which is much closer to thepopulation centers of the industrialized world.)


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Halogens In any case, natural sources, such as volcanoeruptions, continue to add halogen compoundsto the atmosphere.

Bromine

Bromine is a foul-smelling reddish-brown liquidwhose name is derived from a Greek word mean-ing “stink.” With a boiling point much lower thanthat of water—137.84°F (58.8°C)—it readilytransforms into a gas. Like other halogens, itsvapors are highly irritating to the eyes and throat.It is found primarily in deposits of brine, a solu-tion of salt and water. Among the most signifi-cant brine deposits are in Israel’s Dead Sea, aswell as in Arkansas and Michigan.

Credit for the isolation of bromine is usual-ly given to French chemist Antoine-JérômeBalard (1802-1876), though in fact Germanchemist Carl Löwig (1803-1890) actually isolatedit first, in 1825. However, Balard, who publishedhis results a year later, provided a much moredetailed explanation of bromine’s properties.

The first use of bromine actually predatedboth men by several millennia. To make theirfamous purple dyes, the Phoenicians used murexmollusks, which contained bromine. (Like thenames of the halogens, the word “Phoenicians”is derived from Greek—in this case, a wordmeaning “red” or “purple,” which referred totheir dyes.) Today bromine is also used in dyes,and other modern uses include applications inpesticides, disinfectants, medicines, and flameretardants.

At one time, a compound containingbromine was widely used by the petroleumindustry as an additive for gasoline containinglead. Ethylene dibromide reacts with the leadreleased by gasoline to form lead bromide(PbBr2), referred to as a “scavenger,” because ittends to clean the emissions of lead-containinggasoline. However, leaded gasoline was phasedout during the late 1970s and early 1980s; as aresult, demand for ethylene dibromide droppedconsiderably.

HALOGEN LAMPS. The name“halogen” is probably familiar to most peoplebecause of the term “halogen lamp.” Used forautomobile headlights, spotlights, and flood-lights, the halogen lamp is much more effectivethan ordinary incandescent light. Incandescent“heat-producing” light was first developed in the1870s and improved during the early part of the

twentieth century with the replacement of car-bon by tungsten as the principal material in thefilament, the area that is heated.

Tungsten proved much more durable thancarbon when heated, but it has a number ofproblems when combined with the gases in anincandescent bulb. As the light bulb continues toburn for a period of time, the tungsten filamentbegins to thin and will eventually break. At thesame time, tungsten begins to accumulate on thesurface of the bulb, dimming its light. However,by adding bromine and other halogens to thebulb’s gas filling—thus making a halogen lamp—these problems are alleviated.

As tungsten evaporates from the filament, itcombines with the halogen to form a gaseouscompound that circulates within the bulb.Instead of depositing on the surface of the bulb,the compound remains a gas until it comes intocontact with the filament and breaks down. It isthen redeposited on the filament, and the halo-gen gas is free to combine with newly evaporatedtungsten. Though a halogen bulb does eventual-ly break down, it lasts much longer than an ordi-nary incandescent bulb and burns with a muchbrighter light. Also, because of the decreasedtungsten deposits on the surface, it does notbegin to dim as it nears the end of its life.

Iodine

First isolated in 1811 from ashes of seaweed,iodine has a name derived from the Greek wordmeaning “violet-colored”—a reference to the factit forms dark purple crystals. During the 1800s,iodine was obtained commercially from mines inChile, but during the twentieth century wells ofbrine in Japan, Oklahoma, and Michigan haveproven a better source.

Among the best-known properties of iodineis its importance in the human diet. The thyroidgland produces a growth-regulating hormonethat contains iodine, and lack of iodine can causea goiter, a swelling around the neck. Table saltdoes not naturally contain iodine; however,sodium chloride sold in stores usually containsabout 0.01% sodium iodide, added by the manu-facturer.

Iodine was once used in the development ofphotography: during the early days of photo-graphic technology, the daguerreotype processused silver plates sensitized with iodine vapors.


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Halogens






chlorine anion. Rather, an anion involving

a single element is named by adding the

suffix -ide to the name of the original ele-

ment—in this case, “chloride.” Other rules

apply for more complex anions.






CHEMICAL SYMBOL: A one-or two-


element.


ment that exists as molecules composed

of two atoms. All of the halogens are

diatomic.


trical current to cause a chemical reaction.

HALF-LIFE: The length of time it takes

a substance to diminish to one-half its

initial amount.


table of elements, including fluorine, chlo-

rine, bromine, iodine, and astatine. The

halogens are diatomic, and tend to form

salts; hence their name, which comes from

two Greek terms meaning “salt-forming.”

ION: An atom that has lost or gained

one or more electrons, and thus has a net

electric charge.



between ions with opposite electrical

charges.






unstable—that is, radioactive.

PERIODIC TABLE OF ELEMENTS:

A chart that shows the elements arranged

in order of atomic number. Vertical

columns within the periodic table indicate

groups or “families” of elements with sim-

ilar chemical characteristics.

POLYMER: A large molecule contain-

ing many small units that hook together.






“rot”; instead, radioactive isotopes con-

tinue to emit particles, changing into iso-

topes of other elements, until they become

stable.

SALT: A compound formed by the

reaction of an acid with a base. Salts are

usually formed by the joining of a metal

and a nonmetal.



atom, and are involved in chemical bond-

ing. The halogens all have seven valence

electrons.

KEY TERMS

set_vol1_sec6 9/13/01 11:51 AM

Halogens Iodine compounds are used today in chemical

analysis and in synthesis of organic compounds.

Astatine

Just as fluorine has the distinction of being the

most reactive, astatine is the rarest of all the ele-

ments. Long after its existence was predicted,

chemists still had no luck finding it in nature,

and it was only created in 1940 by bombarding

bismuth with alpha particles (positively charged

helium nuclei). The newly isolated element was

given a Greek name meaning “unstable.”

Indeed, none of astatine’s 20 known isotopes

is stable, and the longest-lived has a half-life of

only 8.3 hours. This has only added to the diffi-

culties involved in learning about this strange

element, and therefore it is difficult to say what

applications, if any, astatine may have. The most

promising area involves the use of astatine to

treat a condition known as hyperthyroidism,

related to an overly active thyroid gland.

WHERE TO LEARN MORE

“The Chemistry of the Halogens.” Purdue UniversityDepartment of Chemistry (Web site). <http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/group7.html> (May 20, 2001).

“Halogens.” Chemical Elements.com (Web site).<http://www.chemicalelements.com/groups/halogens.html> (May 20, 2001).

“Halogens.” Corrosion Source (Web site).<http://www.corrosionsource.com/handbook/periodic/halogens.htm> (May 20, 2001).

“Halogens” (Web site).<http://registrar.ies.ncsu.edu/ol_2000/module6/halogen/halogen/htm> (May 20, 2001).

“The Halogens” (Web site).<http://www.nidlink.com/~jfromm/elements/halogen.htm> (May 20, 2001).

Knapp, Brian J. Chlorine, Fluorine, Bromine, and Iodine.Henley-on-Thames, England: Atlantic Europe, 1996.



“Visual Elements: Group VII—The Halogens” (Web site).<http://www.chemsoc.org/viselements/pages/data/intro_groupvii_data.html> (May 20, 2001).


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NOBLE GASESNoble Gases

CONCEPTAlong the extreme right-hand column of theperiodic table of elements is a group known asthe noble gases: helium, neon, argon, krypton,xenon, and radon. Also known as the rare gases,they once were called inert gases, because scien-tists believed them incapable of reacting withother elements. Rare though they are, these gasesare a part of everyday life, as evidenced by thehelium in balloons, the neon in signs—and theharmful radon in some American homes.

HOW IT WORKS

Defining the Noble Gases

The periodic table of elements is ordered by thenumber of protons in the nucleus of an atom fora given element (the atomic number), yet thechart is also arranged in such a way that elementswith similar characteristics are grouped together.Such is the case with Group 8, which is some-times called Group 18, a collection of non-metalsknown as the noble gases. The six noble gases are helium (He), neon (Ne), argon (Ar), kryp-ton (Kr), xenon (Xe), and radon (Rn). Theiratomic numbers are, respectively, 2, 10, 18, 36, 54,and 86.

Several characteristics, aside from theirplacement on the periodic table, define the noblegases. Obviously, all are gases, meaning that theyonly form liquids or solids at extremely low tem-peratures—temperatures that, on Earth at least,are usually only achieved in a laboratory. Theyare colorless, odorless, and tasteless, as well asmonatomic—meaning that they exist as individ-ual atoms, rather than in molecules. (By contrast,

atoms of oxygen—another gas, though not

among this group—usually combine to form a

molecule, O2.)

Low Reactivity

There is a reason why noble gas atoms tend not

to combine: one of the defining characteristics of

the noble gas “family” is their lack of chemical

reactivity. Rather than reacting to, or bonding

with, other elements, the noble gases tend to

remain apart—hence the name “noble,” implying

someone or something that is set apart from the

crowd, as it were. Due to their apparent lack of

reactivity, the noble gases—also known as the

rare gases—were once known as the inert gases.

Indeed, helium, neon, and argon have not

been found to combine with other elements to

form compounds. However, in 1962 English

chemist Neil Bartlett (1932-) succeeded in

preparing a compound of xenon with platinum

and fluorine (XePtF6), thus overturning the idea

that the noble gases were entirely “inert.” Since

that time, numerous compounds of xenon with

other elements, most notably oxygen and fluo-

rine, have been developed. Fluorine has also been

used to form simple compounds with krypton

and radon.

Nonetheless, low reactivity—instead of no

reactivity, as had formerly been thought—char-

acterizes the rare gases. One of the factors gov-

erning the reactivity of an element is its electron

configuration, and the electrons of the noble

gases are arranged in such a way as to discourage

bonding with other elements.

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NobleGases



Isolation of the Noble Gases

HELIUM. Helium is an unusual elementin many respects—not least because it is the onlyelement to have first been identified in the SolarSystem before it was discovered on Earth. This issignificant, because the elements on Earth are thesame as those found in space: thus, it is morethan just an attempt at sounding poetic when sci-entists say that humans, as well as the worldaround them, are made from “the stuff of stars.”

In 1868, a French astronomer named PierreJanssen (1824-1907) was in India to observe atotal solar eclipse. To aid him in his observations,he used a spectroscope, an instrument for ana-lyzing the spectrum of light emitted by an object.What Janssen’s spectroscope showed was surpris-ing: a yellow line in the spectrum, never seenbefore, which seemed to indicate the presence ofa previously undiscovered element. Janssencalled it “helium” after the Greek god Helios,or Apollo, whom the ancients associated with the Sun.

Janssen shared his findings with Englishastronomer Sir Joseph Lockyer (1836-1920), who

had a worldwide reputation for his work in ana-lyzing light waves. Lockyer, too, believed thatwhat Janssen had seen was a new element, and afew months later, he observed the same unusualspectral lines. At that time, the spectroscope wasstill a new invention, and many members of theworldwide scientific community doubted its use-fulness, and therefore, in spite of Lockyer’s repu-tation, they questioned the existence of this“new” element. Yet during their lifetimes, Janssenand Lockyer were proven correct.

NEON, ARGON, KRYPTON,

AND XENON. They had to wait a quartercentury, however. In 1893, English chemist SirWilliam Ramsay (1852-1916) became intriguedby the presence of a mysterious gas bubble leftover when nitrogen from the atmosphere wascombined with oxygen. This was a phenomenonthat had also been noted by English physicistHenry Cavendish (1731-1810) more than a cen-tury before, but Cavendish could offer no expla-nation. Ramsay, on the other hand, had the ben-efit of observations made by English physicistJohn William Strutt, Lord Rayleigh (1842-1919).

Up to that time, scientists believed that airconsisted only of oxygen, carbon dioxide, andwater vapor. However, Rayleigh had noticed thatwhen nitrogen was extracted from air after a

A SCIENTIST USING THE POTASSIUM-ARGON DATING PROCESS TO DETERMINE THE AGE OF MATERIALS. (Dean

Conger/Corbis. Reproduced by permission.)

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NobleGases

process of removing those other components, ithad a slightly higher density than nitrogen pre-pared from a chemical reaction. In light of hisown observations, Ramsay concluded thatwhereas nitrogen obtained from chemical reac-tions was pure, the nitrogen extracted from aircontained trace amounts of an unknown gas.

Ramsay was wrong in only one respect: hid-den with the nitrogen was not one gas, but five.In order to isolate these gases, Ramsay andRayleigh subjected air to a combination of highpressure and low temperature, allowing the vari-ous gases to boil off at different temperatures.One of the gases was helium—the first confirma-tion that the element existed on Earth—but theother four gases were previously unknown. TheGreek roots of the names given to the four gasesreflected scientists’ wonder at discovering thesehard-to-find elements: neos (new), argos (inac-tive), kryptos (hidden), and xenon (stranger).

RADON. Inspired by the studies of Pol-ish-French physicist and chemist Marie Curie(1867-1934) regarding the element radium andthe phenomenon of radioactivity (she discoveredthe element, and coined the latter term), Germanphysicist Friedrich Dorn (1848-1916) becamefascinated with radium. Studying the element, hediscovered that it emitted a radioactive gas,which he dubbed “radium emanation.” Eventual-ly, however, he realized that what was being pro-duced was a new element. This was the first clearproof that one element could become anotherthrough the process of radioactive decay.

Ramsay, who along with Rayleigh hadreceived the Nobel Prize in 1904 for his work onthe noble gases, was able to map the new ele-ment’s spectral lines and determine its densityand atomic mass. A few years later, in 1918,another scientist named C. Schmidt gave it thename “radon.” Due to its behavior and the con-figuration of its electrons, chemists classifiedradon among what they continued to call the“inert gases” for another half-century—untilBartlett’s preparation of xenon compounds in1962.

Presence of the Rare Gaseson Earth

IN THE ATMOSPHERE. Thoughthe rare gases are found in minerals and mete-orites on Earth, their greatest presence is in theplanet’s atmosphere. It is believed that they were


released into the air long ago as a by-product ofdecay on the part of radioactive materials in the Earth’s crust. Within the atmosphere, argon isthe most “abundant”—in comparative terms,given the fact that the “rare gases” are, by defini-tion, rare.

Nitrogen makes up about 78% of Earth’satmosphere and oxygen 21%, meaning that thesetwo elements constitute fully 99% of the airabove the Earth. Argon ranks a distant third, with0.93%. The remaining 0.07% is made up onwater vapor, carbon dioxide, ozone (O3), andtraces of the noble gases. These are present insuch small quantities that the figures for them arenot typically presented as percentages, but ratherin terms of parts per million (ppm). The con-centrations of neon, helium, krypton, and xenonin the atmosphere are 18, 5, 1, and 0.09 ppmrespectively.

IN THE SOIL. Radon in the atmos-phere is virtually negligible, which is a fortunatething, in light of its radioactive qualities. FewAmericans, in fact, even knew of its existenceuntil 1988, when the United States Environmen-tal Protection Agency (EPA) released a reportestimating that some ten million Americanhomes had potentially harmful radon levels. This

LAS VEGAS HAS NUMEROUS EXAMPLES OF SIGNS

ILLUMINATED BY NEON. (Dave Bartruff/Corbis. Reproduced by

permission.)

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set off a scare, and during the late 1980s and1990s, sales of home radon detectors boomed.Meanwhile, the federal government increasedconcerns with additional reports, advising peopleto seal their basements and ventilate their homesif radon exceeded certain levels.

A number of scientists have disputed thegovernment’s claims, yet some regions of theUnited States appear to be at relatively high riskdue to the presence of radon in the soil. The ele-ment seems to be most plentiful in soils contain-ing high concentrations of uranium. If radon ispresent in a home that has been weather-sealedto improve the efficiency of heating and coolingsystems, it is indeed potentially dangerous to theresidents.

Chinese scientists in the 1960s made aninteresting discovery regarding radon and itsapplication to seismography, or the area of theearth sciences devoted to studying and predictingearthquakes. Radon levels in groundwater, theChinese reports showed, rise considerably justbefore an earthquake. Since then, the Chinesehave monitored radon concentrations in water,and used this data to predict earthquakes.

EXTRACTING RARE GASES.

Radon, in fact, is not the only rare gas that can be

obtained as the result of radioactive decay: in1903, Ramsay and British chemist FrederickSoddy (1877-1956) showed that the breakdownof either uranium or radium results in the pro-duction of helium atoms (beta particles). A fewyears later, English physicist Ernest Rutherford(1871-1937) demonstrated that radiation carry-ing a positive electrical charge (alpha rays) wasactually a stream of helium atoms stripped of anelectron.

Many of the noble gases are extracted by liquefying air—that is, by reducing it to temper-atures at which it assumes the properties of a liquid rather than a gas. By controlling tempera-tures in the liquefied air, it is possible to reach theboiling point for a particular noble gas and there-by extract it, much as was done when these gaseswere first isolated in the 1890s.

THE UNIQUE SITUATION OF

HELIUM. Helium is remarkable, in that itonly liquefies at a temperature of -457.6°F (-272°C), just above absolute zero. Absolute zerois the temperature at which the motion of atomsor molecules comes to a virtual stop, but themotion of helium atoms never completely ceases.In order to liquefy it, in fact, even at those low

A RESEARCHER WORKS WITH LIQUID HELIUM AS PART OF AN EXPERIMENT ON COSMIC BACKGROUND RADIATION.CLOSE TO ABSOLUTE ZERO, HELIUM TRANSFORMS INTO A HIGHLY UNUSUAL LIQUID THAT HAS NO MEASURABLE

RESISTANCE TO FLOW. (Roger Ressmeyer/Corbis. Reproduced by permission.)

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NobleGases

temperatures, it must be subjected to pressuresmany times that exerted by Earth’s atmosphere.

Given these facts, it is difficult to extracthelium from air. More often, it is obtained fromnatural gas wells, where it is present in relativelylarge concentrations—between 1% and 7% ofthe natural gas. The majority of the Earth’s heli-um supply belongs to the United States, wherethe greatest abundance of helium-supplyingwells are in Texas, Oklahoma, and Kansas. Dur-ing World War II, the United States took advan-tage of this supply of relatively inexpensive heli-um to provide buoyancy for a fleet of airshipsused for reconnaissance.

There is one place with an abundant supplyof helium, but there are no plans for a miningexpedition any time soon. That place is the Sun,where the nuclear fusion of hydrogen atoms cre-ates helium. Indeed, helium seems to be the mostplentiful element of all, after hydrogen, constitut-ing 23% of the total mass of the universe. Why,then, is it so difficult to obtain on Earth? Mostlikely because it is so light in comparison to air; itsimply floats off into space.

Applications for the NobleGases

RADON, ARGON, KRYPTON,

AND XENON. Though radon is known pri-marily for the hazards it poses to human life andwell-being, it has useful applications. As notedabove, its presence in groundwater appears toprovide a possible means of predicting earth-quakes. In addition, it is used for detecting leaks,measuring flow rates, and inspecting metalwelds.

One interesting use of argon and, in particu-lar, the stable isotope argon-40, is in dating tech-niques used by geologists, paleontologists, andother scientists studying the distant past. Whenvolcanic rocks are subjected to extremely hightemperatures, they release argon, and as the rockscool, argon-40 accumulates. Because argon-40 isformed by the radioactive decay of a potassiumisotope, potassium-40, the amount of argon-40that forms is proportional to the rate of decay forpotassium-40. The latter has a half-life of 1.3 bil-lion years, meaning that it takes 1.3 billion yearsfor half the potassium-40 originally present to beconverted to argon-40. Using argon-40, paleon-tologists have been able to estimate the age of

volcanic layers above and below fossil and arti-fact remains in east Africa.

Krypton has a number of specialized appli-cations—for instance, it is mixed with argon andused in the manufacture of windows with a highlevel of thermal efficiency. Used in lasers, it isoften mixed with a halogen such as fluorine. Inaddition, it is also sometimes used in halogensealed-beam headlights. Many fans of Superman,no doubt, were disappointed at some point intheir lives to discover that there is no such thingas “kryptonite,” the fictional element that causedthe Man of Steel to lose his legendary strength.Yet krypton—the real thing—has applicationsthat are literally out of this world. In the devel-opment of fuel for space exploration, krypton isin competition with its sister element, xenon.Xenon offers better performance, but costs aboutten times more to produce; thus krypton hasbecome more attractive as a fuel for space flight.

In addition to its potential as a space fuel,xenon is used in arc lamps for motion-picturefilm projection, in high-pressure ultraviolet radi-ation lamps, and in specialized flashbulbs usedby photographers. One particular isotope ofxenon is utilized for tracing the movement ofsands along a coastline. Xenon is also applied inhigh-energy physics for detecting nuclear radia-tion in bubble chambers. Furthermore, neurosci-entists are experimenting with the use of xenonin diagnostic procedures to clarify x-ray imagesof the human brain.

NEON. Neon, of course, is best-knownfor its application in neon signs, which producean eye-catching glow when lit up at night. Frenchchemist Georges Claude (1870-1960), intriguedby Ramsay’s discovery of neon, conducted exper-iments that led to the development of the neonlight in 1910. That first neon light was simply aglass tube filled with neon gas, which glowed abright red when charged with electricity.

Claude eventually discovered that mixingother gases with neon produced different colorsof light. He also experimented with variations inthe shapes of glass tubes to create letters and pic-tures. By the 1920s, neon light had come intovogue, and it is still popular today. Modern neonlamps are typically made of plastic rather thanglass, and the range of colors is much greaterthan in Claude’s day: not only the gas filling, butthe coating inside the tube, is varied, resulting ina variety of colors from across the spectrum.


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NobleGases

Though the neon sign is its best-knownapplication, neon is used for many other things.Neon glow lamps are often used to indicateon/off settings on electronic instrument panels,and lightweight neon lamps are found onmachines ranging from computers to voltageregulators. In fact, the first practical color televi-sion, produced in 1928, used a neon tube to pro-duce the red color in the receiver. Green camefrom mercury, but the blue light in that earlycolor TV came from another noble gas, helium.

HELIUM. Helium, of course, is widelyknown for its use in balloons—both for large air-ships and for the balloons that have provided joyand fun to many a small child. Though helium ismuch more expensive than hydrogen as a meansof providing buoyancy to airships, hydrogen isextremely flammable, and after the infamousexplosion of the airship Hindenburg in 1937,helium became the preferred medium for air-ships. As noted earlier, the United States militarymade extensive use of helium-filled airships dur-ing the World War II.

The use of helium for buoyancy is one of themost prominent applications of this noble gas,but far from the only one. In fact, not only havepeople used helium to go up in balloons, butdivers use helium for going down beneath thesurface of the ocean. In that situation, of course,helium is not used for providing buoyancy, but asa means of protecting against the diving-relatedcondition known as “the bends,” which occurswhen nitrogen in the blood bubbles as the diverrises to the surface. Helium is mixed with oxygenin diver’s air tanks because it does not dissolve inthe blood as easily as nitrogen.

Among the most fascinating applications ofhelium relate to its extraordinarily low freezingpoint. Helium has played a significant role in thelow-temperature science known as cryogenics,and has found application in research concerningsuperconductivity: the use of very low tempera-tures to develop materials that conduct electricalpower with vastly greater efficiency than ordi-nary conductors. Close to absolute zero, heliumtransforms into a highly unusual liquid unlikeany known substance, in that it has no measura-ble resistance to flow. This means that it couldcarry an electrical current hundreds of timesmore efficiently than a copper wire.

WHERE TO LEARN MORE

“The Chemistry of the Rare Gases” (Web site).<http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/raregas.html> (May 13, 2001).

“Homework: Science: Chemistry: Gases” Channelone.com(Web site). <http://www.channelone.com/fasttrack/science/chemistry/gases.html> (May 12, 2001).

Knapp, Brian J.; David Woodroffe; David A. Hardy. Ele-ments. Danbury, CT: Grolier Educational, 2000.

Mebane, Robert C. and Thomas R. Rybolt. Air and OtherGases. Illustrations by Anni Matsick. New York:Twenty-First Century Books, 1995.

“Noble Gases”Xrefer.com (Web site). <http://www.xrefer.com/entry/643259> (May 13, 2001).

Rare Gases. Praxair (Web site). <http://www.praxair.com/Praxair.nsf/X1/gase_rarega?openDocument>(May 13, 2001).

Stwertka, Albert. Superconductors: The Irresistible Future.New York: F. Watts, 1991.

Taylor, Ron. Facts on Radon and Asbestos. Illustrated byIan Moores. New York: F. Watts, 1990.


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C ARBONCarbon

CONCEPTThe phrase “carbon-based life forms,” often usedin science-fiction books and movies by aliens todescribe the creatures of Earth, is something of acliché. It is also a redundancy when applied tocreatures on Earth, the only planet known tosupport life: all living things contain carbon.Carbon is also in plenty of things that were onceliving, which makes it useful for dating theremains of past settlements on Earth. Of evengreater usefulness is petroleum, a substance con-taining carbon-based forms that died long ago,became fossilized, and ultimately changed chem-ically into fuels. Then again, not all materialscontaining carbon were once living creatures; yetbecause carbon is a common denominator to allliving things on Earth, the branch of studyknown as organic chemistry is devoted to thestudy of compounds containing carbon. Amongthe most important organic compounds are themany carboxylic acids that are vital to life, butcarbon is also present in numerous importantinorganic compounds—most notably carbondioxide and carbon monoxide.

HOW IT WORKS

The Basics of Carbon

Carbon’s name comes from the Latin word carbo,or charcoal—which, indeed, is almost pure car-bon. Its chemical symbol is C, and it has an atom-ic number of 6, meaning that there are six protonsin its nucleus. Its two stable isotopes are 12C, whichconstitutes 98.9% of all carbon found in nature,and 13C, which accounts for the other 1.1%.

The mass of the 12C atom is the basis for theatomic mass unit (amu), by which mass figures

for all other elements are measured: the amu isdefined as exactly 1/12 the mass of a single 12Catom. The difference in mass between 12C and13C, which is heavier because of its extra neutron,account for the fact that the atomic mass of car-bon is 12.01 amu: were it not for the small quan-tities of 13C present in a sample of carbon, themass would be exactly 12.00 amu.

WHERE CARBON IS FOUND.

Carbon makes up only a small portion of theknown elemental mass in Earth’s crust, oceans,and atmosphere—just 0.08%, or 1/1250 of thewhole—yet it is the fourteenth most abundantelement on the planet. In the human body, car-bon is second only to oxygen in abundance, andaccounts for 18% of the body’s mass. Thus if aperson weighs 100 lb (45.3 kg), she is carryingaround 18 lb (8.2 kg) of carbon—the same mate-rial from which diamonds are made!

Present in the inorganic rocks of the groundand in the living creatures above it, carbon iseverywhere. Combined with other elements, itforms carbonates, most notably calcium carbon-ate (CaCO3), which appears in the form of lime-stone, marble, and chalk. In combination withhydrogen, it creates hydrocarbons, present indeposits of fossil fuels: natural gas, petroleum,and coal. In the environment, carbon—in theform of carbon dioxide (CO2)—is taken in byplants, which undergo the process of photosyn-thesis and release oxygen to animals. Animalsbreathe in oxygen and release carbon dioxide tothe atmosphere.

Carbon and Bonding

Located in Group 4 of the periodic table of ele-ments (Group 14 in the IUPAC system), carbonhas a valence electron configuration of 2s22p2;

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Carbon


likewise, all the members of Group 4—some-times known as the “carbon family”—have con-figurations of ns2np2, where n is the number ofthe period or row that the element occupies onthe table.

There are two elements noted for their abili-ty to form long strings of atoms and seeminglyendless varieties of molecules: one is carbon, andthe other is silicon, directly below it on the peri-odic table. Silicon, found in virtually all types ofrocks except the calcium carbonates (mentionedabove), is to the inorganic world what carbon isto the organic. Yet silicon atoms are about oneand a half times as large as those of carbon; thusnot even silicon can compete with carbon’s abili-ty to form a seemingly limitless array of mole-cules in various shapes and sizes, and having var-ious chemical properties.

BASICS OF CHEMICAL BOND-

ING. Carbon is further distinguished by itshigh value of electronegativity, the relative abilityof an atom to attract valence electrons. Elec-tronegativity increases with an increase in groupnumber, and decreases with an increase in periodnumber. In other words, the elements with thehighest electronegativity values lie in the upperright-hand corner of the periodic table.

Actually, the previous statement requiresone significant qualification: the extreme right-hand side of the periodic table is occupied by ele-ments with negligible electronegativity values.These are the noble gases, which have eightvalence electrons each. Eight, as it turns out, isthe “magic number” for chemical bonding: mostelements follow what is known as the octet rule,meaning that when one element bonds to anoth-er, the two atoms have eight valence electrons.

If the two atoms have an electric charge andthus are ions, they form strong ionic bonds. Ionicbonding occurs when a metal bonds with a non-metal. The other principal type of bond is a cova-lent bond, in which two uncharged atoms shareeight valence electrons. If the electronegativityvalues of the two elements involved are equal,they share the electrons equally; but if one ele-ment has a higher electronegativity value, theelectrons will be more drawn to that element.

ELECTRONEGATIVITY OF CAR-

BON. To return to electronegativity and theperiodic table, let us ignore the noble gases,which are the chemical equivalent of snobs.

(Hence the term “noble,” meaning that they areset apart.) To the left of the noble gases are thehalogens, a wildly gregarious bunch—none moreso than the element that occupies the top of thecolumn, fluorine. With an electronegativity valueof 4.0, fluorine is the most reactive of all ele-ments, and the only one capable of bonding evento a few of the noble gases.

So why is fluorine—capable of formingmultitudinous bonds—not as chemically signifi-cant as carbon? There are a number of answers,but a simple one is this: because fluorine is toostrong, and tends to “overwhelm” other elements,precluding the possibility of forming long chains,it is less chemically significant than carbon. Car-bon, on the other hand, has an electronegativityvalue of 2.5, which places it well behind fluorine.Yet it is still at sixth place (in a tie with iodine andsulfur) on the periodic table, behind only fluo-rine; oxygen (3.5); nitrogen and chlorine (3.0);and bromine (2.8). In addition, with four valenceelectrons, carbon is ideally suited to find otherelements (or other carbon atoms) for formingcovalent bonds according to the octet rule.

MULTIPLE BONDS. Normally, anelement does not necessarily have the ability tobond with as many other elements as it hasvalence electrons, but carbon—with its fourvalence electrons—happens to be tetravalent, orcapable of bonding to four other atoms at once.Additionally, carbon is capable of forming notonly a single bond, but also a double bond, oreven a triple bond, with other elements.

Suppose a carbon atom bonds to two oxygenatoms to form carbon dioxide. Let us imaginethese three atoms side by side, with the oxygen inthe middle. (This, in fact, is how these bonds aredepicted in the Couper and Lewis systems ofchemical symbolism, discussed in the ChemicalBonding essay.) We know that the carbon hasfour valence electrons, that the oxygens have six,and that the goal is for each atom to have eightvalence electrons—some of which it will sharecovalently.

Two of the valence electrons from the car-bon bond with two valence electrons each fromthe oxygen atoms on either side. This means thatthe carbon is doubly bonded to each of the oxy-gen atoms. Therefore, the two oxygens each havefour other unbonded valence electrons, whichmight bond to another atom. It is theoreticallypossible, also, for the carbon to form a triple

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Carbon


bond with one of the oxygens by sharing three ofits valence electrons. It would then have one elec-tron free to share with the other oxygen.


Organic Chemistry

We have stated that carbon forms tetravalentbonds, and makes multiple bonds with a singleatom. In addition, we have mentioned the factthat carbon forms long chains of atoms and vari-eties of shapes. But how does it do these things,and why? These are good questions, but not oneswe will attempt to answer here. In fact, an entirebranch of chemistry is devoted to answeringthese theoretical questions, as well as to deter-mining solutions to a host of other, more practi-cal problems.

Organic chemistry is the study of carbon, itscompounds, and their properties. (There are car-bon-containing compounds that are not consid-ered organic, however. Among these are oxidessuch as carbon dioxide and monoxide; as well ascarbonates, most notably calcium carbonate.) Atone time, chemists thought that “organic” wassynonymous with “living,” and even as recently asthe early nineteenth century, they believed thatorganic substances contained a supernatural “lifeforce.” Then, in 1828, German chemist FriedrichWöhler (1800-1882) cracked the code that dis-tinguished the living from the nonliving, and theorganic from the inorganic.

Wöhler took a sample of ammoniumcyanate (NH4OCN), and by heating it, convertedit into urea (H2N-CO-NH2), a waste product inthe urine of mammals. In other words, he hadturned an inorganic material into a organic one,and he did so, as he observed, “without benefit ofa kidney, a bladder, or a dog.” It was almost asthough he had created life. In fact, what Wöhlerhad glimpsed—and what other scientists whofollowed came to understand, was this: what sep-arates the organic from the inorganic is the man-ner in which the carbon chains are arranged.

Ammonium cyanate and urea have exactlythe same numbers and proportions of atoms, yetthey are different compounds. They are thus iso-mers: substances which have the same formula,but are different chemically. In urea, the carbonforms an organic chain, and in ammonium

cyanate, it does not. Thus, to reduce the specificsof organic chemistry even further, it can be saidthat this area of the field constitutes the study ofcarbon chains, and ways to rearrange them inorder to create new substances.

Rubber, vitamins, cloth, and paper are allorganically based compounds we encounter inour daily lives. In each case, the material comesfrom something that once was living, but whattruly makes these substance organic in nature isthe common denominator of carbon, as well asthe specific arrangements of the atoms. We haveorganic chemistry to thank for any number ofthings: aspirins and all manner of other drugs;preservatives that keep food from spoiling;perfumes and toiletries; dyes and flavorings, andso on.

Allotropes of Carbon

GRAPHITE. Carbon has severalallotropes—different versions of the same ele-ment, distinguished by molecular structure. Thefirst of these is graphite, a soft material with anunusual crystalline structure. Graphite is essen-tially a series of one-atom-thick sheets of carbon,bonded together in a hexagonal pattern, but withonly very weak attractions between adjacentsheets. A piece of graphite is thus like a big, thickstack of carbon paper: on the one hand, the stackis heavy, but the sheets are likely to slide againstone another.

Actually, people born after about 1980 mayhave little experience with carbon paper, whichwas gradually phased out as photocopiersbecame cheaper and more readily available.Today, carbon paper is most often encounteredwhen signing a credit-card receipt: the signaturegoes through the graphite-based backing of thereceipt, onto a customer copy.

In such a situation, one might notice that thecopied image of the signature looks as though itwere signed in pencil. This is not surprising, con-sidering that pencil “lead” is, in fact, a mixture ofgraphite, clay, and wax. In ancient times, peopledid indeed use lead—the heaviest member ofGroup 4, the “carbon family”—for writing,because it left gray marks on a surface. Lead, ofcourse, is poisonous, and is not used today inpencils or in most applications that wouldinvolve prolonged exposure of humans to theelement. Nonetheless, people still use the word“lead” in reference to pencils, much as they still

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Carbon refer to a galvanized steel roof with a zinc coatingas a “tin roof.”

In graphite the atoms of each “sheet” aretightly bonded in a hexagonal, or six-sided, pat-tern, but the attractions between the sheets arenot very strong. This makes it highly useful as alubricant for locks, where oil would tend to bemessy. A good conductor of electricity, graphiteis also utilized for making high-temperature elec-trolysis cells. In addition, the fact that graphiteresists temperatures of up to about 6,332°F(3,500°C) makes it useful in electric motors andgenerators.

DIAMOND. The second allotrope of car-bon is also crystalline in structure. This is dia-mond, most familiar in the form of jewelry, butin fact widely applied for a number of other pur-poses. According to the Moh scale, which meas-ures the hardness of minerals, diamond is a 10—in other words, the hardest type of material. It isused for making drills that bore through solidrock; likewise, small diamonds are used in den-tists’ drills for boring through the ultra-hardenamel on teeth.

Neither diamonds nor graphite are, in thestrictest sense of the term, formed of mole-cules. Their arrangement is definite, as with amolecule, but their size is not: they simply formrepeating patterns that seem to stretch on for-ever. Whereas graphite is in the form ofsheets, a diamond is basically a huge “molecule”composed of carbon atoms strung together bycovalent bonds. The size of this “molecule” corre-sponds to the size of the diamond: a diamond of1 carat, for instance, contains about 1022

(10,000,000,000,000,000,000,000 or 10 billionbillion) carbon atoms.

The diamonds used in industry look quitedifferent from the ones that appear in jewelry.Industrial diamonds are small, dark, and cloudyin appearance, and though they have the samechemical properties as gem-quality diamonds,they are cut with functionality (rather than beau-ty) in mind. A diamond is hard, but brittle: inother words, it can be broken, but it is very diffi-cult to scratch or cut a diamond—except withanother diamond.

The cutting of fine diamonds for jewelry isan art, exemplified in the alluring qualities ofsuch famous gems as the jewels in the BritishCrown or the infamous Hope Diamond in Wash-ington, D.C.’s Smithsonian Institution. Such dia-

monds—as well as the diamonds on an engage-ment ring—are cut to refract or bend light rays,and to disperse the colors of visible light.

B U C K M I N S T E R F U L L E R E N E .

Until 1985, carbon was believed to exist in onlytwo crystalline forms, graphite and diamond. Inthat year, however, chemists at Rice University inHouston, Texas, and at the University of Sussexin England, discovered a third variety of car-bon—and later jointly received a Nobel Prize fortheir work. This “new” carbon molecule com-posed of 60 bonded atoms in the shape of what iscalled a “hollow truncated icosahedron.” In plainlanguage, this is rather like a soccer ball, withinterlocking pentagons and hexagons. However,because the surface of each geometric shape isflat, the “ball” itself is not a perfect sphere.Rather, it describes the shape of a geodesic dome,a design created by American engineer andphilosopher R. Buckminster Fuller (1895-1983).

There are other varieties of buckminster-fullerene molecules, known as fullerenes. Howev-er, the 60-atom shape, designated as 60C, is themost common of all fullerenes, the result of con-densing carbon slowly at high temperatures.Fullerenes potentially have a number of applica-tions, particularly because they exhibit a wholerange of electrical properties: some are insula-tors, while some are conductors, semiconductors,and even superconductors. Due to the high costof producing fullerenes artificially, however, theways in which they are applied remain ratherlimited.

AMORPHOUS CARBON. There isa fourth way in which carbon appears, distin-guished from the other three in that it is amor-phous, as opposed to crystalline, in structure. Anexample of amorphous carbon is carbon black,obtained from smoky flames and used in ink, orfor blacking rubber tires.

Though it retains some of the microscopicstructures of the plant cells in the wood fromwhich it is made, charcoal—wood or other plantmaterial that has been heated without enough airpresent to make it burn—is mostly amorphouscarbon. One form of charcoal is activated char-coal, in which steam is used to remove the stickyproducts of wood decomposition. What remainsare porous grains of pure carbon with enormousmicroscopic surface areas. These are used inwater purifiers and gas masks.


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Carbon

Coal and coke are particularly significantvarieties of amorphous carbon. Formed by thedecay of fossils, coal was one of the first “fossilfuels” (for example, petroleum) used to provideheat and power for industrial societies. Indeed,when the words “industrial revolution” are men-tioned, many people picture tall black smoke-stacks belching smoke from coal fires. Fortunate-ly— from an environmental standpoint—coal isnot nearly so widely used today, and when it is(as for instance in electric power plants), themethods for burning it are much more efficientthan those applied in the nineteenth century.

Actually, much of what those smokestacks ofyesteryear burned was coke, a refined version ofcoal that contains almost pure carbon. Producedby heating soft coal in the absence of air, coke hasa much greater heat value than coal, and is stillwidely used as a reducing agent in the produc-tion of steel and other alloys.

Carbon Dioxide

Carbon forms many millions of compounds,some families of which will be discussed below.Two others, formed by the bonding of carbonatoms with oxygen atoms, are of particular sig-nificance. In carbon dioxide, a single carbon

joins with two oxygens to produce a gas essentialto plant life. In carbon monoxide (CO), a singleoxygen joins the carbon, creating a toxic—butnonetheless important—compound.

The first gas to be distinguished from ordi-nary air, carbon dioxide is an essential compo-nent in the natural balance between plant andanimal life. Animals, including humans, producecarbon dioxide by breathing, and humans fur-ther produce it by burning wood and other fuels.Plants use carbon dioxide when they store ener-gy in the form of food, and they release oxygen tobe used by animals.

DISCOVERY. Flemish chemist andphysicist Johannes van Helmont (1579-1644)discovered in 1630 that air was not, as had beenthought up to that time, a single element: it con-tained a second substance, produced in the burn-ing of wood, which he called “gas sylvestre.” Thushe is recognized as the first scientist to note theexistence of carbon dioxide.

More than a century later, in 1756, Scottishchemist Joseph Black (1728-1799) showed thatcarbon dioxide—which he called “fixed air”—combines with other chemicals to form com-pounds. This and other determinations Blackmade concerning carbon dioxide led to enor-


A GEODESIC DOME IN MONTREAL, QUEBEC, CANADA. SUCH DOMES ARE NAMED AFTER THEIR DESIGNER, R. BUCK-MINSTER FULLER, AND EVENTUALLY PROVIDED THE NAME FOR THE CARBON MOLECULES KNOWN AS BUCKMINSTER-FULLERENES. (Lee Snider/Corbis. Reproduced by permission.)

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Carbon

mous progress in the discovery of gases by vari-ous chemists of the late eighteenth century.

By that time, chemists had begun to arrive ata greater degree of understanding with regard tothe relationship between plant life and carbondioxide. Up until that time, it had been believedthat plants purify the air by day, and poison it atnight. Carbon dioxide and its role in the connec-tion between animal and plant life provided amuch more sophisticated explanation as to theways plants “breathe.”

USES. Around the same time that Blackmade his observations on carbon dioxide, Eng-lish chemist Joseph Priestley (1733-1804)became the first scientist to put the chemical touse. Dissolving it in water, he created carbonatedwater, which today is used in making soft drinks.Not only does the gas add bubbles to drinks, italso acts as a preservative.

Though the natural uses of carbon dioxideare by far the most important, the compound hasnumerous industrial and commercial applica-tions. Used in fire extinguishers, carbon dioxideis ideal for controlling electrical and oil fires,which cannot be put out with water. Heavierthan air, carbon dioxide blankets the flames andsmothers them.

In the solid form of dry ice, carbon dioxide

is used for chilling perishable food during trans-

port. It is also one of the only compounds that

experiences sublimation, or the instantaneous

transformation of a solid to a gas without passing

through an intermediate liquid state, at condi-

tions of ordinary pressure and temperature. Dry

ice has often been used in movies to generate

“mists” or “smoke” in a particular scene.

Carbon Monoxide

During the late eighteenth century, Priestley dis-

covered a carbon-oxygen compound different

from carbon dioxide: carbon monoxide. Scien-

tists had actually known of this toxic gas, released

in the incomplete combustion of wood, from the

Middle Ages onward, but Priestley was the first to

identify it scientifically.

Industry uses carbon monoxide in a number

of ways. By blowing air across very hot coke, the

result is producer gas, which, along with water

gas (made by passing hot steam over coal) is an

important fuel. Producer gas constitutes a 6:1:18

mixture of carbon monoxide, carbon dioxide,

and nitrogen, while water gas is 40% carbon

monoxide, 50% hydrogen, and 10% carbon diox-

ide and other gases.


SOFT DRINKS ARE MADE POSSIBLE BY THE USE OF CARBONATED WATER. (Sergio Dorantes/Corbis. Reproduced by permission.)

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Carbon




ular structure.

AMORPHOUS: Having no definite

structure.

CARBOHYDRATES: Naturally occur-

ring compounds of carbon, hydrogen, and

oxygen. These are primarily produced by

green plants through the process of photo-

synthesis.

CELLULAR RESPIRATION: The

process whereby nutrients from plants are

broken down in an animal’s body to create

carbon dioxide.




CRYSTALLINE: A term describing a

type of solid in which the constituent

parts have a simple and definite geo-

metric arrangement that is repeated in all

directions.

DOUBLE BOND: A form of bonding

in which two atoms share two pairs of

valence electrons. Carbon is also capable of

single bonds and triple bonds.



electrons.







charges.

ISOMERS: Substances which have the

same chemical formula, but which are dif-

ferent chemically due to differences in the












ments, which usually end up with eight

valence electrons.

ORGANIC CHEMISTRY: The study of

carbon, its compounds, and their proper-

ties. (Many carbon-containing oxides and

carbonates are not considered organic,

however.)

PHOTOSYNTHESIS: The biological

conversion of light energy (that is, electro-

magnetic energy) to chemical energy in

plants.









becomes stable.

SINGLE BOND: A form of bonding in

which two atoms share one pair of valence

electrons. Carbon is also capable of double

bonds and triple bonds.

TETRAVALENT: Capable of bonding

to four other elements.

KEY TERMS

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Carbon


Not only are producer and water gas used forfuel, they are also applied as reducing agents.Thus, when carbon monoxide is passed over hotiron oxides, the oxides are reduced to metalliciron, while the carbon monoxide is oxidized toform carbon dioxide. Carbon monoxide is alsoused in reactions with metals such as nickel, iron,and cobalt to form some types of carbonyls.

Carbon monoxide—produced by burningpetroleum in automobiles, as well as by the com-bustion of wood, coal, and other carbon-con-taining fuels—is extremely hazardous to humanhealth. It bonds with iron in hemoglobin, thesubstance in red blood cells that transports oxy-gen throughout the body, and in effect fools thebody into thinking that it is receiving oxygenatedhemoglobin, or oxyhemoglobin. Upon reachingthe cells, carbon monoxide has much less ten-dency than oxygen to break down, and thereforeit continues to circulate throughout the body.Low concentrations can cause nausea, vomiting,and other effects, while prolonged exposure tohigh concentrations can result in death.

Carbon and the Environment

Carbon is released into the atmosphere by one ofthree means: cellular respiration; the burning offossil fuels; and the eruption of volcanoes. Whenplants take in carbon dioxide from the atmos-phere, they combine this with water and manu-facture organic compounds using energy theyhave trapped from sunlight by means of photo-synthesis—the conversion of light to chemicalenergy through biological means. As a by-prod-uct of photosynthesis, plants release oxygen intothe atmosphere.

In the process of undergoing photosynthe-sis, plants produce carbohydrates, which are var-ious compounds of carbon, hydrogen, and oxy-gen essential to life. The other two fundamentalcomponents of a diet are fats and proteins, both

carbon-based as well. Animals eat the plants, oreat other animals that eat the plants, and thusincorporate the fats, proteins, and sugars (a formof carbohydrate) from the plants into their bod-ies. Cellular respiration is the process wherebythese nutrients are broken down to create carbondioxide.

Photosynthesis and cellular respiration arethus linked in what is known as the carbon cycle.Cellular respiration also releases carbon into theatmosphere through the action of decom-posers—bacteria and fungi that feed on theremains of plants and animals. The decomposersextract the energy in the chemical bonds of thedecomposing matter, thus releasing more carbondioxide into the atmosphere.

When creatures die and are buried in such away that they cannot be reached by decom-posers—for instance, at the bottom of the ocean,or beneath layers of rock—the carbon in theirbodies is eventually converted to fossil fuels,including petroleum, natural gas, and coal. Theburning of fossil fuels releases carbon (bothmonoxide and dioxide) into the atmosphere.

Because the rate of such burning hasincreased dramatically since the late nineteenthcentury, this has raised fears that carbon dioxidein the atmosphere may create a greenhouseeffect, leading to global warming. On the otherhand, volcanoes release tons of carbon into theatmosphere regardless of whether humans burnfossil fuels or not.

Radiocarbon Dating

Radiocarbon dating is used to date the age ofcharcoal, wood, and other biological materials.When an organism is alive, it incorporates a cer-tain ratio of carbon-12 in proportion to theamount of the radioisotope (that is, radioactiveisotope) carbon-14 that it receives from the

TRIPLE BOND: A form of bonding in

which two atoms share three pairs of


single bonds and double bonds.






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Carbonatmosphere. As soon as the organism dies, how-

ever, it stops incorporating new carbon, and the

ratio between carbon-12 and carbon-14 will

begin to change as the carbon-14 decays to form

nitrogen-14.

Carbon-14 has a half-life of 5,730 years,

meaning that it takes that long for half the iso-

topes in a sample to decay to nitrogen-14. There-

fore a scientist can use the ratios of carbon-12,

carbon-14, and nitrogen-14 to guess the age of an

organic sample. The problem with radiocarbon

dating, however, is that there is a good likelihood

the sample can become contaminated by addi-

tional carbon from the soil. Furthermore, it can-

not be said with certainty that the ratio of car-

bon-12 to carbon-14 in the atmosphere has been

constant throughout time.

WHERE TO LEARN MORE

Blashfield, Jean F. Carbon. Austin, TX: Raintree Steck-Vaughn, 1999.

“Carbon.” Xrefer (Web site). <http://www.xrefer.com/entry/639742> (May 30, 2001).

“Diamonds.” American Museum of Natural History (Website). <http://www.amnh.org/exhibitions/diamonds/structure.html> (May 30, 2001).

Knapp, Brian J. Carbon Chemistry. Illustrated by DavidWoodroffe. Danbury, CT: Grolier Educational, 1998.

Loudon, G. Marc. Organic Chemistry. Menlo Park, CA:Benjamin/Cummings, 1988.

“Organic Chemistry” (Web site). <http://edie.cprost.sfu.ca/~rhlogan/organic.html> (May 30, 2001).

“Organic Chemistry.” Frostburg State University Chem-istry Helper (Web site). <http://www.chemhelper.com/> (May 30, 2001).

Sparrow, Giles. Carbon. New York: Benchmark Books,1999.

Stille, Darlene. The Respiratory System. New York: Chil-dren’s Press, 1997.


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H YDROGENHydrogen

CONCEPTFirst element on the periodic table, hydrogen istruly in a class by itself. It does not belong to anyfamily of elements, and though it is a nonmetal,it appears on the left side of the periodic tablewith the metals. The other elements with it inGroup 1 form the alkali metal family, but obvi-ously, hydrogen does not belong with them.Indeed, if there is any element similar to hydro-gen in simplicity and abundance, it is the onlyother one on the first row, or period, of the peri-odic table: helium. Together, these two elementsmake up 99.9% of all known matter in the entireuniverse, because hydrogen atoms in stars fuse tocreate helium. Yet whereas helium is a noble gas,and therefore chemically unreactive, hydrogenbonds with all sorts of other elements. In onesuch variety of bond, with carbon, hydrogenforms the backbone for a vast collection oforganic molecules, known as hydrocarbons andtheir derivatives. Bonded with oxygen, hydrogenforms the single most important compound onEarth, and the most important complex sub-stance other than air: water. Yet when it bondswith sulfur, it creates toxic hydrogen sulfide; andon its own, hydrogen is extremely flammable.The only element whose isotopes have names,hydrogen has long been considered as a potentialsource of power and transportation: once upon atime for airships, later as a component in nuclearreactions—and, perhaps in the future, as a sourceof abundant clean energy.

HOW IT WORKS

The Essentials

The atomic number of hydrogen is 1, meaningthat it has a single proton in its nucleus. With its

single electron, hydrogen is the simplest elementof all. Because it is such a basic elemental build-ing block, figures for the mass of other elementswere once based on hydrogen, but the standardtoday is set by 12C or carbon-12, the most com-mon isotope of carbon.

Hydrogen has two stable isotopes—forms ofthe element that differ in mass. The first of these,protium, is simply hydrogen in its most commonform, with no neutrons in its nucleus. Protium(the name is only used to distinguish it from theother isotopes) accounts for 99.985% of all thehydrogen that appears in nature. The second sta-ble isotope, deuterium, has one neutron, andmakes up 0.015% of all hydrogen atoms. Tritium,hydrogen’s one radioactive isotope, will be dis-cussed below.

The fact that hydrogen’s isotopes have sepa-rate names, whereas all other isotopes are desig-nated merely by element name and mass number(for example, “carbon-12”) says something aboutthe prominence of hydrogen as an element. Notonly is its atomic number 1, but in many ways, itis like the number 1 itself—the essential piecefrom which all others are ultimately constructed.Indeed, nuclear fusion of hydrogen in the stars isthe ultimate source for the 90-odd elements thatoccur in nature.

The mass of this number-one element is not,however, 1: it is 1.008 amu, reflecting the smallquantities of deuterium, or “heavy hydrogen,”present in a typical sample. A gas at ordinarytemperatures, hydrogen turns to a liquid at -423.2°F (-252.9°C), and to a solid at –434.°F(–259.3°C). These figures are its boiling pointand melting point respectively; only the figuresfor helium are lower. As noted earlier, these twoelements make up all but 0.01% of the known

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Hydrogen


elemental mass of the universe, and are the prin-cipal materials from which stars are formed.

Normally hydrogen is diatomic, meaningthat its molecules are formed by two atoms. Atthe interior of a star, however, where the temper-ature is many millions of degrees, H2 moleculesare separated into atoms, and these atomsbecome ionized. In other words, the electron sep-arates from the proton, resulting in an ion with apositive charge, along with a free electron. Thepositive ions experience fusion—that is, theirnuclei bond, releasing enormous amounts ofenergy as they form new elements.

Because the principal isotopic form of heli-um has two protons in the nucleus, it is naturalthat helium is the element usually formed; yet itis nonetheless true—amazing as it may seem—that all the elements found on Earth were onceformed in stars. On Earth, however, hydrogenranks ninth in its percentage of the planet’sknown elemental mass: just 0.87%. In the humanbody, on the other hand, it is third, after oxygenand carbon, making up 10% of human elementalbody mass.

Hydrogen and Bonding

Having just one electron, hydrogen can bond toother atoms in one of two ways. The first option

is to combine its electron with one from the atomof a nonmetallic element to make a covalentbond, in which the two electrons are shared.Hydrogen is unusual in this regard, because mostatoms conform to the octet rule, ending up witheight valence electrons. The bonding behavior ofhydrogen follows the duet rule, resulting in justtwo electrons for bonding.

Examples of this first type of bond includewater (H2O), hydrogen sulfide (H2S), andammonia (NH3), as well as the many organiccompounds formed on a hydrogen-carbon back-bone. But hydrogen can form a second type ofbond, in which it gains an extra electron tobecome the negative ion H-, or hydride. It is thenable to combine with a metallic positive ion toform an ionic bond. Ionic hydrides are conven-ient sources of hydrogen gas: for instance, calci-um hydride, or CaH2, is sold commercially, andprovides a very convenient means of hydrogengeneration. The hydrogen gas produced by thereaction of calcium hydride with water can beused to inflate life rafts.

The presence of hydrogen in certain types ofmolecules can also be a factor in intermolecularbonding. Intermolecular bonding is the attrac-tion between molecules, as opposed to the bond-

AMONG THE USES OF TRITIUM IS IN FUEL FOR HYDROGEN BOMBS. SHOWN HERE IS THE HANFORD PLANT IN THE

STATE OF WASHINGTON, WHERE SUCH FUEL IS PRODUCED. (Roger Ressmeyer/Corbis. Reproduced by permission.)

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Hydrogen


ing within molecules, which is usually whatchemists mean when they talk about “bonding.”

Hydrogen’s Early History

Because it bonds so readily with other elements,hydrogen almost never appears in pure elementalform on Earth. Yet by the late fifteenth century,chemists recognized that by adding a metal to anacid, hydrogen was produced. Only in 1766, how-ever, did English chemist and physicist HenryCavendish (1731-1810) recognize hydrogen as asubstance distinct from all other “airs,” as gasesthen were called.

Seventeen years later, in 1783, Frenchchemist Antoine Lavoisier (1743-1794) namedthe substance after two Greek words: hydro(water) and genes (born or formed). It wasanother two decades before English chemist JohnDalton (1766-1844) formed his atomic theory ofmatter, and despite the great strides he made forscience, Dalton remained convinced that hydro-gen and oxygen in water formed “water atoms.”Around the same time, however, Italian physicistAmedeo Avogadro (1776-1856) clarified the dis-tinction between atoms and molecules, thoughthis theory would not be generally accepted untilthe 1850s.

Contemporary to Dalton and Avogadro wasSwedish chemist Jons Berzelius (1779-1848),who developed a system of comparing the massof various atoms in relation to hydrogen. Thismethod remained in use for more than a centu-ry, until the discovery of neutrons, protons, andisotopes pointed the way toward a means of mak-ing more accurate determinations of atomicmass. In 1931, American chemist and physicistHarold Urey (1893-1981) made the first separa-tion of an isotope: deuterium, from ordinarywater.


Deuterium and Tritium

Designated as 2H, deuterium is a stable isotope,whereas tritium—3H—is unstable, or radioac-tive. Not only do these two have names; they evenhave chemical symbols (D and T, respectively), asthough they were elements on the periodic table.Just as hydrogen represents the most basic pro-ton-electron combination against which otheratoms are compared, these two are respectivelythe most basic isotope containing a single neu-

THE USE OF HYDROGEN IN AIRSHIPS EFFECTIVELY ENDED WITH THE CRASH OF THE HINDENBURG in 1937. (UPI/Cor-

bis-Bettmann. Reproduced by permission.)

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Hydrogentron, and the most basic radioisotope, orradioactive isotope.

Deuterium is sometimes called “heavyhydrogen,” and its nucleus is called a deuteron. Inseparating deuterium—an achievement forwhich he won the 1934 Nobel Prize—Urey col-lected a relatively large sample of liquid hydro-gen: 4.2 qt (4 l). He then allowed the liquid toevaporate very slowly, predicting that the moreabundant protium would evaporate more quick-ly than the heavier isotope. When all but 0.034 oz(1 ml) of the sample had evaporated, he submit-ted the remainder to a form of analysis calledspectroscopy, adding a burst of energy to theatoms and then analyzing the light spectrum theyemitted for evidence of differing varieties ofatoms.

With an atomic mass of 2.014102 amu, deu-terium is almost exactly twice as heavy as pro-tium, which has an atomic mass of 1.007825. Itsmelting points and boiling points, respectively–426°F (–254°C) and –417°F (–249°C), are high-er than for protium. Often, deuterium is appliedas a tracer, an atom or group of atoms whose par-ticipation in a chemical, physical, or biologicalreaction can be easily observed.

DEUTERIUM IN WAR AND

PEACE. In nuclear power plants, deuteriumis combined with oxygen to form “heavy water”(D2O), which likewise has higher boiling andmelting points than ordinary water. Heavy wateris often used in nuclear fission reactors to slowdown the fission process, or the splitting ofatoms. Deuterium is also present in nuclearfusion, both on the Sun and in laboratories.

During the period shortly after World WarII, physicists developed a means of duplicatingthe thermonuclear fusion process. The result wasthe hydrogen bomb—more properly called afusion bomb—whose detonating device was acompound of lithium and deuterium called lithi-um deuteride. Vastly more powerful than the“atomic” (that is, fission) bombs dropped by theUnited States over Japan (Nagaski and Hiroshi-ma) in 1945, the hydrogen bomb greatlyincreased the threat of worldwide nuclear anni-hilation in the postwar years.

Yet the power that could destroy the worldalso has the potential to provide safe, abundantfusion energy from power plants—a dream as yetunrealized. Physicists studying nuclear fusion areattempting several approaches, including a

process involving the fusion of two deuterons.This fusion would result in a triton, the nucleusof tritium, along with a single proton. Theoreti-cally, the triton and deuteron would then befused to create a helium nucleus, resulting in theproduction of vast amounts of energy.

TRITIUM. Whereas deuterium has a sin-gle neutron, tritium—as its mass number of 3indicates—has two. And just as deuterium hasapproximately twice the mass of protium, tri-tium has about three times the mass, or 3.016amu. Its melting and boiling points are higherstill than those of deuterium: thus tritium heavywater (T2O) melts at 40°F (4.5°C), as comparedwith 32°F (0°C) for H2O.

Tritium has a half-life (the length of time ittakes for half the radioisotopes in a sample tobecome stable) of 12.26 years. As it decays, itsnucleus emits a low-energy beta particle, which iseither an electron or a subatomic particle called apositron, resulting in the creation of the helium-3 isotope. Due to the low energy levels involved,the radioactive decay of tritium poses little dan-ger to humans. In fact, there is always a smallquantity of tritium in the atmosphere, and thisquantity is constantly being replenished by cos-mic rays.

Like deuterium, tritium is applied in nuclearfusion, but due to its scarcity, it is usually com-bined with deuterium. Sometimes it is released insmall quantities into the groundwater as a meansof monitoring subterranean water flow. It is alsoused as a tracer in biochemical processes, and asan ingredient in luminous paints.

Hydrogen and Oxygen

WATER. Water, of course, is the mostwell-known compound involving hydrogen.Nonetheless, it is worthwhile to consider theinteraction between hydrogen and oxygen, thetwo ingredients in water, which provides aninteresting illustration of chemistry in action.

Chemically bonded as water, hydrogen andoxygen can put out any type of fire except an oilor electrical fire; as separate substances, however,hydrogen and oxygen are highly flammable. In anoxyhydrogen torch, the potentially explosivereaction between the two gases is controlled by agradual feeding process, which produces com-bustion instead of the more violent explosionthat sometimes occurs when hydrogen and oxy-gen come into contact.


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Hydrogen HYDROGEN PEROXIDE. Asidefrom water, another commonly used hydrogen-oxygen compound is hydrogen peroxide, orH2O2. A colorless liquid, hydrogen peroxide ischemically unstable (not “unstable” in the waythat a radioisotope is), and decomposes slowly toform water and oxygen gas. In high concentra-tions, it can be used as rocket fuel.

By contrast, the hydrogen peroxide used inhomes as a disinfectant and bleaching agent isonly a 3% solution. The formation of oxygen gasmolecules causes hydrogen peroxide to bubble,and this bubbling is quite rapid when the perox-ide is placed on cuts, because the enzymes inblood act as a catalyst to speed up the reaction.

Hydrogen Chloride

Another significant compound involving hydro-gen is hydrogen chloride, or HCl—in otherwords, one hydrogen atom bonded to chlorine, amember of the halogens family. Dissolved inwater, it produces hydrochloric acid, used in lab-oratories for analyses involving other acids. Nor-mally, hydrogen chloride is produced by the reac-tion of salt with sulfuric acid, though it can alsobe created by direct bonding of hydrogen andchlorine at temperatures above 428°F (250°C).

Hydrogen chloride and hydrochloric acidhave numerous applications in metallurgy, aswell as in the manufacture of pharmaceuticals,dyes, and synthetic rubber. They are used, forinstance, in making pharmaceutical hydrochlo-rides, water-soluble drugs that dissolve wheningested. Other applications include the produc-tion of fertilizers, synthetic silk, paint pigments,soap, and numerous other products.

Not all hydrochloric acid is produced byindustry, or by chemists in laboratories. Activevolcanoes, as well as waters from volcanic moun-tain sources, contain traces of the acid. So, too,does the human body, which generates it duringdigestion. However, too much hydrochloric acidin the digestive system can cause the formation ofgastric ulcers.

Hydrogen Sulfide

It may not be a pleasant subject, but hydrogen—in the form of hydrogen sulfide—is also presentin intestinal gas. The fact that hydrogen sulfide isan extremely malodorous substance once againillustrates the strange things that happen when

elements bond: neither hydrogen nor sulfur hasany smell on its own, yet together they form anextremely noxious—and toxic—substance.

Pockets of hydrogen sulfide occur in nature.If a person were to breathe the vapors for verylong, it could be fatal, but usually, the foul odorkeeps people away. The May 2001 National Geo-graphic included two stories relating to such nat-ural hydrogen-sulfide deposits, on opposite sidesof the Earth, and in both cases the presence ofthese toxic fumes created interesting results.

In southern Mexico is a system of cavesknown as Villa Luz, through which run some 20underground springs, many of them carryinglarge quantities of hydrogen sulfide. The Nation-al Geographic Society’s team had to enter thecaves wearing gas masks, yet the area teems withstrange varieties of life. Among these are fish thatare red from high concentrations of hemoglobin,or red blood cells. The creatures need this extradose of hemoglobin, necessary to move oxygenthrough the body, in order to survive on the scantoxygen supplies. The waters of the cave are fur-ther populated by microorganisms that oxidizethe hydrogen sulfide and turn it into sulfuricacid, which dissolves the rock walls and continu-ally enlarges the cave.

Thousands of miles away, in the Black Sea,explorers supported by a grant from the Nation-al Geographic Society examined evidence sug-gesting that there indeed had been a great ancientflood in the area, much like the one depicted inthe Bible. In their efforts, they had an unlikelyally: hydrogen sulfide, which had formed at thebottom of the sea, and was covered by dense lay-ers of salt water. Because the Black Sea lacks thetemperature differences that cause water to cir-culate from the bottom upward, the hydrogensulfide stayed at the bottom.

Under normal circumstances, the wreck of a1,500-year-old wooden ship would not have beenpreserved; but because oxygen could not reachthe bottom of the Black Sea—and thus wood-boring worms could not live in the toxic envi-ronment—the ship was left undisturbed. Thanksto the presence of hydrogen sulfide, explorerswere able to study the ship, the first fully intactancient shipwreck to be discovered.

Hydrocarbons

Together with carbon, hydrogen forms a hugearray of organic materials known as hydrocar-


set_vol1_sec6 9/13/01 11:51 AM

Hydrogenbons—chemical compounds whose moleculesare made up of nothing but carbon and hydrogenatoms. Theoretically, there is no limit to thenumber of possible hydrocarbons. Not only doescarbon form itself into seemingly limitlessmolecular shapes, but hydrogen is a particularlygood partner. Because it has the smallest atom ofany element on the periodic table, it can bond toone of carbon’s valence electrons without gettingin the way of the others.

Hydrocarbons may either be saturated orunsaturated. A saturated hydrocarbon is one inwhich the carbon atom is already bonded to fourother atoms, and thus cannot bond to any others.In an unsaturated hydrocarbon, however, not allthe valence electrons of the carbon atom arebonded to other atoms.

Hydrogenation is a term describing anychemical reaction in which hydrogen atoms areadded to carbon multiple bonds. There are manyapplications of hydrogenation, but one that isparticularly relevant to daily life involves its usein turning unsaturated hydrocarbons into satu-rated ones. When treated with hydrogen gas,unsaturated fats (fats are complex substancesthat involve hydrocarbons bonded to other mol-ecules) become saturated fats, which are softerand more stable, and stand up better to the heatof frying. Many foods contain hydrogenated veg-etable oil; however, saturated fats have beenlinked with a rise in blood cholesterol levels—and with an increased risk of heart disease.

PETROCHEMICALS AND FUNC-

TIONAL GROUPS. One important varietyof hydrocarbons is described under the collectiveheading of petrochemicals—that is, derivativesof petroleum. These include natural gas; petrole-um ether, a solvent; naphtha, a solvent (for exam-ple, paint thinner); gasoline; kerosene; fuel forheating and diesel fuel; lubricating oils; petrole-um jelly; paraffin wax; and pitch, or tar. A host ofother organic chemicals, including various drugs,plastics, paints, adhesives, fibers, detergents, syn-thetic rubber, and agricultural chemicals, owetheir existence to petrochemicals.

Then there are the many hydrocarbon deriv-atives formed by the bonding of hydrocarbons tovarious functional groups—broad arrays of mol-ecule types involving other elements. Amongthese are alcohols—both ethanol (the alcohol inbeer and other drinks) and methanol, used inadhesives, fibers, and plastics, and as a fuel. Other

functional groups include aldehydes, ketones,carboxylic acids, and esters. Products of thesefunctional groups range from aspirin to butyricacid, which is in part responsible for the smellboth of rancid butter and human sweat. Hydro-carbons also form the basis for polymer plasticssuch as Nylon and Teflon.

Hydrogen for Transporta-tion and Power

We have already seen that hydrogen is a compo-nent of petroleum, and that hydrogen is used increating nuclear power—both deadly and peace-ful varieties. But hydrogen has been applied inmany other ways in the transportation andpower industries.

There are only three gases practical for lift-ing a balloon: hydrogen, helium, and hot air.Each is much less dense than ordinary air, andthis gives them their buoyancy. Because hydrogenis the lightest known gas and is relatively cheap toproduce, it initially seemed the ideal choice, par-ticularly for airships, which made their debutnear the end of the nineteenth century.

For a few decades in the early twentieth cen-tury, airships were widely used, first in warfareand later as the equivalent of luxury liners in theskies. One of the greatest such craft was Ger-many’s Hindenburg, which used hydrogen to pro-vide buoyancy. Then, on May 6, 1937, the Hin-denburg caught fire while mooring at Lakehurst,New Jersey, and 36 people were killed—a tragicand dramatic event that effectively ended the useof hydrogen in airships.

Adding to the pathos of the Hindenburgcrash was the voice of radio announcer HerbMorrison, whose audio report has become a clas-sic of radio history. Morrison had come to Lake-hurst to report on the landing of the famous air-ship, but ended up with the biggest—and mosthorrifying—story of his career. As the ship burstinto flames, Morrison’s voice broke, and heuttered words that have become famous: “Oh, thehumanity!”

Half a century later, a hydrogen-related dis-aster destroyed a craft much more sophisticatedthan the Hindenburg, and this time, the mediumof television provided an entire nation with aview of the ensuing horror. The event was theexplosion of the space shuttle Challenger on Jan-uary 28, 1986, and the cause was the failure of arubber seal in the shuttle’s fuel tanks. As a result,


set_vol1_sec6 9/13/01 11:51 AM

Hydrogen







two atoms.

DUET RULE: A term describing the

distribution of valence electrons when

hydrogen atoms—which end up with only

two valence electrons—experience chemi-

cal bonding with other atoms. Most other

elements follow the octet rule.


tric current to cause a chemical reaction.

FISSION: A nuclear reaction involving

the splitting of atoms.

FUSION: A nuclear reaction that

involves the joining of atomic nuclei.

HYDROCARBON: Any chemical com-

pound whose molecules are made up of

nothing but carbon and hydrogen atoms.

HYDROGENATION: A chemical reac-

tion in which hydrogen atoms are added to

carbon multiple bonds, as in a hydro-

carbon.



and thus has a net electrical charge.




charges. The bonding of a metal to a non-

metal such as hydrogen is ionic.





mass. An isotope may either be stable or

radioactive.









electrons. Hydrogen is an exception, and

follows the duet rule.

ORGANIC: At one time, chemists used

the term “organic” only in reference to liv-

ing things. Now the word is applied to

most compounds containing carbon, with

the exception of calcium carbonate (lime-

stone) and oxides such as carbon dioxide.

RADIOISOTOPE: An isotope subject

to the decay associated with radioactivity.

A radioisotope is thus an unstable isotope.

SATURATED: A term describing a

hydrocarbon in which each carbon is

already bound to four other atoms.

TRACER: An atom or group of atoms

whose participation in a chemical, physi-

cal, or biological reaction can be easily

observed. Radioisotopes are often used as

tracers.

UNSATURATED: A term describing a

hydrocarbon, in which the carbons

involved in a multiple bond are free to

bond with other atoms.





KEY TERMS

set_vol1_sec6 9/13/01 11:51 AM

Hydrogenhydrogen gas flooded out of the craft and straightinto the jet of flame behind the rocket. All sevenastronauts aboard were killed.

THE FUTURE OF HYDROGEN

POWER. Despite the misfortunes that haveoccurred as a result of hydrogen’s high flamma-bility, the element nonetheless holds out thepromise of cheap, safe power. Just as it made pos-sible the fusion, or hydrogen, bomb—which for-tunately has never been dropped in wartime, butis estimated to be many hundreds of times morelethal than the fission bombs dropped onJapan—hydrogen may be the key to the harness-ing of nuclear fusion, which could make possiblealmost unlimited power.

A number of individuals and agencies advo-cate another form of hydrogen power, created bythe controlled burning of hydrogen in air. Notonly is hydrogen an incredibly clean fuel, pro-ducing no by-products other than water vapor, itis available in vast quantities from water. In orderto separate it from the oxygen atoms, electrolysiswould have to be applied—and this is one of thechallenges that must be addressed before hydro-gen fuel can become a reality.

Electrolysis requires enormous amounts ofelectricity, which would have to be producedbefore the benefits of hydrogen fuel could berealized. Furthermore, though the burning ofhydrogen could be controlled, there are the dan-

gers associated with transporting it across coun-try in pipelines. Nonetheless, a number of advo-cacy groups—some of whose Web sites are listedbelow—continue to promote efforts toward real-izing the dream of nonpolluting, virtually limit-less, fuel.

WHERE TO LEARN MORE

American Hydrogen Association (Web site).<http://www.clean-air.org> (June 1, 2001).

Blashfield, Jean F. Hydrogen. Austin, TX: Raintree Steck-Vaughn, 1999.

Farndon, John. Hydrogen. New York: Benchmark Books,2001.

“Hydrogen” (Web site).<http://pearl1.lanl.gov/periodic/elements/1.html>(June 1, 2001).

Hydrogen Energy Center (Web site).<http://www.h2eco.org/> (June 1, 2001).

Hydrogen Information Network (Web site).<http://www.eren.doe.gov/hydrogen/> (June 1,2001).



National Hydrogen Association (Web site).<http://www.ttcorp.com/nha/> (June 1, 2001).

Uehling, Mark. The Story of Hydrogen. New York:Franklin Watts, 1995.


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261


BOND ING ANDREACT IONS

BOND ING ANDREACT IONS

CHEM ICAL BOND ING

COMPOUNDS

CHEM ICAL REACT IONS

OX IDAT ION -REDUCT ION REACT IONS

CHEM ICAL EQU I L IBR IUM

CATALYSTS

AC IDS AND BASES

AC ID -BASE REACT IONS



CHEM ICAL BOND INGChemical Bonding

CONCEPTAlmost everything a person sees or touches indaily life—the air we breathe, the food we eat, theclothes we wear, and so on—is the result of achemical bond, or, more accurately, many chem-ical bonds. Though a knowledge of atoms andelements is essential to comprehend the subjectschemistry addresses, the world is generally notcomposed of isolated atoms; rather, atoms bondto one another to form molecules and hencechemical compounds. Not all chemical bonds arecreated equal: some are weak, and some verystrong, a difference that depends primarily onthe interactions of electrons between atoms.

HOW IT WORKS

Early Ideas of Bonding

The theory that all of matter is composed ofatoms did not originate in modern times: theatomic model actually dates back to the fifth cen-tury B.C. in Greece. The leading exponent ofatomic theory in ancient times was Democritus(c. 460-370 B.C.), who proposed that mattercould not be infinitely subdivided. At its deepestsubstructure, Democritus maintained, the mate-rial world was made up of tiny fragments hecalled atomos, a Greek term meaning “no cut” or“indivisible”

Forward-thinking though it was, Democri-tus’s idea was not what modern scientists todaywould describe as a proper scientific hypothesis.His “atoms” were not purely physical units, butrather idealized philosophical constructs, andthus, he was not really approaching the subjectfrom the perspective of a scientist. In any case,

there was no way for Democritus to test hishypothesis even if he had wanted to: by their verynature, the atoms he described were far too smallto observe. Even today, what scientists knowabout atomic behavior comes not from directobservation, but indirect means.

Hence, Democritus and the few otherancients who subscribed to atomic theory wentmore on instinct than by scientific methods. Yet,some of them were remarkably prescient in theirdescription of the bonding of atoms, in view ofthe primitive scientific methods they had at theirdisposal. No other scientist came close to theaccuracy of their theory for about 2,000 years.

ASCLEPIADES AND LUCRETI-

US DISCUSS BONDS. The physicianAsclepiades of Prusa (c.130-40 B.C.) drew on theideas of the Greek philosopher Epicurus (341-270 B.C., another proponent of atomism. Asclepi-ades speculated on the ways in which atomsinteract, and discussed “clusters of atoms,”though, of course, he had no idea what forceattracted the atoms to one another.

A few years after Asclepiades, the Romanphilosopher and poet Lucretius (c.95-c.55 B.C.espoused views that combined atomism with theidea of the “four elements”—earth, air, fire, andwater. In his great work De rerum natura (“Onthe Nature of Things”), Lucretius describedatoms as tiny spheres attached to each other byfishhook-like appendages that became entangledwith one another.

LACK OF PROGRESS UNTIL

1800. Unfortunately, the competing idea ofthe four elements, handed down by the greatphilosopher Aristotle (384-322 B.C.), prevailedover the atomic model. As the Roman Empire


ChemicalBonding


began to decline after A.D. 200, the pace of scien-tific inquiry slowed and—in Western Europe atleast—eventually came to a virtual halt. Hence,the four elements theory, which had its own fan-ciful explanations as to why certain “elements”bonded with one another, held sway in Europeuntil the beginning of the modern era.

During the seventeenth century, a mountingarray of facts from the realms of astronomy andphysics collectively disproved the Aristotelianmodel. In the area of chemistry, English physicistand chemist Robert Boyle (1627-1691) showedthat the four elements were not elements at all,because they could be broken down into simplersubstances. Yet, no one really understood whatconstituted an element until the very beginningof the nineteenth century, and until that questionwas addressed, it was difficult to move on to themystery of why certain atoms bonded with oneanother.

Early Modern Advances inBonding Theory

DALTON’S ATOMIC THEORY.

The birth of atomic theory in modern timesoccurred in 1803, when English chemist JohnDalton (1766-1844) formulated the idea that all

elements are composed of tiny, indestructibleparticles. These he called by the name Democri-tus had given them nearly 23 centuries earlier:atoms. All known substances, he said, are com-posed of some combination of atoms, which dif-fer from one another only in mass.

Though Dalton’s theory paved the way forenormous advances in the years that followed,there were a number of flaws in it. Mass alone, forinstance, is not really what differentiates oneatom from another: differences in mass reflectthe presence of subatomic particles—protonsand neutrons—of whose existence scientists wereunaware at the time.

Furthermore, the properties of atoms thatcause them to bond relate to a third subatomicparticle, the electron, which, though it con-tributes little to the mass of the atom, is all-important to the energy it possesses. As for howatoms bond to one another, Dalton had little tosay: in his conception of the atomic model, atomssimply sit adjacent to one another without form-ing true bonds, as such.

AVOGADRO AND THE MOLE-

CULE. Though Dalton recognized that thestructure of atoms in a particular element orcompound is uniform, he maintained that com-pounds are made up of compound atoms: thus,

ELECTRON-DOT DIAGRAMS OF SOME ELEMENTS IN THE PERIODIC TABLE. (Robert L. Wolke. Reproduced by permission.)


ChemicalBonding


water is a compound of “water atoms.” However,water is not an element, and therefore, there hadto be some structure—still very small, but largerthan the atom—in which atoms coalesced toform the basic materials of a compound.

That structure was the molecule, firstdescribed by Italian physicist Amedeo Avogadro(1776-1856). For several decades, Avogadro, whooriginated the idea of the mole as a means ofcomparing large groups of atoms or molecules,remained a more or less unsung hero. Only in1860, four years after his death, was his idea of

the molecule resurrected by Italian chemistStanislao Cannizzaro (1826-1910). Cannizzaro’swork was occasioned by disagreement amongscientists regarding the determination of atomicmass; however, the establishment of the molecu-lar model had far-reaching implications for the-ories of bonding.

SYMBOLIZING ATOMIC BONDS.

In 1858, German chemist Friedrich AugustKekulé (1829-1896) made the first attempt todefine the concept of valency, or the property anatom of one element possesses that determines

HYDROGEN BONDING IN HF, H2O, AND NH3. (Robert L. Wolke. Reproduced by permission.)


ChemicalBonding


its ability to bond with atoms of other elements.

A pioneer in organic chemistry, which deals with

chemical structures containing carbon, Kekulé

described the carbon atom as tetravalent, mean-

ing that it can bond to four other atoms. (The

Latin prefix tetra- means “four.”) He also specu-

lated that carbon atoms are capable of bonding

with one another in long chains.

“PURE” WATER FROM A MOUNTAIN STREAM IS ACTUALLY FILLED WITH TRACES OF THE ROCKS OVER WHICH IT

HAS FLOWED. IN FACT, WATER IS ALMOST IMPOSSIBLE TO FIND IN PURE FORM EXCEPT BY PURIFYING IT IN A LABO-RATORY. (David Muench/Corbis. Reproduced by permission.)


ChemicalBonding

This was one of the first attempts to examinethe subject of bonding using modern scientificterminology, complete with hypotheses thatcould be tested by experimentation. Kekulé alsorecognized that in order to discuss bonds under-standably, there needed to be some means of rep-resenting those bonds with symbols. He evenwent so far as to develop a system for showing thearrangement of bonds in space; however, his sys-tem was so elaborate that it was replaced in favorof a simpler one developed by Scottish chemistArchibald Scott Couper (1831-1892).

Couper, who also studied valency and thetetravalent carbon bond—he is usually givenequal credit with Kekulé for these ideas—createdan extremely straightforward schematic repre-sentation still in use by chemists today. InCouper’s system, short dashed lines serve to des-ignate chemical bonds. Hence, the bond betweentwo hydrogen atoms and an oxygen atom in awater molecule would be represented thus: H-O-H. As the understanding of bonds progressed inmodern times, this system was modified to takeinto account multiple bonds, discussed below.


Atoms, Electrons, and Ions

Today, chemical bonding is understood as thejoining of atoms through electromagnetic force.Before that understanding could be achieved,however, scientists had to unlock the secret of theelectromagnetic interactions that take placewithin an atom.

The key to bonding is the electron, discov-ered in 1897 by English physicist J. J. Thomson(1856-1940). Atomic structure in general, andthe properties of the electron in particular, arediscussed at length elsewhere in this volume.However, because these specifics are critical tobonding, they will be presented here in the short-est possible form.

At the center of an atom is a nucleus, con-sisting of protons, with a positive electricalcharge; and neutrons, which have no charge.These form the bulk of the atom’s mass, but theyhave little to do with bonding. In fact, the neu-tron has nothing to do with it, while the protonplays only a passive role, rather like a flower beingpollinated by a bee. The “bee” is the electron,

and, like a bee, it buzzes to and fro, carrying apowerful “sting”—its negative electric charge,which attracts it to the positively charged proton.

ELECTRONS AND IONS. Thoughthe electron weighs much, much less than a pro-ton, it possesses enough electric charge to coun-terbalance the positive charge of the proton. Allatoms have the same number of protons as elec-trons, and hence the net electric charge is zero.However, as befits their highly active role, elec-trons are capable of moving from one atom toanother under the proper circumstances. Anatom that loses or acquires electrons has an elec-tric charge, and is called an ion.

The atom that has lost an electron or elec-trons becomes a positively charged ion, or cation.On the other hand, an atom that gains an elec-tron or electrons becomes a negatively chargedion, or anion. As we shall see, ionic bonds, suchas those that join sodium and chlorine atoms toform NaCl, or salt, are extremely powerful.

ELECTRON CONFIGURATION.

Even in covalent bonding, which does not involveions, the configurations of electrons in twoatoms are highly important. The basics of elec-tron configuration are explained in the Electronsessay, though even there, this information is pre-sented with the statement that the student shouldconsult a chemistry textbook for a more exhaus-tive explanation.

In the simplest possible terms, electron con-figuration refers to the distribution of electronsat various positions in an atom. However,because the behavior of electrons cannot be fullypredicted, this distribution can only be expressedin terms of probability. An electron movingaround the nucleus of an atom can be comparedto a fly buzzing around some form of attractant(e.g., food or a female fly, if the moving fly ismale) at the center of a sealed room. We can statepositively that the fly is in the room, and we canpredict that he will be most attracted to the cen-ter, but we can never predict his location at anygiven moment.

As one moves along the periodic table of ele-ments, electron configurations become evermore complex. The reason is that with anincrease in atomic number, there is an increase inthe energy levels of atoms. This indicates agreater range of energies that electrons can occu-py, as well as a greater range of motion. Electronsoccupying the highest energy level in an atom are



ChemicalBonding

called valence electrons, and these are the onlyones involved in chemical bonding. By contrast,the core electrons, or the ones closest to thenucleus, play no role in the bonding of atoms.

Ionic and Covalent Bonds

THE GOAL OF EIGHT VA-

LENCE ELECTRONS. The above dis-cussion of the atom, and the electron’s place in it,refers to much that was unknown at the timeThomson discovered the electron. Protons werenot discovered for several more years, and neu-trons several decades after that. Nonetheless, theelectron proved the key to solving the riddle ofhow substances bond, and not long after Thom-son’s discovery, German chemist Richard Abegg(1869-1910) suggested as much.

While studying noble gases, noted for theirtendency not to bond, Abegg discovered thatthese gases always have eight valence electrons.His observation led to one of the most importantprinciples of chemical bonding: atoms bond insuch a way that they achieve the electron config-uration of a noble gas. This has been shown to bethe case in most stable chemical compounds.

TWO DIFFERENT TYPES OF

BONDS. Perhaps, Abegg hypothesized, atomscombine with one another because theyexchange electrons in such a way that both endup with eight valence electrons. This was an earlymodel of ionic bonding, which results fromattractions between ions with opposite electriccharges: when they bond, these ions “complete”one another.

Ionic bonds, which occur when a metalbonds with a nonmetal are extremely strong. Asnoted earlier, salt is an example of an ionic bond:the metal sodium loses an electron, forming acation; meanwhile, the nonmetal chlorine gainsthe electron to become an anion. Their ionicbond results from the attraction of oppositecharges.

Ionic bonding, however, could not explainall types of chemical bonds for the simple reasonthat not all compounds are ionic. A few yearsafter Abegg’s death, American chemist GilbertNewton Lewis (1875-1946) discovered a very dif-ferent type of bond, in which nonionic com-pounds share electrons. The result, once again, iseight valence electrons for each atom, but in thiscase, the nuclei of the two atoms share electrons.

In ionic bonding, two ions start out with dif-ferent charges and end up forming a bond inwhich both have eight valence electrons. In thetype of bond Lewis described, a covalent bond,two atoms start out as atoms do, with a netcharge of zero. Each ends up possessing eightvalence electrons, but neither atom “owns” them;rather, they share electrons.

LEWIS STRUCTURES. In additionto discovering the concept of covalent bonding,Lewis developed the Lewis structure, a means ofshowing schematically how valence electrons arearranged among the atoms in a molecule. Alsoknown as the electron-dot system, Lewis struc-tures represent the valence electrons as dots sur-rounding the chemical symbols of the atomsinvolved. These dots, which look rather like acolon, may be above or below, or on either sideof, the chemical symbol. (The dots above orbelow the chemical symbol are side-by-side, likea colon turned at a 90°-angle.)

To obtain the Lewis structure representing achemical bond, it is first necessary to know thenumber of valence electrons involved. One pairof electrons is always placed between elements,indicating the bond between them. Sometimesthis pair of valence electrons is symbolized by adashed line, as in the system developed byCouper. The remaining electrons are distributedaccording to the rules by which specific elementsbond.

MULTIPLE BONDS. Hydrogenbonds according to what is known as the duetrule, meaning that a hydrogen atom has only twovalence electrons. In most other elements—thereare exceptions, but these will not be discussedhere—atoms end up with eight valence electrons,and thus are said to follow the octet rule. If thebond is covalent, the total number of valenceelectrons will not be a multiple of eight, howev-er, because the atoms share some electrons.

When carbon bonds to two oxygen atoms toform carbon dioxide (CO2), it is represented inthe Couper system as O-C-O. The Lewis struc-ture also uses dashed lines, which stand for twovalence electrons shared between atoms. In thiscase, then, the dashed line to the left of the car-bon atom indicates a bond of two electrons withthe oxygen atom to the left, and the dashed lineto the right of it indicates a bond of two electronswith the oxygen atom on that side.



ChemicalBonding

The non-bonding valence electrons in theoxygen atoms can be represented by sets of twodots above, below, and on the outside of eachatom, for a total of six each. Combined with thetwo dots for the electrons that bond them to car-bon, this gives each oxygen atom a total of eightvalence electrons. So much for the oxygen atoms,but something is wrong with the representationof the carbon atom, which, up to this point, isshown only with four electrons surrounding it,not eight.

In fact carbon in this particular configura-tion forms not a single bond, but a double bond,which is represented by two dashed lines—asymbol that looks like an equals sign. By showingthe double bonds joining the carbon atom to thetwo oxygen atoms on either side, the carbonatom has the required number of eight valenceelectrons. The carbon atom may also form atriple bond (represented by three dashed lines,one above the other) with an oxygen atom, inwhich case the oxygen atom would have only twoother valence electrons.

Electronegativity and PolarCovalent Bonds

Today, chemists understand that most bonds areneither purely ionic nor purely covalent; rather,there is a wide range of hybrids between the twoextremes. Credit for this discovery belongs toAmerican chemist Linus Pauling (1901-1994),who, in the 1930s, developed the concept of elec-tronegativity—the relative ability of an atom toattract valence electrons.

Elements capable of bonding are assigned anelectronegativity value ranging from a minimumof 0.7 for cesium to a maximum of 4.0 for fluo-rine. Fluorine is capable of bonding with somenoble gases, which do not bond with any otherelements or each other. The greater the elec-tronegativity value, the greater the tendency ofan element to draw valence electrons to itself.

If fluorine and cesium bond, then, the bondwould be purely ionic, because the fluorine exertsso much more attraction for the valence elec-trons. But if two elements have equal electroneg-ativity values—for instance, cobalt and silicon,both of which are rated at 1.9—the bond is pure-ly covalent. Most bonds, as stated earlier, fallsomewhere in between these two extremes.

POLAR COVALENT BONDING.

When substances of differing electronegativity

values form a covalent bond, this is described aspolar covalent bonding. Sometimes these aresimply called “polar bonds,” but that is not asaccurate: all ionic bonds, after all, are polar, dueto the extreme differences in electronegativity.The term “polar covalent bond” is much morespecific, describing a bond, for instance, betweenhydrogen (2.1) and sulfur (2.6). Because sulfurhas a slightly greater electronegativity value, thevalence electrons will be slightly more attractedto the sulfur atom than to the hydrogen atom.

Another example of a polar covalent bond isthe one that forms between hydrogen and oxygen(3.5) to form H2O or water, which has a numberof interesting properties. For instance, the polarquality of a water molecule gives it a great attrac-tion for ions, and thus ionic substances such assalt dissolve easily in water. “Pure” water from amountain stream is actually filled with traces ofthe rocks over which it has flowed. In fact,water—sometimes called the “universal sol-vent”—is almost impossible to find in pure form,except when it is purified in a laboratory.

By contrast, molecules of petroleum (CH2)tend to be nonpolar, because carbon and hydro-gen have almost identical electronegativity val-ues—2.5 and 2.1 respectively. Thus, an oil mole-cule offers no electric charge to bond it with awater molecule, and for this reason, oil and waterdo not mix. It is a good thing that water mole-cules attract each other so strongly, because thismeans that a great amount of energy is requiredto change water from a liquid to a gas. If this werenot so, the oceans and rivers would vaporize, andlife on Earth could not exist as it does.

BOND ENERGY. The last two para-graphs allude to attractions between molecules,which is not the same as (nor is it as strong as)the attraction between atoms within a molecule.In fact, the bond energy—the energy required topull apart the atoms in a chemical bond—is lowfor water. This is due to the presence of hydrogenatoms, with their two (rather than eight) valenceelectrons. It is thus relatively easy to separatewater into its constituent parts of hydrogen andoxygen, through a process known as electrolysis.

Covalent bonds that involve hydrogen areamong the weakest bonds between atoms.(Again, this is different from bonds betweenmolecules.) Stronger than hydrogen bonds areregular, octet-rule covalent bonds: as one mightexpect, double covalent bonds are stronger than



ChemicalBonding






chlorine anion. Rather, for an anion

involving a single element, it is named by





ment. An atom can exist either alone or

in combination with other atoms in a

molecule.






BOND ENERGY: The energy required

to pull apart the atoms in a chemical bond.



A cation (pronounced “KAT-ie-un”) is

named after the element of which it is an

ion and thus is called, for instance, the alu-

minum ion or the aluminum cation.



representing different elements. The prin-

cipal types of bonds are covalent bonding

and ionic bonding, though few bonds are

purely one or the other. Rather, there is a

wide range of “hybrid” bonds, in accor-

dance with the electronegativity values of

the elements involved.

CHEMICAL SYMBOL: A one-or two-


element.






share valence electrons. Atoms may bond

by single, double, or triple covalent bonds,

which, in representations of Lewis struc-

tures, are shown by single, double, or triple

dashed lines. (The double dashed line

looks like an equals sign.) When atoms

have differing values of electronegativity,

they form polar covalent bonds.

DUET RULE: A term describing the

distribution of valence electrons when

hydrogen atoms—which end up with only

two valence electrons—experience chemi-

cal bonding with other atoms. Most other

elements follow the octet rule.




make up the atom’s nucleus, are essential to

chemical bonding.



electrons.


only one kind of atom. Unlike compounds,

elements cannot be broken down chemi-

cally into other substances.

KEY TERMS


ChemicalBonding


ION: An atom that has lost or gained

one or more electrons, and thus has a net

electrical charge. Ions may either be anions

or cations.



between ions with opposite electrical

charges.

LEWIS STRUCTURE: A means of

showing schematically how valence elec-

trons are distributed among the atoms in a

molecule. Also known as the electron-dot

system, Lewis structure represents pairs of

electrons with a symbol rather like a colon,

which—depending on the situation—can

be placed above, below, or on either side of

the chemical symbol. In the Lewis struc-

ture, the pairs of electrons involved in

chemical bonds are usually represented by

a dashed line.


ly, but not always, representing more



molecules.


has no electrical charge. Neutrons are

found at the nucleus of an atom, alongside

protons.








electrons. Hydrogen is an exception, and

follows the duet rule. A few elements fol-

low other rules, and some (most notably

the noble gases) do not typically bond with

other elements.

PERIODIC TABLE OF ELEMENTS:

A chart showing the elements arranged in




chemical characteristics.

POLAR COVALENT BONDING: The

type of chemical bonding between atoms

that have differing values of electronegativ-

ity. If the difference is extreme, of course,

the bond is not a covalent bond at all, but

an ionic bond. Thus, although these are

sometimes called polar bonds, they are

more properly identified as polar covalent

bonds.


in an atom.




in chemical bonding. By contrast, the core

electrons, or the ones at lower energy lev-

els, play no role in the bonding of atoms.

VALENCY: The property of the atom of

one element that determines its ability to

bond with atoms of other elements.



ChemicalBonding

single ones, and triple covalent bonds are

stronger still. Strongest of all are ionic bonds,

involved in the bonding of a metal to a metal, or

a metal to a nonmetal, as in salt. The strength of

the bond energy in salt is reflected by its boiling

point of 1,472°F (800°C), much higher than that

of water, at 212°F (100°C).

WHERE TO LEARN MORE

“Chemical Bond” (Web site). <http://www.science.uwaterloo.ca/~cchieh/cact/c120/chembond.html>(May 18, 2001).

“Chemical Bonding” Oklahoma State University (Website). <http://www.okstate.edu/jgelder/bondtable.html> (May 19, 2001).

“Chemical Bonding” (Web site). <http://pc65.frontier.osrhe.edu/hs/science/pbond.htm> (May 19, 2001).

“Chemical Bonding” (Web site). <http://users.senet.com.au/~rowanb/chem/chembond.htm> (May 19,2001).

“The Colored Periodic Table and Chemical Bonding” (Website). <http://students.washington.edu/manteca/>(May 19, 2001).


Linus Pauling and the Twentieth Century: An Exhibition(Web site). <http://www.paulingexhibit.org> (May19, 2001).

“Valence Shell Electron Pair Repulsion (VSEPR).”<http://www.shef.ac.uk/~chem/vsepr/chime/vsepr.html> (May 18, 2001).

White, Florence Meiman. Linus Pauling, Scientist andCrusader. New York: Walker, 1980.





COMPOUNDSCompounds

CONCEPTA compound is a chemical substance in whichatoms combine in such a way that the compoundalways has the same composition, unless it ischemically altered in some way. Elements makeup compounds, and although there are onlyabout 90 elements that occur in nature, there areliterally many millions of compounds. A com-pound is not the same as a mixture, which has avariable composition, but until chemists under-stood the atomic and molecular substructure ofcompounds, the distinction was not always clear.When atoms of one element bond to atoms ofanother, they form substances quite differentfrom either element. Sugar, for instance, is madeup of carbon, the material in coal and graphite,along with two gases, hydrogen and oxygen.None of these is sweet, yet when they are broughttogether in just the right way, they make the com-pound that sweetens everything from breakfastcereals to colas. A compound cannot be under-stood purely in terms of its constituent elements;an awareness of the structure is also needed.Likewise, it is important to know just how toname a compound, using a uniform terminolo-gy, since there are far too many substances in theworld to give each an individual name.

HOW IT WORKS

Elements and Compounds

A compound is a substance made up of atomsrepresenting more than one element, and theseatoms are typically joined in molecules. Thecomposition of a compound is always the same:for instance, water always contains molecules

composed of two hydrogen atoms bonded to asingle oxygen atom. It can be frozen or boiled,but the molecules themselves are unchanged,because freezing and boiling are merely physicalprocesses. To transform one compound intoanother, on the other hand, requires a chemicalchange or reaction.

Likewise, a chemical change is required tobreak down a compound into its constituent ele-ments. Water can be broken down by passing apowerful electric current through it. Thisprocess, known as electrolysis, separates waterinto hydrogen and oxygen, both of which arehighly flammable gases. The fact that they can bejoined to form water, a substance used for put-ting out most kinds of fires, illustrates the typesof changes elements undergo when they join toform compounds.

The atoms in a compound do not changeinto other atoms; or to put it another way, theelements do not change into other elements.Though two hydrogen atoms bond with an oxy-gen atom to form water, they are still hydrogens,and the oxygen is still an oxygen. The atoms’ ele-mental identity thus remains intact, and if theseelements are separated, they can join with otherelements to form entirely different compounds.

Letters and Elements;Words and Compounds

The nature of a compound is almost paradoxical:the constituent elements remain the same, yet thecompound typically bears little resemblance tothe elements that form it. How can this be? Per-haps an analogy to language, and the symbolsused to express it, will serve to show that this


Compounds


There are between 88 and 92 elements thatappear in nature; figures vary, because a handfulof the elements with atomic numbers less than 92have not actually been found in nature, but wereproduced in laboratories. (All elements with anatomic number higher than that of uranium,which is 92, are artificial.) In any case, there arefar more elements than there are letters in theEnglish alphabet; yet even with 26 characters, it ispossible to form an almost limitless vocabulary.

Let us focus on a single letter, g. Phonetical-ly, it can make a hard sound, as in great, or a softone, as in geranium. It may be silent, as in light,or it may combine with another letter to makeeither a typical “g” sound (edge) or a totallyunexpected one (rough). It can slide into a vowelsound smoothly, as in singer, or with an almostimperceptible pause, as in finger. The complexi-ties multiply when we consider the range of pos-sibilities for the letter g, and all the resultantmeanings involved: from God to dog, or fromglory to degeneration.

WHAT THIS ANALOGY SHOWS

ABOUT CHEMISTRY. One could go onendlessly in this vein, and all with just one letter;however, a few points need to be made regardingthe analogy between letters of the alphabet andelements. First of all, in all of the examples givenabove, or any number of others, the g stillremained a g. In the same way, an atom of hydro-gen (another g-word!) is always a hydrogen,whether it combines with oxygen to form water,nitrogen to make ammonia, or carbon to pro-duce petroleum.

Secondly, note that this illustration wouldhave been more difficult if the letter chosen hadbeen q or z, since these are used more rarely. Like-wise, there are elements such as technetium orpraseodymium that seldom come up in discus-sions of compounds. On the other hand, onecannot go far in the study of chemistry withoutrunning across compounds involving hydrogen,carbon, oxygen, nitrogen, various metals, orhalogens such as fluorine.

A third point to consider is the fact thatthere is nothing about g itself that provides anyinformation as to the meaning of the word ithelps to form. Here the analogy to elements isimperfect, because the nature of the element’ssubatomic structure—particularly the configura-tion of electrons on the outer shell of the atom—

apparent contradiction is not a contradiction at

all—it is merely evidence of the complexities that

occur when a simple particle contributes to a

larger combination.

A MIXTURE VERSUS A COMPOUND. (Robert L. Wolke. Repro-



Compoundsprovides a great deal of information as to how itwill combine with other atoms.

Still, it is evident, as we have seen, that theproperties of a compound cannot easily be pre-dicted by studying the properties of its con-stituent elements—just as one cannot begin todefine a word simply because it contains a g. Onthe other hand, knowledge of the sounds that gmakes may help us to pronounce a word con-taining that letter. In the same way, the fact that acompound includes carbon provides a clue thatthe compound might be organic.

Even when we have all the letters to form aword, we still need to know how they areordered. Loge and ogle contain the same letters,but one is a noun referring to a theater box,whereas the other is a verb meaning the act oflooking at someone in an improper fashion. Inchemistry, these are called isomers: two sub-stances having identical chemical formulas, butdiffering chemical structures.

It so happens that g is also part of the suffix-ing, used in forming a gerund. Aside from beingyet another g-word, a gerund is a substantive, ornoun form of a verb: for instance, going. Onceagain, there is a similarity between words andcompounds. Just as certain classes of words areformed by the addition of regular suffixes, so thenaming of whole classes of compounds can beachieved by means of a uniform system ofnomenclature.


Distinguishing Between Com-pounds and Mixtures

To continue the analogy used above just one stepfurther, a word is not just a collection of letters; ithas to have a meaning. Likewise, the fact that var-ious substances are mixed together does not nec-essarily make them a compound. Actually, thedifference between a word and a mere collectionof letters is somewhat greater than the differencebetween a compound and a mixture—a sub-stance in which elements are not chemicallybonded, and in which the composition is vari-able. A nonsensical string of characters serves nolinguistic purpose; on the other hand, mixturesare an integral part of life.

Tea, whether iced or hot, is a mixture. So iscoffee, or even wine. In each case, substances areadded together and subjected to a process, butthe result is not a compound. We know thisbecause the composition varies. Depending onthe coffee beans used, for instance, coffee canhave a wide variety of flavors. If, in the brewingprocess, too much coffee is used in proportion tothe water, the resulting mixture will be strong orbitter; on the hand, an insufficient coffee-to-water ratio will produce coffee that is too weak.

Note that a number of terms have been usedhere that, from a scientific standpoint at least, arevague. How weak is “too weak”? That all dependson the tastes of the person making the coffee. Butas long as coffee beans and hot water are used, nomatter what the proportion, the mixture is stillcoffee. On the other hand, when two oxygenatoms, rather than one, are chemically combinedwith two hydrogen atoms, the result is not“strong water.” Nor is it “oxygenated water”: it ishydrogen peroxide, a substance no one shoulddrink.

Three principal characteristics serve to dif-ferentiate a compound from a mixture. First, aswe have seen, a compound has a definite andconstant composition, whereas a mixture canexist with virtually any proportion between itsconstituent parts. Second, elements lose theircharacteristic elemental properties once theyform a compound, but the parts of a mixture donot. (For example, when mixed with water, sugaris still sweet.) Third, the formation of a com-pound involves a chemical reaction, whereas amixture can be created simply by the physical actof stirring items together.

The Atomic and MolecularKeys

The means by which compounds are formed arediscussed numerous times, and in various ways,throughout this book. A few of those particularswill be mentioned briefly below, in relation to thenaming of compounds, but for the most part,there will be no attempt to explain the details ofthe processes involved in chemical bonding. Thereader is therefore encouraged to consult theessays on Chemical Bonding and Electrons.

It is important, nonetheless, to recognizethat chemists’ knowledge is based on their under-standing of the atom and the ways that electrons,negatively charged particles in the atom, bring



Compounds about bonds between elements. Awareness ofthese specifics emerged only at the beginning ofthe twentieth century, with the discovery of sub-atomic particles. Another important thresholdhad been crossed a century before, with thedevelopment of atomic theory by Englishchemist John Dalton (1766-1844), and of molec-ular theory by Italian physicist Amedeo Avogadro(1776-1856).

Around the same time, French chemistsAntoine Lavoisier (1743-1794) and Joseph-LouisProust (1754-1826), respectively, clarified thedefinitions of “element” and “compound.” Untilthen, the idea of a compound had little precisemeaning for chemists, who often used the termto describe a mixture. Thus, French chemistClaude Louis Berthollet (1748-1822) assertedthat compounds have variable composition, andfor evidence he pointed to the fact that whensome metals are heated, they form oxides, inwhich the percentage of oxygen increases withtemperature.

Proust, on the other hand, maintained thatcompounds must have a constant composition,an argument supported by Dalton’s atomic theo-ry. Proust worked to counter Berthollet’s posi-tions on a number of particulars, but was stillunable to explain why metals form variablealloys, or combinations of metals; no chemist atthe time understood that an alloy is a mixture,not a compound.

Nonetheless, Proust was right in his theoryof constant composition, and Berthollet wasincorrect on this score. With the discovery ofsubatomic structures, it became possible todevelop highly sophisticated theories of chemicalbonding, which in turn facilitated understandingof compounds.

Types of Compounds

ORGANIC COMPOUNDS. Thoughthere are millions of compounds, these can begrouped into just a few categories. Organic com-pounds, of which there are many millions, arecompounds containing carbon. The only majorgroupings of carbon-containing compounds thatare not considered organic are carbonates, suchas limestone, and oxides, such as carbon dioxide.Organic compounds can be further subdividedinto a number of functional groups, such as alco-hols. Within the realm of organic compounds,whether natural or artificial, are petroleum and

its many products, as well as plastics and othersynthetic materials.

The term “organic,” as applied in chemistry,does not necessarily refer to living things, sincethe definition is based on the presence of carbon.Nonetheless, all living organisms are organic, andthe biochemical compounds in living thingsform an important subset of organic com-pounds. Biochemical compounds are, in turn,divided into four families: carbohydrates, pro-teins, nucleic acids, and lipids.

INORGANIC COMPOUNDS. In-organic compounds can be classified accordingto five major groupings: acids, bases, salts, oxides,and others. An acid is a compound which, whendissolved in water, produces positive ions (atomswith an electric charge) of hydrogen. Bases aresubstances that produce negatively chargedhydroxide (OH–) ions when dissolved in water.An oxide is a compound in which the only nega-tively charged ion is an oxygen, and a salt isformed by the reaction of an acid with a base.Generally speaking, a salt is any combination of ametal and a nonmetal, and it can contain ions ofany element but hydrogen.

The remaining inorganic compounds, classi-fied as “others,” are those that do not fit into anyof the groupings described above. An importantsubset of this broad category are the coordina-tion compounds, formed when one or more ionsor molecules contributes both electrons neces-sary for a bonding pair, in order to bond with ametallic ion or atom.

Naming Compounds

In the early days of chemistry as a science, com-mon names were applied to compounds. Water isan example of a common name; so is sugar, aswell as salt. These names work well enough ineveryday life, and in fact, chemists still refer towater simply as “water.” (The only other com-mon name still used in chemistry is ammonia.)

But as the number of compounds discoveredand developed by chemists began to proliferate,the need for a systematic means of naming thembecame apparent. With millions of compounds,it would be nearly impossible to come up withnames for every one. Furthermore, commonnames tell chemists nothing about the chemicalproperties of a particular substance.



CompoundsToday, chemists use a system of nomencla-ture that is rather detailed but fairly easy tounderstand, once the rules are understood. Wewill examine this system briefly, primarily as itrelates to binary compounds—compounds con-taining just two elements. Binary compounds aredivided into three groups. The first two are ioniccompounds, involving metals that form positive-ly charged ions, or cations (pronounced KAT-ie-unz). The third consists of compounds that con-tain only nonmetals. These groups are:

• Type I: Ionic compounds involving a metalthat always forms a cation of a certain elec-tric charge.

• Type II: Ionic compounds involving a metal(typically a transition metal) that formscations with differing charges.

• Type III: Compounds containing only non-metals.

CATIONS AND ANIONS. Cationsare represented symbolically thus: H+, or Mg2+.The first, a hydrogen cation, has a positive chargeof 1, but note that the 1 is not shown—just as thefirst power of a number is never designated inmathematics. In the second example, a magne-sium cation, the superscript number, combinedwith the plus sign (which can either follow thenumber, as is shown here, or proceed it) indicatesthat the atom has a positive charge of two. Thus,even if one were not told that this is a cation,it would be easy enough to discern from thenotation.

Anions (AN-ie-unz), or negatively chargedions, are represented in a similar way: H- forhydride, an anion of hydrogen; or O2– for oxide.Note, however, that the naming of anions andcations is different. Cations are always called, forexample, a hydrogen cation, or a magnesiumcation. On the other hand, names of simpleanions (involving a single atom) are formed bytaking the root of the element name and addingan -ide: for example, fluoride (F–).

TYPE I BINARY COMPOUNDS.

In a binary ionic compound, a metal combineswith a nonmetal. The metal loses one or moreelectrons to become a cation, while the nonmetalgains one or more electrons to become an anion.Thus, table salt is formed by the joining of acation of the metal sodium (Na+) and an anionof the nonmetal chlorine (Cl–). Instead of thecommon name “salt,” which can apply to a range

of substances, its chemical name is sodium chloride.

In naming Type I binary compounds, ofwhich sodium chloride is an example, the cationis always represented first by the name of the ele-ment. The anion follows, with the root name ofthe element attached to the -ide suffix, asdescribed above. Another example is calcium sul-fide, formed by a cation of the metal calcium(Ca2+) and an anion of the nonmetal sulfur (S2–).

TYPE II BINARY COMPOUNDS.

The chemical nomenclature for type II binarycompounds is somewhat more complicated,because they involve metals that can have multi-ple positive charges. This is particularly true ofthe transition metals, a family of 40 elements inthe middle of the periodic table distinguishedfrom other elements by a number of characteris-tics. The name “transition” thus implies a breakin the even pattern of the periodic table.

Iron (Fe), for instance, is a transition metal,and it can form cations with charges of 2+ or 3+,while copper (Cu) can form cations with chargesof 1+ or 2+. When encountering positivelycharged cations, it is not enough to say, forinstance, “iron oxide,” or “copper sulfide,”because it is not clear which iron cation isinvolved. To solve the problem, chemists use asystem of Roman numerals.

According to this system, the cationsreferred to above are expressed in the name of aType II binary compound thus: iron (II), iron(III), copper (I), and copper (II). This is followedwith the name of the anion, as before, using the -ide suffix. Note that the Roman numeral is usu-ally the same as the number of positive charges inthe cation.

It should be noted, also, that there is an oldersystem for naming Type II binary compoundswith terms that incorporate the element name—often the Latin original, reflected in the chemicalsymbol—with a suffix. For instance, this systemuses the word “ferrous” for iron (II) and “ferric”for iron (III). However, this method of nomen-clature is increasingly being replaced by the onewe have described here.

TYPE III BINARY COMPOUNDS.

In a Type III binary compound involving onlynonmetals, the first element in the formula isreferred to simply by its element name, as thoughit were a cation, while the second element is givenan -ide suffix, as though it were an anion. If there



Compounds is more than one atom present, prefixes are usedto indicate the number of atoms. These prefixesare listed below. It should be noted that mono- isnever used for naming the first element in a typeIII binary compound.

• mono-: 1

• di-: 2

• tri-: 3

• tetra-: 4

• penta-: 5

• hexa-: 6

• hepta-: 7

• octa-: 8

Thus CO2 is called carbon dioxide, indicat-ing one carbon atom and two oxygens. Again,mono- is not used for the first element, but it isused for the second: hence, the name of the com-pound with the formula CO is carbon monoxide.It is also possible to know the formula for a com-pound simply from the name: if confronted witha name such as “dinitrogen pentoxide,” forinstance, it is fairly easy to apply the rules gov-erning these prefixes to discern that the sub-stance is N2O5. Note that vowels at the end of aprefix are dropped when the name of the elementthat follows it also begins with a vowel: we saymonoxide, not “monooxide”; and pentoxide, not“pentaoxide.” This makes pronunciation mucheasier.

POLYATOMIC IONS AND ACIDS.

More complicated rules apply for polyatomicions, which are charged groupings of atoms suchas NH4NO3, or ammonium nitrate. The only wayto learn how to name polyatomic ions is bymemorizing the names of the constituent parts.In the above example, for instance, the first partof the formula, which has a positive charge, isalways called ammonium, while the second part,which has a negative charge, is always callednitrate. A good chemistry textbook should pro-vide a table listing the names of common poly-atomic ions.

A number of polyatomic ions are calledoxyanions, meaning that they include varyingnumbers of oxygen atoms combined with atoms of other elements. There are rules for des-ignating the names of these polyatomic ions,some of which are listed in the essay on Ions andIonization.

Still other rules govern the naming of acids;here again, the operative factor is the presence ofoxygen. Thus, if the anion does not contain oxy-

gen, the name of the acid is created by adding theprefix hydro- and the suffix -ic, as in hydrochlo-ric acid. If the anion does contain oxygen, theroot name of the principal element is joined to asuffix: depending on the relative numbers of oxy-gen atoms, this may be -ic or -ous.

Isomers

As observed much earlier, isomers are like twowords with the same letters, but arranged in dif-ferent ways. Specifically, isomers are chemicalcompounds having the same formula, but inwhich the atoms are arranged differently, therebyforming different compounds. There are twoprincipal types of isomer: structural isomers,which differ according to the attachment ofatoms on the molecule, and stereoisomers, whichdiffer according to the locations of the atoms inspace.

An example of structural isomerism is thedifference between propyl alcohol and isopropyl(rubbing) alcohol: these two have differing prop-erties, because their alcohol functional groupsare not attached to the same carbon atom in thecarbon chain to which the functional group isattached.

In a stereoisomer, on the other hand, atomsare attached in the same order, but have differentspatial relationships. If functional groups arealigned on the same side of a double bondbetween two carbon atoms, this is called a cis iso-mer, from a Latin word meaning “on this side.” Ifthey are on opposite sides, it is called a trans(“across”) isomer. Hence the term “trans fats,”which are saturated fats that improve certainproperties of foods—including taste—but whichmay contribute to heart disease.

WHERE TO LEARN MORE

“Chemical Compounds” (Web site). <http://www.netac-cess.on.ca/~dbc/cic_hamilton/comp.html> (June 2,2001).

“Chemtutor Compounds.” Chemtutor (Web site).<http://www.chemtutor.com/compoun.htm> (June2, 2001).

Fullick, Ann. Matter. Des Plaines, IL: HeinemannLibrary, 1999.

“Glossary of Products with Hazards A to Z” Environmen-tal Protection Agency (Web site). <http://www.epa.gov/grtlakes/seahome/housewaste/house/products.htm> (June 2, 2001).

Knapp, Brian J. Elements, Compounds, and Mixtures, Vol-ume 2. Edited by Mary Sanders. Danbury, CT: Grolier



Compounds


ACID: An inorganic compound that,

when dissolved in water, produces positive

ions—that is, cations—of hydrogen, desig-

nated symbolically as H+ ions. (This is just

one definition; see Acids and Bases; Acid-

Base Reactions for more.)

ANION: The negatively charged ion

that results when an atom gains one or

more electrons. The word is pronounced

“AN-ie-un.”

BASE: An inorganic compound that

produces negative hydroxide ions when it

is dissolved in water. These anions are des-

ignated by the symbol OH-. (This is just

one definition; see Acids and Bases; Acid-

Base Reactions for more.)

BINARY COMPOUND: A compound

that contains just two elements. For

the purposes of establishing compound

names, binary compounds are divided into

Type I, Type II, and Type III.

CATION: The positively charged ion

that results when an atom loses one or

more electrons. The word “cation” is pro-

nounced “KAT-ie-un.”



representing different elements.


atoms of more than one element, which are

chemically bonded and usually joined in

molecules. The composition of a com-

pound is always the same, unless it is

changed chemically.

COORDINATION COMPOUNDS:

Inorganic compounds formed when one

or more ions or molecules contribute both

electrons necessary for a bonding pair in

order to bond with a metallic ion or atom.

ELECTRON: A negatively charged par-

ticle in an atom.

INORGANIC COMPOUND: For the

most part, inorganic compounds are any

compounds that do not contain carbon.

However, carbonates and carbon oxides are

also inorganic compounds. Compare with

organic compounds.







charges.

IONIC COMPOUND: A compound in

which ions are present. Ionic compounds

contain at least one metal joined to anoth-

er element by an ionic bond.

ISOMERS: Substances which have the

same chemical formula, but which have

different chemical properties due to differ-

ences in the arrangement of atoms.

MIXTURE: A substance in which ele-

ments are not chemically bonded, and in

which the composition is variable. A mix-

ture is distinguished from a compound.

ORGANIC COMPOUND: Generally

speaking, a compound containing carbon.

The only exceptions are the carbonates (for

example, calcium carbonate or limestone)

and oxides, such as carbon dioxide.

OXIDE: An inorganic compound in

which the only negatively charged ion is an

oxygen.

KEY TERMS


Compounds


Educational, 1998.

“List of Compounds” (Web site). <http://www.speclab.com/compound/chemabc.htm> (June 2, 2001).

Maton, Anthea. Exploring Physical Science. Upper SaddleRiver, N.J.: Prentice Hall, 1997.

“Molecules and Compounds.” General Chemistry Online

(Web site). <http://antoine.fsu.umd.edu/chem/

senese/101/compounds/index.shtml> (June 2, 2001).

Oxlade, Chris. Elements and Compounds. Chicago, IL:

Heinemann Library, 2001.



SALT: An inorganic compound formed

by the reaction of an acid with a base. Gen-

erally speaking, a salt is a combination of a

metal and a nonmetal, and it can contain

ions of any element but hydrogen.

TYPE I BINARY COMPOUNDS:

Ionic compounds involving a metal that

always forms a cation of a certain electric

charge.

TYPE II BINARY COMPOUNDS:

Ionic compounds involving a metal (typi-

cally a transition metal) that forms cations

with differing charges.

TYPE III BINARY COMPOUNDS:

Compounds containing only nonmetals.








CHEM ICAL REACT IONSChemical Reactions

CONCEPTIf chemistry were compared to a sport, then thestudy of atomic and molecular properties, alongwith learning about the elements and how theyrelate on the periodic table, would be like goingto practice. Learning about chemical reactions,which includes observing them and sometimesproducing them in a laboratory situation, is likestepping out onto the field for the game itself.Just as every sport has its “vocabulary”—the con-cepts of offense and defense, as well as variousrules and strategies—the study of chemical reac-tions involves a large set of terms. Some aspectsof reactions may seem rather abstract, but theeffects are not. Every day, we witness evidence ofchemical reactions—for instance, when a fireburns, or metal rusts. To an even greater extent,we are surrounded by the products of chemicalreactions: the colors in the clothes we wear, orartificial materials such as polymers, used ineverything from nylon running jackets to plasticmilk containers.

HOW IT WORKS

What Is a Chemical Reaction?

If liquid water is boiled, it is still water; likewisefrozen water, or ice, is still water. Melting, boiling,or freezing simply by the application of a changein temperature are examples of physical changes,because they do not affect the internal composi-tion of the item or items involved. A chemicalchange, on the other hand, occurs when the actu-al composition changes—that is, when one sub-stance is transformed into another. Water can bechemically changed, for instance, when an elec-

tric current is run through a sample, separating itinto oxygen and hydrogen gas.

Chemical change requires a chemical reac-tion, a process whereby the chemical propertiesof a substance are altered by a rearrangement ofthe atoms in the substance. Of course we cannotsee atoms with the naked eye, but fortunately,there are a number of clues that tell us when achemical reaction has occurred. In many chemi-cal reactions, for instance, the substance mayexperience a change of state or phase—as forinstance when liquid water turns into gaseousoxygen and hydrogen as a result of electrolysis.

HOW DO WE KNOW WHEN A

CHEMICAL REACTION HAS OC-

CURRED? Changes of state may of course bemerely physical—as for example when liquidwater is boiled to form a vapor. (These and otherexamples of physical changes resulting from tem-perature changes are discussed in the essays onProperties of Matter; Temperature and Heat.)The vapor produced by boiling water, as notedabove, is still water; on the other hand, when liq-uid water is turned into the elemental gaseshydrogen and oxygen, a more profound changehas occurred.

Likewise the addition of liquid potassiumchromate (K2CrO4) to a solution of bariumnitrate (Ba[NO3]2 forms solid barium chromate(BaCrO4). In the reaction described, a solution isalso formed, but the fact remains that the mix-ture of two solids has resulted in the formation ofa solid in a different solution. Again, this is a farmore complex phenomenon than the mere freez-ing of water to form ice: here the fundamentalproperties of the materials involved havechanged.


ChemicalReactions


The physical change of water to ice or steam,

of course, involves changes in temperature; like-

wise, chemical changes are often accompanied by

changes in temperature, the crucial difference

being that these changes are the result of alter-

ations in the chemical properties of the sub-

stances involved. Such is the case, for instance,

when wood burns in the presence of oxygen:

once wood is turned to ash, it has become an

entirely different mixture than it was before.

Obviously, the ashes cannot be simply frozen to

turn them back into wood again. This is an

example of an irreversible chemical reaction.

Chemical reactions may also involve changes

in color. In specific proportions and under the

right conditions, carbon—which is black—can

be combined with colorless hydrogen and oxygen

to produce white sugar. This suggests another

kind of change: a change in taste. (Of course, not

every product of a chemical reaction should be

tasted—some of the compounds produced may

be toxic, or at the very least, extremely unpleasant

to the taste buds.) Smell, too, can change. Sulfur

is odorless in its elemental form, but when

combined with hydrogen to form hydrogen sul-

fide (H2S), it becomes an evil-smelling, highly

toxic gas.

The bubbling of a substance is yet another

clue that a chemical reaction has occurred.

Though water bubbles when it boils, this is mere-

A MICROSCOPIC CLOSE-UP OF NYLON FIBERS. NYLON IS FORMED THROUGH A SYNTHESIS REACTION. (Science Pictures

Limited/Corbis. Reproduced by permission.)


ChemicalReactions

ly because heat has been added to the water,increasing the kinetic energy of its molecules. Butwhen hydrogen peroxide bubbles when exposedto oxygen, no heat has been added. As with manyof the characteristics of a chemical reactiondescribed above, bubbling does not always occurwhen two chemicals react; however, when one ofthese clues is present, it tells us that a chemicalreaction may have taken place.


Chemical Equations

In every chemical reaction, there are participantsknown as reactants, which, by chemically react-ing to one another, result in the creation of aproduct or products. As stated earlier, a chemicalreaction involves changes in the arrangement ofatoms. The atoms in the reactants (or, if the reac-tant is a compound, the atoms in its molecules)are rearranged. The atomic or molecular struc-ture of the product is different from that of eitherreactant.

Note, however, that the number of atomsdoes not change. Atoms themselves are neithercreated nor destroyed, and in a chemical reac-tion, they merely change partners, or lose part-ners altogether as they return to their elementalform. This is a critical principle in chemistry, onethat proves that medieval alchemists’ dream ofturning lead into gold was based on a fallacy.Lead and gold are both elements, meaning thateach has different atoms. To imagine a chemicalreaction in which one becomes the other is likesaying “one plus one equals one.”

SYMBOLS IN A CHEMICAL

EQUATION. In a mathematical equation,the sums of the numbers on one side of theequals sign must be the same as the sum of thenumbers on the other side. The same is true of achemical equation, a representation of a chemi-cal reaction in which the chemical symbols onthe left stand for the reactants, and those on theright are the product or products. Instead of anequals sign separating them, an arrow, pointingto the right to indicate the direction of the reac-tion, is used.

Chemical equations usually include notationindicating the state or phase of matter for the


reactants and products. These symbols are as follows:

• (s): solid

• (l): liquid

• (g): gas

• (aq): dissolved in water (an aqueous solu-tion)

The fourth symbol, of course, does not indi-cate a phase of matter per se (though obviously itappears to be a liquid); but as we shall see, aque-ous solutions play a role in so many chemicalreactions that these have their own symbol. Atany rate, using this notation, we begin to sym-bolize the reaction of hydrogen and oxygen toform water thus: H(g) + O(g) 4H2O(l).

This equation as written, however, needs tobe modified in several ways. First of all, neitherhydrogen nor oxygen is monatomic. In otherwords, in their elemental form, neither appearsas a single atom; rather, these form diatomic(two-atom) molecules. Therefore, the equationmust be rewritten as H2(g) + O2(g) 4H2O(l). Butthis is still not correct, as a little rudimentaryanalysis will show.

A CLOSE-UP OF RUST ON A PIECE OF METAL. RUST IS

AN EXAMPLE OF A COMBINATION REACTION, NOT A

DECOMPOSITION REACTION. (Marko Modic/Corbis. Reproduced

by permission.)


ChemicalReactions

Balancing Chemical Equations

When checking a chemical equation, one shouldalways break it down into its constituent ele-ments, to determine whether all the atoms on the left side reappear on the right side; otherwise,the result may be an incorrect equation, along thelines of “1 + 1 = 1.” That is exactly what has hap-pened here. On the left side, we have two hydro-gen atoms and two oxygen atoms; on the rightside, however, there is only one oxygen atom togo with the two hydrogens.

Obviously, this equation needs to be correct-ed to account for the second oxygen atom, andthe best way to do that is to show a second watermolecule on the right side. This will be repre-sented by a 2 before the H2O, indicating that twowater molecules now have been created. The 2, orany other number used for showing more thanone of a particular chemical species in a chemicalequation, is called a coefficient. Now we haveH2(g) + O2(g) 42H2O(l).

Is this right? Once again, it is time to analyzethe equation, to see if the number of atoms onthe left equals the number on the right. Suchanalysis can be done in a number of ways: forinstance, by symbolizing each chemical species asa circle with chemical symbols for each elementin it. Thus a single water molecule would beshown as a circle containing two H’s and one O.

Whatever the method used, analysis willreveal that the problem of the oxygen imbalancehas been solved: now there are two oxygens onthe left, and two on the right. But solving thatproblem has created another, because now thereare four hydrogen atoms on the right, as com-pared with two on the left. Obviously, anothercoefficient of 2 is needed, this time in front of thehydrogen molecule on the left. The changedequation is thus written as: 2H2(g) + O2(g) 42H2O(l). Now, finally, the equation is correct.

THE PROCESS OF BALANC-

ING CHEMICAL EQUATIONS. Whatwe have done is to balance an unbalanced equa-tion. An unbalanced equation is one in which thenumbers of atoms on the left are not the same asthe number of atoms on the right. Though anunbalanced equation is incorrect, it is sometimesa necessary step in the process of finding the bal-anced equation—one in which the number ofatoms in the reactants and those in the productare equal.

In writing and balancing a chemical equa-tion, the first step is to ascertain the identities, byformula, of the chemical species involved, as wellas their states of matter. After identifying thereactants and product, the next step is to write anunbalanced equation. After that, the unbalancedequation should be subjected to analysis, asdemonstrated above.

The example used, of course, involves a fair-ly simple substance, but often, much more com-plex molecules will be part of the equation. Inperforming analysis to balance the equation, it isbest to start with the most complex molecule,and determine whether the same numbers andproportions of elements appear in the product orproducts. After the most complicated moleculehas been dealt with, the second-most complexcan then be addressed, and so on.

Assuming the numbers of atoms in the reac-tant and product do not match, it will be neces-sary to place coefficients before one or morechemical species. After this has been done, theequation should again be checked, because as wehave seen, the use of a coefficient to straightenout one discrepancy may create another. Notethat only coefficients can be changed; the formu-las of the species themselves (assuming they werecorrect to begin with) should not be changed.

After the equation has been fully balanced,one final step is necessary. The coefficients mustbe checked to ensure that the smallest integerspossible have been used. Suppose, in the aboveexercise, we had ended up with an equation thatlooked like this: 12H2(g) + 6O2(g) 412H2O(l).This is correct, but not very “clean.” Just as a frac-tion such as 12/24 needs to be reduced to its sim-plest form, 1/2, the same is true of a chemicalequation. The coefficients should thus always bethe smallest number that can be used to yield acorrect result.

Types of Chemical Reactions

Note that in chemical equations, one of the sym-bols used is (aq), which indicates a chemicalspecies that has been dissolved in water—that is,an aqueous solution. The fact that this has itsown special symbol indicates that aqueous solu-tions are an important part of chemistry. Exam-ples of reactions in aqueous solutions are dis-cussed, for instance, in the essays on Acid-BaseReactions; Chemical Equilibrium; Solutions.



ChemicalReactions

Another extremely important type of reac-tion is an oxidation-reduction reaction. Some-times called a redox reaction, an oxidation-reduction reaction occurs during the transfer ofelectrons. The rusting of iron is an example of anoxidation-reduction reaction; so too is combus-tion. Indeed, combustion reactions—in whichoxygen produces energy so rapidly that a flameor even an explosion results—are an importantsubset of oxidation-reduction reactions.

REACTIONS THAT FORM WA-

TER, SOLIDS, OR GASES. Anothertype of reaction is an acid-base reaction, in whichan acid is mixed with a base, resulting in the for-mation of water along with a salt.

Other reactions form gases, as for instancewhen water is separated into hydrogen and oxy-gen. Similarly, heating calcium carbonate (lime-stone) to make calcium oxide or lime for cementalso yields gaseous carbon dioxide: CaCO3(s) +heat 4CaO(s) + CO2(g).

There are also reactions that form a solid,such as the one mentioned much earlier, inwhich solid BaCrO4(s) is formed. Such reactionsare called precipitation reactions. But this is alsoa reaction in an aqueous solution, and there isanother product: 2KNO3(aq), or potassiumnitrate dissolved in water.

SINGLE AND DOUBLE DIS-

PLACEMENT. The reaction referred to inthe preceding paragraph also happens to be anexample of another type of reaction, because twoanions (negatively charged ions) have beenexchanged. Initially K+ and CrO4

2- were together,and these reacted with a compound in whichBa2+ and NO3

-were combined. The anionschanged places, an instance of a double-displace-ment reaction, which is symbolized thus: AB +CD 4AD + CB.

It is also possible to have a single-displace-ment reaction, in which an element reacts with acompound, and one of the elements in the com-pound is released as a free element. This can berepresented symbolically as A + BC 4B + AC.Single-displacement reactions often occur withmetals and with halogens. For instance, a metal(A) reacts with an acid (BC) to produce hydro-gen (B) and a salt (AC).

COMBINATION AND DECOM-

POSITION. A synthesis, or combination,reaction is one in which a compound is formedfrom simpler materials—whether those materi-

als be elements or simple compounds. A basicexample of this is the reaction described earlierin relation to chemical equations, when hydro-gen and oxygen combine to form water. On theother hand, some extremely complex substances,such as the polymers in plastics and syn-thetic fabrics such as nylon, also involve synthesisreactions.

When iron rusts (in other words, it oxidizesin the presence of air), this is both an oxidation-reduction and a synthesis reaction. This also rep-resents one of many instances in which the lan-guage of science is quite different from everydaylanguage. If a piece of iron—say, a railing on abalcony—rusts due to the fact that the paint haspeeled off, it would seem from an unscientificstandpoint that the iron has “decomposed.”However, rust (or rather, metal oxide) is a morecomplex substance than the iron, so this is actu-ally a synthesis or combination reaction.

A true decomposition reaction occurs whena compound is broken down into simpler com-pounds, or even into elements. When water issubjected to electrolysis such that the hydrogenand oxygen are separated, this is a decompositionreaction. The fermentation of grapes to makewine is also a form of decomposition.

And then, of course, there are the processesthat normally come to mind when we think of“decomposition”: the decay or rotting of a for-merly living thing. This could also include thedecay of something, such as an item of food,made from a formerly living thing. In suchinstances, an organic substance is eventually bro-ken down through a number of processes, mostnotably the activity of bacteria, until it ultimate-ly becomes carbon, nitrogen, oxygen, and otherelements that are returned to the environment.

SOME OTHER PARAMETERS.

Obviously, there are numerous ways to classifychemical reactions. Just to complicate things alittle more, they can also be identified as towhether they produce heat (exothermic) orabsorb heat (endothermic). Combustion is clear-ly an example of an exothermic reaction, whilean endothermic reaction can be exemplified bythe process that takes place in a cold pack. Usedfor instance to prevent swelling on an injuredankle, a cold pack contains an ampule thatabsorbs heat when broken.

Still another way to identify chemical reac-tions is in terms of the phases of matter involved.



ChemicalReactions

We have already seen that some reactions formgases, some solids, and some yield water as one ofthe products. If reactants in one phase of matterproduce a substance or substances in the samephase (liquid, solid, or gas), this is called a homo-geneous reaction. On the other hand, if the reac-tants are in different phases of matter, or if theyproduce a substance or substances that are in adifferent phase, this is called a heterogeneousreaction.

An example of a homogeneous reactionoccurs when gaseous nitrogen combines withoxygen, also a gas, to produce nitrous oxide, or“laughing gas.” Similarly, nitrogen and hydrogencombine to form ammonia, also a gas. But whenhydrogen and oxygen form water, this is a het-erogeneous reaction. Likewise, when a metalundergoes an oxidation-reduction reaction, a gasand a solid react, resulting in a changed form ofthe metal, along with the production of newgases.

Finally, a chemical reaction can be eitherreversible or irreversible. Much earlier, wedescribed how wood experiences combustion,resulting in the production of ash. This is clearlyan example of an irreversible reaction. The atomsin the wood and the air that oxidized it have notbeen destroyed, but it would be impossible to putthe ash back together to make a piece of wood. Bycontrast, the formation of water by hydrogen andoxygen is reversible by means of electrolysis.

KEEPING IT ALL STRAIGHT.

The different classifications of reactions dis-cussed above are clearly not mutually exclusive;they simply identify specific aspects of the samething. This is rather like the many physical char-acteristics that describe a person: gender, height,weight, eye color, hair color, race, and so on. Justbecause someone is blonde, for instance, doesnot mean that the person cannot also be brown-eyed; these are two different parameters that aremore or less independent.

On the other hand, there is some relationbetween these parameters in specific instances:for example, females over six feet tall are rare,simply because women tend to be shorter thanmen. But there are women who are six feet tall, oreven considerably taller. In the same way, it isunlikely that a reaction in an aqueous solutionwill be a combustion reaction—yet it does hap-pen, as for instance when potassium reacts withwater.

Studying Chemical Reactions

Several aspects or subdisciplines of chemistry arebrought to bear in the study of chemical reac-tions. One is stoichiometry (stoy-kee-AH-muh-tree), which is concerned with the relationshipsamong the amounts of reactants and products ina chemical reaction. The balancing of the chemi-cal equation for water earlier in this essay is anexample of basic stoichiometry.

Chemical thermodynamics is the area ofchemistry that addresses the amounts of heatand other forms of energy associated with chem-ical reactions. Thermodynamics is also a branchof physics, but in that realm, it is concerned pure-ly with physical processes involving heat andenergy. Likewise physicists study kinetics, associ-ated with the movement of objects. Chemicalkinetics, on the other hand, involves the study ofthe collisions between molecules that produce achemical reaction, and is specifically concernedwith the rates and mechanisms of reaction.

SPEEDING UP A CHEMICAL

REACTION. Essentially, a chemical reactionis the result of collisions between molecules.According to this collision model, if the collisionis strong enough, it can break the chemical bondsin the reactants, resulting in a rearrangement ofthe atoms to form products. The more the mole-cules collide, the faster the reaction. Increase inthe numbers of collisions can be produced in twoways: either the concentrations of the reactantsare increased, or the temperature is increased. Ineither case, more molecules are colliding.

Increases of concentration and temperaturecan be applied together to produce an even fasterreaction, but rates of reaction can also beincreased by use of a catalyst, a substance thatspeeds up the reaction without participating in iteither as a reactant or product. Catalysts are thusnot consumed in the reaction. One very impor-tant example of a catalyst is an enzyme, whichspeeds up complex reactions in the human body.At ordinary body temperatures, these reactionsare too slow, but the enzyme hastens them along.Thus human life can be said to depend on chemical reactions aided by a wondrous form ofcatalyst.

WHERE TO LEARN MORE

Bender, Hal. “Chemical Reactions.” Clackamas Communi-ty College (Web site). <http://dl.clackamas.cc.or.us/ch104-01/chemical.htm> (June 3, 2001).



ChemicalReactions


ACID-BASE REACTION: A chemical

reaction in which an acid is mixed with a

base, resulting in the formation of water

along with a salt.

AQUEOUS SOLUTIONS: A mixture

of water and any substance that is solvent

in it.

BALANCED EQUATION: A chemical

equation in which the numbers of atoms in

the reactants and those in the product are

equal. In the course of balancing an equa-

tion, coefficients may need to be applied to

one or more of the chemical species

involved; however, the actual formulas of

the species cannot be changed.

CATALYST: A substance that speeds up

a chemical reaction without participating

in it either as a reactant or product.

Catalysts are thus not consumed in the

reaction.

CHEMICAL EQUATION: A represen-

tation of a chemical reaction in which

the chemical symbols on the left stand for

the reactants, and those on the right for the

product or products. On paper, a chemical

equation looks much like a mathematical

one; however, instead of an equals sign, a

chemical equation uses an arrow to show

the direction of the reaction.

CHEMICAL KINETICS: the study of

the rate at which chemical reactions occur.

CHEMICAL REACTION: A process

whereby the chemical properties of a sub-

stance are changed by a rearrangement of

the atoms in the substance.

CHEMICAL SPECIES: A generic term

used for any substance studied in chem-

istry—whether it be an element, com-

pound, mixture, atom, molecule, ion, and

so forth.

CHEMICAL THERMODYNAMICS:

The study of the amounts of heat and

other forms of energy associated with

chemical reactions.

COEFFICIENT: A number used to

indicate the presence of more than one

unit—typically, more than one molecule—

of a chemical species in a chemical equa-

tion. For instance, 2H2O indicates two

water molecules. (Note that 1 is never used

as a coefficient.)

COLLISION MODEL: The theory that

chemical reactions are the result of colli-

sions between molecules that are strong

enough to break bonds in the reactants,

resulting in a rearrangement of atoms to

form a product or products.

DECOMPOSITION REACTION: A

chemical reaction in which a compound is

broken down into simpler compounds, or

even into elements. This is the opposite of

a synthesis or combination reaction.

DOUBLE-DISPLACEMENT REAC-

TION: A chemical reaction in which

the partners in two compounds change

places. This can be symbolized as

AB + CD 4AD + CB. Compare single-

displacement reaction.

ENDOTHERMIC: A term describing a

chemical reaction in which heat is

absorbed or consumed.

EXOTHERMIC: A term describing a


produced.

KEY TERMS


ChemicalReactions


“Catalysis, Separations, and Reactions.” Accelrys (Website). <http://www.accelrys.com/chemicals/catalysis/>(June 3, 2001).

Goo, Edward. “Chemical Reactions” (Web site).<http://www-classes.usc.edu/engr/ms/125/MDA125/reactions/> (June 3, 2001).

Knapp, Brian J. Oxidation and Reduction. Illustrated byDavid Woodroffe. Danbury, CT: Grolier Educational,1998.

Knapp, Brian J. Energy and Chemical Change. Illustratedby David Woodroffe. Danbury, CT: Grolier Educa-tional, 1998.

Newmark, Ann. Chemistry. New York: Dorling Kinders-ley, 1993.

“Periodic Table: Chemical Reaction Data.” WebElements

(Web site). <http://www.webelements.com/webele-

ments/elements/text/periodic-table/chem.html>

(June 3, 2001).

Richards, Jon. Chemicals and Reactions. Brookfield, CT:

Copper Beech Books, 2000.

“Types of Chemical Reactions” (Web site). <http://www.

usoe.k12.ut.us/curr/science/sciber00/8th/matter/

sciber/chemtype.htm> (June 3, 2001).



HETEROGENEOUS: A term describ-

ing a chemical reaction in which the reac-

tants are in different phases of matter (liq-

uid, solid, or gas), or one in which the

product is in a different phase from that of

the reactants.

HOMOGENEOUS: A term describing

a chemical reaction in which the reactants

and the product are all in the same phase of

matter (liquid, solid, or gas).

OXIDATION-REDUCTION REACTION:

A chemical reaction involving the transfer

of electrons.

PRECIPITATION REACTION: A

chemical reaction in which a solid is

formed.

PRODUCT: The substance or sub-

stances that result from a chemical

reaction.

REACTANT: A substance that inter-

acts with another substance in a chemical

reaction, resulting in the formation of a

product.

SINGLE-DISPLACEMENT REAC-

TION: A chemical reaction in which an

element reacts with a compound, and one

of the elements in the compound is

released as a free element. This can be rep-

resented symbolically as A + BC 4B + AC.

Compare double-displacement reaction.

STOICHIOMETRY: The study of the

relationships among the amounts of reac-

tants and products in a chemical reaction.

Producing a balanced equation requires

application of stoichiometry (pronounced

“stoy-kee-AH-muh-tree”).

SYNTHESIS OR COMBINATION

REACTION: A chemical reaction in

which a compound is formed from simpler

materials—either elements or simple com-

pounds. It is the opposite of a decomposi-

tion reaction.

UNBALANCED EQUATION: A chem-

ical equation in which the sum of atoms in

the product or products does not equal the

sum of atoms in the reactants. Initial

observations of a chemical reaction usually

produce an unbalanced equation, which

needs to be analyzed and corrected (by the

use of coefficients) to yield a balanced

equation.




OX IDAT ION -REDUCT IONREACT IONSOxidation-Reduction Reactions

CONCEPTMost people have heard the term “oxidation” atsome point or another, and, from the sound ofthe word, may have developed the impressionthat it has something to do with oxygen. Indeedit does, because oxygen has a tendency to drawelectrons to itself. This tendency, rather than thepresence of oxygen itself, is actually what identi-fies oxidation, defined as a process in which asubstance loses electrons. The oxidation of onesubstance is always accompanied by reduction, orthe gaining of electrons, on the part of anothersubstance—hence the term “oxidation-reductionreaction,” sometimes called a redox reaction. Theworld is full of examples of this highly significantform of chemical reaction. One such example iscombustion, or an even more rapid form of com-bustion, explosion. Likewise the metabolism offood, as well as other biological processes,involves oxidation and reduction reactions. So,too, do a number of processes that take place onthe surfaces of metals: when iron rusts; whencopper turns green; or when aluminum forms acoating of aluminum oxide that prevents it fromrusting. Oxidation-reduction reactions also playa major role in electrochemistry, which has ahighly useful application to daily life in the formof batteries.

HOW IT WORKS

Chemical Reactions

A chemical reaction is a process whereby thechemical properties of a substance are changedby a rearrangement its atoms. The change pro-duced by a chemical reaction is quite differentfrom a purely physical change, which does not

affect the fundamental properties of the sub-stance itself. A piece of copper can be heated,melted, beaten into different shapes, and soforth, yet throughout all those changes, itremains pure copper, an element of the transi-tion metals family.

But suppose a copper roof is exposed to theelements for many years. Copper is famous for itshighly noncorrosive quality, and this, combinedwith its beauty, has made it a favored material foruse in the roofs of imposing buildings. (Becauseit is relatively expensive, few middle-class peopletoday can afford a roof entirely made of copper,but sometimes it is used as a decorative touch—for instance, over the entryway of a house.) Even-tually, however, copper does begin to corrodewhen exposed to air for long periods of time.

Over the years, exposed copper develops athin layer of black copper oxide, and as timepasses, traces of carbon dioxide in the air con-tribute to the formation of greenish copper car-bonate. This explains why the Statue of Liberty,covered in sheets of copper, is green, rather thanhaving the reddish-golden hue of new, uncorrod-ed copper.

EXTERNAL VS. INTERNAL

CHANGE. The preceding paragraphsdescribe two very different phenomena. The firstwas a physical change in which the chemicalproperties of a substance—copper—remainedunaltered. The second, on the other hand,involved a chemical change on the surface of thecopper, as copper atoms bonded with carbon andoxygen atoms in the air to form something dif-ferent from copper. The difference between thesetwo types of changes can be likened to varietiesof changes in a person’s life—an external change


Oxidation-ReductionReactions

on the one hand, and a deeply rooted change onthe other.

A person may move to another house, job,school, or town, yet the person remains the same.Many sayings in the English language express thisfact: for instance, “Wherever you go, there youare,” or “You can take the boy out of the country,but you can’t take the country out of the boy.”Moving is simply a physical change. On the otherhand, if a person changes belief systems, over-comes old feelings (or succumbs to new ones),changes lifestyles in a profound manner, or inany other way changes his or her mind aboutsomething important—this is analogous to achemical change. In these instances, the person,like the surface of the copper described above,has changed not merely in external properties,but in inner composition.

“LEO the Lion Says ‘GER’”

Chemical reactions are addressed in depth with-in the essay devoted to that subject, which dis-cusses—among other subjects—many ways ofclassifying chemical reactions. These varieties ofchemical reaction are not all mutually exclusive,as they relate to different aspects of the reaction.As noted in the review of various reaction types,one of the most significant is an oxidation-reduction reaction (sometimes called a redoxreaction) involving the transfer of electrons.

As its name implies, an oxidation-reductionreaction is really two processes: oxidation, inwhich electrons are lost, and reduction, in whichelectrons are gained. Though these are definedseparately here, they do not occur independent-ly; hence the larger reaction of which each is apart is called an oxidation-reduction reaction.In order to keep the two straight, chemistryteachers long ago developed a useful, if nonsensi-cal, mnemonic device: “LEO the lion says ‘GER’.” LEO stands for “Loss of Electrons, Oxida-tion,” and “GER” means “Gain of Electrons,Reduction.”

Many, though not all, oxidation-reductionreactions involve oxygen. Oxygen combinesreadily with other elements, and in so doing, ittends to grab electrons from those other ele-ments’ atoms. As a result, the oxygen atombecomes an ion (an atom with an electriccharge)—specifically, an anion, or negativelycharged ion.

In interacting with another element, oxygenbecomes reduced, while the other element is oxi-dized to become a cation, or a positively chargedion. This, too, is easy to remember: oxygen itself,obviously, cannot be oxidized, so it must be theone being reduced. But since not all oxidation-reduction reactions involve oxygen, perhaps thefollowing is a better way to remember it. Elec-trons are negatively charged, and the elementthat takes them on in an oxidation-reductionreaction is reduced—just as a person who thinksnegative thoughts are “reduced” if those negativethoughts overcome positive ones.

Oxidation Numbers

An oxidation number (sometimes called an oxi-dation state) is a whole-number integer assignedto each atom in an oxidation-reduction reaction.This makes it easier to keep track of the electronsinvolved, and to observe the ways in which theychange positions. Here are some rules for deter-mining oxidation number.

• 1. The oxidation number for an atom of anelement not combined with other elementsin a compound is always zero.

• 2. For an ion of any element, the oxidationnumber is the same as its charge. Thus asodium ion, which has a charge of +1 and isdesignated symbolically as Na+, has an oxi-dation number of +1.

• 3. Certain elements or families form ions inpredictable ways:

• a. Alkali metals, such as sodium, alwaysform a +1 ion; oxidation number = +1.

• b. Alkaline earth metals, such as magne-sium, always form a +2 ion; oxidation num-ber = +2.

• c. Halogens, such as fluorine, form –1 ions;oxidation number = –1.

• d. Other elements have predictable ways toform ions; but some, such as nitrogen, canhave numerous oxidation numbers.

• 4. The oxidation number for oxygen is –2for most compounds involving covalentbonds.

• 5. When hydrogen is involved in covalentbonds with nonmetals, its oxidation num-ber is +1.

• 6. In binary compounds (compounds withtwo elements), the element having greaterelectronegativity is assigned a negative oxi-dation number that is the same as its charge





when it appears as an anion in ionic com-pounds.

• 7. When a compound is electrically neutral,the sum of its elements’ oxidation states iszero.

• 8. In an ionic chemical species, the sum ofthe oxidation states for its constituent ele-ments must equal the overall charge.

These rules will not be discussed here;rather, they are presented to show some of thecomplexities involved in analyzing an oxidation-reduction reaction from a structural stand-point—that is, in terms of the atomic or molecu-lar reactions. For the most part, we will beobserving oxidation-reductions phenomenologi-cally, or in terms of their outward effects. A goodchemistry textbook should provide a moredetailed review of these rules, along with a tableshowing oxidation numbers of elements andbinary compounds.

Oxidation-reduction reactions are easier tounderstand if they are studied as though theywere two half-reactions. Half the reactioninvolves what happens to the substances andelectrons in the oxidizing portion, while theother half-reaction indicates the activities of sub-stances and electrons in the reduction portion.


Combustion and Explosions

As with any type of chemical reaction, combus-tion takes place when chemical bonds are brokenand new bonds are formed. It so happens thatcombustion is a particularly dramatic type ofoxidation-reduction reaction: whereas we can-not watch iron rust, combustion is a noticeableevent. Even more dramatic is combustion thattakes place at a rate so rapid that it results in anexplosion.

Coal is almost pure carbon, and its combus-tion in air is a textbook example of oxidation-reduction. Although there is far more nitrogenthan oxygen in air (which is a mixture ratherthan a compound), nitrogen is very unreactive atlow temperatures. For this reason, it can be usedto clean empty fuel tanks, a situation in whichthe presence of pure oxygen is extremely danger-ous. In any case, when a substance burns, it isreacting with the oxygen in air.

As one might expect from what has alreadybeen said about oxidation-reduction, the oxygenis reduced while the carbon is oxidized. In termsof oxidation numbers, the oxidation number of

Batteries use oxidation-reduction reactions to produce electrical current. (Lester V. Bergman/Corbis. Reproduced by permission.)


tary application, incidentally, was “Greek fire,”created by the Byzantines in the seventh centuryA.D. A mixture of petroleum, potassium nitrate,and possibly quicklime, Greek fire could burn onwater, and was used in naval battles to destroyenemy ships.

For the most part, however, the range ofactivities to which combustion could be appliedwas fairly narrow until the development of thesteam engine in the period from the late seven-teenth century to the early nineteenth century.The steam engine applied the combustion of coalto the production of heat for boiling water, whichin turn provided the power to run machinery. Bythe beginning of the twentieth century, combus-tion had found a new application in the internalcombustion engine, used to power automobiles.

EXPLOSIONS AND EXPLO-

SIVES. An internal combustion engine doesnot simply burn fuel; rather, by the combinedaction of the fuel injectors (in a modern vehicle),in concert with the pistons, cylinders, and sparkplugs, it actually produces small explosions in themolecules of gasoline. These produce the outputof power necessary to turn the crankshaft, andultimately the wheels.

An explosion, in simple terms, is a sped-upform of combustion. The first explosives wereinvented by the Chinese during the Middle Ages,and these included not only fireworks and explosive rockets, but gunpowder. Ironically,however, China rejected the use of gunpowder inwarfare for many centuries, while Europeanstook to it with enthusiasm. Needless to say, Euro-peans’ possession of firearms aided their con-quest of the Americas, as well as much of Africa,Asia, and the Pacific, during the period fromabout 1500 to 1900.

The late nineteenth and early twentieth cen-turies saw the development of new explosives,such as TNT or trinitrotoluene, a hydrocarbon.Then in the mid-twentieth century came themost fearsome explosive of all: the nuclear bomb.A nuclear explosion is not itself the result of anoxidation-reduction reaction, but of somethingmuch more complex—either the splitting ofatoms (fission) or the forcing together of atomicnuclei (fusion).

Nuclear bombs release far more energy thanany ordinary explosive, but the resulting blastalso causes plenty of ordinary combustion. Whenthe United States dropped atomic bombs on the


carbon jumps from 0 to 4, while that of oxygen isreduced to –2. As they burn, these two form car-bon dioxide or CO2, in which the two –2 chargesof the oxygen atoms cancel out the +4 charge ofthe carbon atom to yield a compound that iselectrically neutral.

COMBUSTION IN HUMAN EX-

PERIENCE. Combustion has been a signifi-cant part of human life ever since our prehistoricancestors learned how to harness the power offire to cook food and light their caves. We tend tothink of premodern times—to use the memo-rable title of a book by American historianWilliam Manchester, about the Middle Ages—asA World Lit Only By Fire. In fact, our modern ageis even more combustion-driven than that of ourforebears.

For centuries, burning animal fat—in torch-es, lamps, and eventually in candles—providedlight for humans. Wood fires supplied warmth, aswell as a means to cook meals. These were themain uses of combustion, aside from the occa-sional use of fire in warfare or for other purpos-es (including that ghastly medieval form of exe-cution, burning at the stake). One notable mili-

Oxidation-reduction reactions fuel thespace-shuttle at take-off. (Corbis. Reproduced bypermission.)




Japanese cities of Hiroshima and Nagasaki in

August 1945, those cities suffered not only the

effects of the immediate blast, but also massive

fires resulting from the explosion itself.

FUELING THE SPACE SHUT-

TLE. Oxidation-reduction reactions also fuel

the most advanced form of transportation

known today, the space shuttle. The actual

orbiter vehicle is relatively small compared to its

external power apparatus, which consists of two

solid rocket boosters on either side, along with anexternal fuel tank.

Inside the solid rocket boosters are ammoni-um perchlorate (NH4ClO4) and powdered alu-minum, which undergo an oxidation-reductionreaction that gives the shuttle enormousamounts of extra thrust. As for the larger singleexternal fuel tank, this contains the gases thatpower the rocket: hydrogen and oxygen.

Because these two are extremely explosive,they must be kept in separate compartments.

If not polished regularly, silver responds to oxidation by tarnishing—forming a surface ofsilver sulfide, or Ag2S. (Mimmo Jodice/Corbis. Reproduced by permission.)





When they react, they form water, of course, butin doing so, they also release vast quantities ofenergy. The chemical equation for this is: 2H2 +O2 42H2O + energy.

On January 28, 1986, something went terri-bly wrong with this arrangement on the spaceshuttle Challenger. Cold weather had fatigued theO-rings that sealed the hydrogen and oxygencompartments, and the gases fed straight into theflames behind the shuttle itself. This produced apowerful and uncontrolled oxidation-reductionreaction, an explosion that took the lives of allseven astronauts aboard the shuttle.

The Environment and Human Health

Combustion, though it can do much good, canalso do much harm. This goes beyond the obvi-ous: by burning fossil fuels or hydrocarbons,excess carbon (in the form of carbon dioxide andcarbon monoxide) is released to the atmosphere,with a damaging effect on the environment.

In fact, oxidation-reduction reactions areintimately connected with the functioning of thenatural environment. For example, photosynthe-sis, the conversion of light to chemical energy byplants, is a form of oxidation-reduction reactionthat produces two essentials of human life: oxy-gen and carbohydrates. Likewise cellular respira-tion, which along with photosynthesis is dis-cussed in the Carbon essay, is an oxidation-reduction reaction in which living things breakdown molecules of food to produce energy, car-bon dioxide, and water.

Enzymes in the human body regulate oxida-tion-reduction reactions. These complex pro-teins, of which several hundred are known, act ascatalysts, speeding up chemical processes in thebody. Oxidation-reduction reactions also takeplace in the metabolism of food for energy, withsubstances in the food broken down into compo-nents the body can use.

OXIDATION: SPOILING AND

AGING. At the same time, oxidation-reduc-tion reactions are responsible for the spoiling offood, the culprit here being the oxidation portionof the reaction. To prevent spoilage, manufactur-ers of food items often add preservatives, whichact as reducing agents.

Oxidation may also be linked with the effectsof aging in humans, as well as with other condi-

tions such as cancer, hardening of the arteries,and rheumatoid arthritis. It appears that oxygenmolecules and other oxidizing agents, alwayshungry for electrons, extract these from themembranes in human cells. Over time, this cancause a gradual breakdown in the body’simmune system.

To forestall the effects of oxidation, somedoctors and scientists recommend antioxi-dants—natural reducing agents such as vitaminC and vitamin E. The vitamin C in lemon juicecan be used to prevent oxidizing on the cut sur-face of an apple, to keep it from turning brown.Perhaps, some experts maintain, natural reduc-ing agents can also slow the pace of oxidation inthe human body.

Forming a New Surface on Metal

Clearly, oxidization can have a corrosive effect,and nowhere is this more obvious than in thecorrosion of metals by exposure to oxidizingagents—primarily oxygen itself. Most metalsreact with O2, and might corrode so quickly thatthey become useless, were it not for the forma-tion of a protective coating—an oxide.

Iron forms an oxide, commonly known asrust, but this in fact does little to protect it fromcorrosion, because the oxide tends to flake off,exposing fresh surfaces to further oxidation.Every year, businesses and governments devotemillions of dollars to protecting iron and steelfrom oxidation by means of painting and othermeasures, such as galvanizing with zinc. In fact,oxidation-reduction reactions virtually definethe world of iron. Found naturally only in ores,the element is purified by heating the ore withcoke (impure carbon) in the presence of oxygen,such that the coke reduces the iron.

COINAGE METALS. Copper, as wehave seen, responds to oxidation by corroding ina different way: not by rusting, but by changingcolor. A similar effect occurs in silver, which tar-nishes, forming a surface of silver sulfide, orAg2S. Copper and silver are two of the “coinagemetals,” so named because they have often beenused to mint coins. They have been used for thispurpose not only because of their beauty, butalso due to their relative resistance to corrosion.This resistance has, in fact, earned them the nick-name “noble metals.”



The third member of this mini-family is

gold, which is virtually noncorrosive. Wonderful

as gold is in this respect, however, no one is like-

ly to use it as a roofing material, or for any such

large-scale application involving its resistance to

oxidation. Aside from the obvious expense, gold

is soft, and not very good for structural uses, even

if it were much cheaper. Yet there is such a “won-

der metal”: one that experiences virtually no cor-

rosion, is cheap, and strong enough in alloys

to be used for structural purposes. Its name is

aluminum.

ALUMINUM. There was a time, in fact,

when aluminum was even more expensive than

gold. When the French emperor Napoleon III

wanted to impress a dinner guest, he arranged for

the person to be served with aluminum utensils,

while less distinguished personages had to settle

for “ordinary” gold and silver.

In 1855, aluminum sold for $100,000 a

pound, whereas in 1990, the going rate was about

$0.74. Demand did not go down—in fact, it

increased exponentially—but rather, supply

increased, thanks to the development of an inex-


ANION: An ion with a negative charge;

pronounced “AN-ie-un.”

BINARY COMPOUND: A compound

involving two elements.









so forth.




ELECTROCHEMISTRY: The study of

the relationship between chemical and

electrical energy.


tric current to produce a chemical change.



electrons.

CATION: An ion with a positive charge;

pronounced “KAT-ie-un.”


has acquired an electric charge due to the

loss or gain of electrons.

OXIDATION: A chemical reaction in

which a substance loses electrons. It is

always accompanied by reduction; hence

the term oxidation-reduction reaction.

OXIDATION NUMBER: A number

assigned to each atom in an oxidation-

reduction reaction as a means of keeping

track of the electrons involved. Another

term for oxidation number is “oxidation

state.”

OXIDATION-REDUCTION REACTION:

A chemical reaction involving the transfer

of electrons.

REDOX REACTION: Another name

for an oxidation-reduction reaction.

REDUCTION: A chemical reaction in

which a substance gains electrons. It is

always accompanied by oxidation; hence

the term oxidation-reduction reaction.





KEY TERMS



pensive aluminum-reduction process. Two men,one American and one French, discovered thisprocess at the same time: interestingly, their yearsof birth and death were the same.

Aluminum was once a precious metalbecause it proved extremely difficult to separatefrom oxygen. The Hall-Heroult process overcamethe problem by applying electrolysis—the use ofan electric current to produce a chemicalchange—as a way of reducing Al3+ ions (whichhave a high affinity for oxygen) to neutral alu-minum atoms. In the United States today, 4.5%of the total electricity output is used for the pro-duction of aluminum through electrolysis.

The foregoing statistic is staggering, consid-ering just how much electricity Americans use,and it indicates the importance of this once-pre-cious metal. Actually, aluminum oxidizes just likeany other metal—and does so quite quickly, as amatter of fact, by forming a coating of aluminumoxide (Al2O3). But unlike rust, the aluminumoxide is invisible, and acts as a protective coating.Chromium, nickel, and tin react to oxygen in asimilar way, but these are not as inexpensive asaluminum.

Electrochemistry and Batteries

Electrochemistry is the study of the relationshipbetween chemical and electrical energy. Amongits applications is the creation of batteries, whichuse oxidation-reduction reactions to produce anelectric current.

A basic battery can be pictured schematical-ly as two beakers of solution connected by a wire.In one solution is the oxidizing agent; in theother, a reducing agent. The wire allows electronsto pass back and forth between the two solutions,but to ensure that the flow goes both ways, thetwo solutions are also connected by a “saltbridge.” The salt bridge contains a gel or solutionthat permits ions to pass back and forth, but aporous membrane prevents the solutions fromactually mixing.

In the lead storage battery of an automobile,lead itself is the reducing agent, while lead (IV)

oxide (PbO2) acts as the oxidizing agent. A high-ly efficient type of battery, able to withstand wideextremes in temperature, the lead storage batteryhas been in use since 1915. Along the way, fea-tures have been altered, but the basic principleshave remained—a testament to the soundness ofits original design.

The batteries people use for powering allkinds of portable appliances, from flashlights toboom boxes, are called dry cell batteries. In con-trast to the model described above, using solu-tions, a dry cell (as its name implies) involves noliquid components. Instead, it utilizes variouselements in a range of combinations, includingzinc, magnesium, mercury, silver, nickel, andcadmium. The last two are applied in the nickel-cadmium battery, which is particularly usefulbecause it can be recharged over and over againby an external current. The current turns theproducts of the chemical reactions in the batteryback into reactants.

WHERE TO LEARN MORE

“Batteries.” Oregon State University Department of Chem-istry (Web site). <http://www.chem.orst.edu/ch411/scbatt.htm> (June 4, 2001).

“Batteries and Fuel Cells” (Web site). <http://vectorsite.com/ttfuelc.html> (June 4, 2001).

Borton, Paula and Vicky Cave. The Usborne Book of Bat-teries and Magnets. Tulsa, OK: EDC Publishing, 1995.

Craats, Rennay. The Science of Fire. Milwaukee, WI:Gareth Stevens Publishing, 2000.

Knapp, Brian J. Oxidation and Reduction. Danbury, CT:Grolier Educational, 1998.

“Oxidation Reduction Links” (Web site). <http://users.erols.com/merosen/redox.htm> (June 4, 2001).

“Oxidation-Reduction Reactions.” General Chemistry(Web site). <http://ull.chemistry.uakron.edu/genchem/11/> (June 4, 2001).

“Oxidation-Reduction Reactions: Redox.” UNC-ChapelHill Chemistry Fundamentals (Web site).<http://www.shodor.org/unchem/advanced/redox/>(June 4, 2001).

Yount, Lisa. Antoine Lavoisier: Founder of Modern Chem-istry. Springfield, NJ: Enslow Publishers, 1997.

Zumdahl, Steven S. Introductory Chemistry: A Founda-tion, fourth edition. Boston: Houghton Mifflin, 2000.




CHEM ICAL EQU I L IBR IUMChemical Equilibrium

CONCEPTReactions are the “verbs” of chemistry—theactivity that chemists study. Many reactionsmove to their conclusion and then stop, meaningthat the reactants have been completely trans-formed into products, with no means of return-ing to their original state. In some cases, the reac-tion truly is irreversible, as for instance whencombustion changes both the physical andchemical properties of a substance. There areplenty of other circumstances, however, in whicha reverse reaction is not only possible but anongoing process, as the products of the first reac-tion become the reactants in a second one. Thisdynamic state, in which the concentrations ofreactants and products remains constant, isreferred to as equilibrium. It is possible to predictthe behavior of substances in equilibriumthrough the use of certain laws, which are appliedin industries seeking to lower the costs of pro-ducing specific chemicals. Equilibrium is alsouseful in understanding processes that pre-serve—or potentially threaten—human health.

HOW IT WORKS

Chemical Reactions in Brief

What follows is a highly condensed discussion ofchemical reactions, and particularly the methodsfor writing equations to describe them. For amore detailed explanation of these principles, thereader is encouraged to consult the ChemicalReactions essay.

A chemical reaction is a process whereby thechemical properties of a substance are altered bya rearrangement of the atoms in the substance.

The changes produced by a chemical reaction arefundamentally different from physical changes,such as boiling or melting liquid water, changesthat alter the physical properties of water withoutaffecting its molecular structure.

INDICATIONS THAT A CHEMI-

CAL REACTION HAS OCCURRED.

Though chemical reactions are most effectivelyanalyzed in terms of molecular properties andbehaviors, there are numerous indicators thatsuggest to us when a chemical reaction hasoccurred. It is unlikely that all of these will resultfrom any one reaction, and in fact chances arethat a particular reaction will manifest only oneor two of these effects. Nonetheless, these offer ushints that a reaction has taken place:

Signs that a substance has undergone achemical reaction:

• Water is produced

• A solid forms

• Gases are produced

• Bubbles are formed

• There is a change in color

• The temperature changes

• The taste of a consumable substancechanges

• The smell changes

CHEMICAL CHANGES CON-

TRASTED WITH PHYSICAL CHANG-

ES OF TEMPERATURE. Many of theseeffects can be produced simply by changing thetemperature of a substance, but again, the mereact of applying heat from outside (or removingheat from the substance itself) does not consti-tute a chemical change. Water can be “produced”by melting ice, but the water was already there—it only changed form. By contrast, when an acid


ChemicalEquilibrium

and a base react to form water and a salt, that is atrue chemical reaction.

Similarly, the freezing of water forms a solid,but no new substance has been formed; in achemical reaction, by contrast, two liquids canreact to form a solid. When water boils throughthe application of heat, bubbles form, and a gasor vapor is produced; yet in chemical changes,these effects are not the direct result of applyingheat.

In this context, a change in temperature,noted as another sign that a reaction has takenplace, is a change of temperature from within thesubstance itself. Chemical reactions can be classi-fied as heat-producing (exothermic) or heat-absorbing (endothermic). In either case, thetransfer of heat is not accomplished simply bycreating a temperature differential, as wouldoccur if heat were transferred merely throughphysical means.

Why Do Chemical ReactionsOccur?

At one time, chemists could only study reactionsfrom “outside,” as it were—purely in terms ofeffects noticeable through the senses. Between

the early nineteenth and the early twentieth cen-turies, however, the entire character of chemistrychanged, as did the terms in which chemists dis-cussed reactions. Today, those reactions are ana-lyzed primarily in terms of subatomic, atomic,and molecular properties and activities.

Despite all this progress, however, chemistsstill do not know exactly what happens in achemical reaction—but they do have a goodapproximation. This is the collision model,which explains chemical reactions in terms ofcollisions between molecules. If the collision isstrong enough, it can break the chemical bondsin the reactants, resulting in a re-formation ofatoms within different molecules. The more themolecules collide, the more bonds are being bro-ken, and the faster the reaction.

Increase in the numbers of collisions can beproduced in two ways: either the concentrationsof the reactants are increased, or the temperatureis increased. By raising the temperature, thespeeds of the molecules themselves increase, andthe collisions possess more energy. A certainenergy threshold, the activation energy (symbol-ized Ea) must be crossed in order for a reaction tooccur. A temperature increase raises the likeli-hood that a given collision will cross the activa-tion-energy threshold, producing the energy tobreak the molecular bonds and promote thechemical reaction.

Raising the temperature and the concentra-tions of reactants can increase the energy andhasten the reactions, but in some cases it is notpossible to do either. Fortunately, the rate ofreaction can be increased in a third way, throughthe introduction of a catalyst, a substance thatspeeds up the reaction without participating in iteither as a reactant or product. Catalysts are thusnot consumed in the reaction. Many chemistrytextbooks discuss catalysts within the context ofequilibrium; however, because catalysts play suchan important role in human life, in this bookthey are the subject of a separate essay.

Chemical Equations InvolvingEquilibrium

A chemical equation, like a mathematical equa-tion, symbolizes an interaction between entitiesthat produces a particular result. In the case of achemical equation, the entities are not numbersbut reactants, and they interact with each othernot through addition or multiplication, but by


MANY MOUNTAIN CLIMBERS USE PRESSURIZED OXYGEN

AT VERY HIGH ALTITUDES TO MAINTAIN PROPER HEMO-GLOBIN-OXYGEN EQUILIBRIUM. (Galen Rowell/Corbis. Repro-



ChemicalEquilibrium

chemical reaction. Yet just as a product is theresult of multiplication in mathematics, a prod-uct in a chemical equation is the substance orsubstances that result from the reaction.

Instead of an equals sign between the reac-tants and the product, an arrow is used. Whenthe arrow points to the right, this indicates a for-ward reaction; conversely, an arrow pointing tothe left symbolizes a reverse reaction. In a reversereaction, the products of a forward reaction havebecome the reactants, and the reactants of theforward reaction are now the products. This isindicated by an arrow that points toward the left.

Chemical equilibrium, which occurs whenthe ratio between the reactants and products isconstant, and in which the forward and reversereactions take place at the same rate, is symbol-ized thus: 1. Note that the arrows, the upper onepointing right and the lower one pointing left,are of the same length. There may be certaincases, discussed below, in which it is necessary toshow these arrows as unequal in length as ameans of indicating the dominance of either theforward or reverse reaction.

Chemical equations usually include notationindicating the state or phase of matter for thereactants and products: (s) for a solid; (l) for aliquid; (g) for a gas. A fourth symbol, (aq), indi-cates a substance dissolved in water—that is, anaqueous solution. In the following paragraphs,we will apply a chemical equation to the demon-stration of equilibrium, but will not discuss thebalancing of equations. The reader is encouragedto consult the passage in the Chemical Reactionsessay that addresses that process, vital to therecording of accurate data.

A SIMPLE EQUILIBRIUM EQUA-

TION. Let us now consider a simple equation,involving the reaction between water and carbonmonoxide (CO) at high temperatures in a closedcontainer. The initial equation is written thus:H2O(g) + CO(g) 4H2(g) + CO2g). In plain Eng-lish, water in the gas phase (steam) has reactedwith carbon monoxide to produce hydrogen gasand carbon dioxide.

As the reaction proceeds, the amount ofreactants decreases, and the concentration ofproducts increases. At some point, however, therewill be a balance between the numbers of prod-ucts and reactants—a state of chemical equilibri-um represented by changing the right-pointingarrow to an equilibrium symbol: H2O(g) +

CO(g) ! H2(g) + CO2g). Assuming that the sys-tem is not disturbed (that is, that the container iskept closed and no outside substances are intro-duced), equilibrium will continue to be main-tained, because the reverse reaction is occurringat the same rate as the forward one.

Note what has been said here: reactions arestill occurring, but the forward and rearwardreactions balance one another. Thus equilibriumis not a static condition, but a dynamic one, andindeed, chemical equilibrium is sometimesreferred to as “dynamic equilibrium.” On theother hand, some chemists refer to chemicalequilibrium simply as equilibrium, but here thequalifier chemical has been used to distinguishthis from the type of equilibrium studied inphysics. Physical equilibrium, which involves fac-tors such as center of gravity, does help us tounderstand chemical equilibrium, but it is a dif-ferent phenomenon.


Homogeneous and Heteroge-neous Equilibria

It should be noted that the equation used aboveidentifies a situation of homogeneous equilibri-um, in which all the substances are in the samephase or state of matter—gas, in this case. It isalso possible to achieve chemical equilibrium in areaction involving substances in more than onephase of matter.

An example of such heterogeneous equilib-rium is the decomposition of calcium carbonatefor the production of lime, a process that involvesthe application of heat. Here the equation wouldbe written thus: CaCO3(s) ! CaO(s) + CO2(g).Both the calcium carbonate (CaCO3) and thelime (CaO) are solids, whereas the carbon diox-ide produced in this reaction is a gas.

The Equilibrium Constant

In 1863, Norwegian chemists Cato MaximilianGuldberg (1836-1902) and Peter Waage (1833-1900)—who happened to be brothers-in-law—formulated what they called the law of massaction. Today, this is called the law of chemicalequilibrium, which states that the direction takenby a reaction is dependant not merely on themass of the various components of the reaction,



ChemicalEquilibrium

but also upon the concentration—that is, themass present in a given volume.

This can be expressed by the formula aA +bB 1 cC +dD, where the capital letters representchemical species, and the italicized lowercase let-ters indicate their coefficients. The equation[C]c[D]d/[A]a[B]b yields what is called an equi-librium constant, symbolized K.

The above formula expresses the equilibri-um constant in terms of molarity, the amount ofsolute in a given volume of solution, but in thecase of gaseous reactants and products, the equi-librium constant can also be expressed in termsof partial pressures. In the reaction of water andcarbon monoxide to produce hydrogen mole-cules and carbon dioxide (H2O + CO ! H2 +CO2). In chemical reactions involving solids,however, the concentration of the solid—becauseit is considered to be invariant—does not appearin the equilibrium constant. In the reactiondescribed earlier, in which calcium carbonatewas in equilibrium with solid lime and gaseouscarbon dioxide, K = pressure of CO2.

We will not attempt here to explore the equi-librium constant in any depth, but it is importantto recognize its usefulness. For a particular reac-tion at a specific temperature, the ratio of con-centrations between reactants and products willalways have the same value—the equilibriumconstant, or K. Because it is not dependant on theamounts of reactants and products mixedtogether initially, K remains the same: the con-centrations themselves may vary, but the ratiosbetween the concentrations in a given situationdo not.

Le Châtelier’s Principle

Not all situations of equilibrium are alike:depending on certain factors, the position ofequilibrium may favor one side of the equationor the other. If a company is producing chemicalsfor sale, for example, its production managerswill attempt to influence reactions in such a wayas to favor the forward reaction. In such a situa-tion, it is said that the equilibrium position hasbeen shifted to the right. In terms of physicalequilibrium, mentioned above, this would beanalogous to what would happen if you wereholding your arms out on either side of yourbody, with a heavy lead weight in your left handand a much smaller weight in the right hand.

Your center of gravity, or equilibrium position,would shift to the left to account for the greaterforce exerted by the heavier weight.

A value of K significantly above 1 causes ashift to the right, meaning that at equilibrium,there will be more products than reactants. Thisis a situation favorable to a chemical company’smanagers, who desire to create more of the prod-uct from less of the reactants. However, natureabhors an imbalance, as expressed in Le Châte-lier’s principle. Named after French chemistHenri Le Châtelier (1850-1936), this principlemaintains that whenever a stress or change isimposed on a chemical system in equilibrium,the system will adjust the amounts of the varioussubstances to reduce the impact of that stress.

Suppose we add more of a particular sub-stance to increase the rate of the forward reac-tion. In an equation for this reaction, the equilib-rium symbol is altered, with a longer arrowpointing to the right to indicate that the forwardreaction is favored. Again, the equilibrium posi-tion has shifted to the right—just as one makesphysical adjustments to account for an imbal-anced weight. The system responds by workingto consume more of the reactant, thus adjustingto the stress that was placed on it by the additionof more of that substance. By the same token, ifwe were to remove a particular reactant or prod-uct, the system would shift in the direction of thedetached component.

Note that Le Châtelier’s principle is mathe-matically related to the equilibrium constant.Suppose we have a basic equilibrium equation ofA + B 1 C, with A and B each having molaritiesof 1, and C a molarity of 4. This tells us that K isequal to the molarity of C divided by that of Amultiplied by B = 4/(1 • 1). Suppose, now, thatenough of C were added to bring its concentra-tion up to 6. This would mean that the systemwas no longer at equilibrium, because C/(A • B)no longer equals 4. In order to return the ratio to4, the numerator (C) must be decreased, whilethe denominator (A • B) is increased. The reac-tion thus shifts from right to left.

CHANGES IN VOLUME AND

TEMPERATURE. If the volume of gases ina closed container is decreased, the pressureincreases. An equilibrium system will thereforeshift in the direction that reduces the pressure;but if the volume is increased, thus reducing thepressure, the system will respond by shifting to



increase pressure. Note, however, that not allincreases in pressure lead to a shift in the equilib-rium. If the pressure were increased by the addi-tion of a noble gas, the gas itself—since these ele-ments are noted for their lack of reactivity—would not be part of the reaction. Thus thespecies added would not be part of the equilibri-um constant expression, and there would be nochange in the equilibrium.

In any case, no change in volume alters theequilibrium constant K; but where changes intemperature are involved, K is indeed altered. Inan exothermic, or heat-producing reaction, theheat is treated as a product. Thus, when nitrogenand hydrogen react, they produce not onlyammonia, but a certain quantity of heat. If thissystem is at equilibrium, Le Châtelier’s principleshows that the addition of heat will induce a shiftin equilibrium to the left—in the direction thatconsumes heat or energy.

The reverse is true in an endothermic, orheat-absorbing reaction. As in the processdescribed earlier, the thermal decomposition ofcalcium carbonate produces lime and carbondioxide. Because heat is used to cause this reac-tion, the amount of heat applied is treated as areactant, and an increase in temperature willcause the equilibrium position to shift to theright.

Equilibrium and Health

Discussions of chemical equilibrium tend tobe rather abstract, as the foregoing sections onthe equilibrium constant and Le Châtelier’s prin-ciple illustrate. (The reader is encouraged to con-sult additional sources on these topics, whichinvolve a number of particulars that have beentouched upon only briefly here.) Despite thechallenges involved in addressing the subject ofequilibrium, the results of chemical equilibriumcan be seen in processes involving human health.

The cooling of food with refrigerators, alongwith means of food preservation that do notinvolve changes in temperature, maintains chem-ical equilibrium in the foods and thereby pre-vents or at least retards spoilage. Even moreimportant is the maintenance of equilibrium inreactions between hemoglobin and oxygen inhuman blood.

HEMOGLOBIN AND OXYGEN.

Hemoglobin, a protein containing iron, is the

material in red blood cells responsible for trans-porting oxygen to the cells. Each hemoglobinmolecule attaches to four oxygen atoms, and theequilibrium conditions of the hemoglobin-oxy-gen interaction can be expressed thus: Hb(aq) +4O2(g) ! Hb(O2)4(aq), where “Hb” stands forhemoglobin. As long as there is sufficient oxygenin the air, a healthy equilibrium is maintained; butat high altitudes, considerable changes occur.

At significant elevations above sea level, theair pressure is lowered, and thus it is more diffi-cult to obtain the oxygen one needs. The result,in accordance with Le Châtelier’s principle, is ashift in equilibrium to the left, away from theoxygenated hemoglobin. Without adequate oxy-gen fed to the body’s cells and tissues, a persontends to feel light-headed.

When someone not physically prepared forthe change is exposed to high altitudes, it may benecessary to introduce pressurized oxygen froman oxygen tank. This shifts the equilibrium to theright. For people born and raised at high alti-tudes, however, the body’s chemistry performsthe equilibrium shift—by producing morehemoglobin, which also shifts equilibrium to theright.

HEMOGLOBIN AND CARBON

MONOXIDE. When someone is exposed tocarbon monoxide gas, a frightening variation onthe normal hemoglobin-oxygen interactionoccurs. Carbon monoxide “fools” hemoglobininto mistaking it for oxygen because it also bondsto hemoglobin in groups of four, and the equi-librium expression thus becomes: Hb(aq) +4CO(g) ! Hb(CO)4(aq). Instead of hemoglobin,what has been produced is called carboxyhemo-globin, which is even redder than hemoglobin.Therefore, one sign of carbon monoxide poison-ing is a flushed face.

The bonds between carbon monoxide andhemoglobin are about 300 times as strong asthose between hemoglobin and oxygen, and thismeans a shift in equilibrium toward the right sideof the equation—the carboxyhemoglobin side. Italso means that K for the hemoglobin-carbonmonoxide reaction is much higher than for thehemoglobin-oxygen reaction. Due to the affinityof hemoglobin for carbon monoxide, the hemo-globin puts a priority on carbon monoxidebonds, and hemoglobin that has bonded withcarbon monoxide is no longer available to carryoxygen.


ChemicalEquilibrium



ACTIVATION ENERGY: The minimal

energy required to convert reactants into

products, symbolized Ea.

AQUEOUS SOLUTION: A mixture of

water and a substance that is dissolved in it.


a chemical reaction without participating

in it, either as a reactant or product.


reaction.

CHEMICAL EQUATION: A represen-

tation of a chemical reaction in which the

chemical symbols on the left stand for the

reactants, and those on the right are the

product or products.

CHEMICAL EQUILIBRIUM: A situa-

tion in which the ratio between the reac-

tants and products in a chemical reaction is

constant, and in which the forward reac-

tions and reverse reactions take place at the

same rate.









so forth.

COEFFICIENT: a number used to indi-

cate the presence of more than one unit—

typically, more than one molecule—of a

chemical species in a chemical equation.



sions between molecules that are strong

enough to break bonds in the reactants,

resulting in a reformation of atoms.

ENDOTHERMIC: A term describing a


absorbed or consumed.

KEY TERMSChemicalEquilibrium

Carbon monoxide in small quantities can

cause headaches and dizziness, but larger con-

centrations can be fatal. To reverse the effects of

the carbon monoxide, pure oxygen must be

introduced to the body. It will react with the car-

boxyhemoglobin to produce properly oxygenat-

ed hemoglobin, along with carbon monoxide:

Hb(CO)4(aq) + 4O2(g) ! Hb(O2)4(aq) +

4CO(g). The gaseous carbon monoxide thus pro-

duced is dissipated when the person exhales.

WHERE TO LEARN MORE

“Catalysts” (Web site). <http://edie.cprost.sfu.ca/~rhlogan/catalyst.html> (June 9, 2001).


“Chemical Equilibrium.” Davidson College Department ofChemistry (Web site). <http://www.chm.davidson.edu/ronutt/che115/EquKin.htm> (June 9, 2001).

“Chemical Equilibrium in the Gas Phase.” Virginia TechChemistry Department (Web site). <http://www.chem.vt.edu/RVGS/ACT/notes/chem-eqm.html>(June 9, 2001).

“Chemical Sciences: Mechanism of Catalysis.” University ofAlberta Department of Chemistry (Web site).<http://www.chem.ualberta.ca/~plambeck/che/p102/p02174.htm> (June 9, 2001).


Hauser, Jill Frankel. Super Science Concoctions: 50 Myste-rious Mixtures for Fabulous Fun. Charlotte, VT:Williamson Publishing, 1996.

“Mark Rosen’s Chemical Equilibrium Links” (Web site).<http://users.erols.com/merosen/equilib.htm> (June9, 2001).


ChemicalEquilibrium

Oxlade, Chris. Chemistry. Illustrated by Chris Fair-clough. Austin, TX: Raintree Steck-Vaughn, 1999.



EXOTHERMIC: A term describing a


produced.

FORWARD REACTION: A chemical

reaction symbolized by a chemical equa-

tion in which the reactants and product are

separated by an arrow that points to the

right, toward the products.

HETEROGENEOUS EQUILIBRIUM:

Chemical equilibrium in which the sub-

stances involved are in different phases of

matter (solid, liquid, gas.)

HOMOGENEOUS EQUILIBRIUM:

Chemical equilibrium in which the sub-

stances involved are in the same phase of

matter.

LE CHÂTELIER’S PRINCIPLE: A

statement, formulated by French chemist

Henri Le Châtelier (1850-1936), which

holds that whenever a stress or change is

imposed on a system in chemical equilibri-

um, the system will adjust the amounts of

the various substances in such a way as to

reduce the impact of that stress.



reaction.

REACTANT: A substance that interacts

with another substance in a chemical reac-

tion, resulting in a product.

REVERSE REACTION: A chemical

reaction symbolized by a chemical equa-

tion in which the products of a forward

reaction have become the reactants, and

the reactants of the forward reaction are

now the products. This is indicated by an

arrow that points toward the left.











C ATALYSTSCatalysts

CONCEPTIn most of the processes studied within the phys-ical sciences, the lesson again and again is thatnature provides no “free lunch”; in other words,it is not possible to get something for nothing. Achemical reaction, for instance, involves the cre-ation of substances different from those thatreacted in the first place, but the number ofatoms involved does not change. In view ofnature’s inherently conservative tendencies, then,the idea of a catalyst—a substance that speeds upa reaction without being consumed—seemsalmost like a magic trick. But catalysts are veryreal, and their presence in the human body helpsto sustain life. Similarly, catalysts enable the synthesis of foods, and catalytic converters inautomobiles protect the environment from dan-gerous exhaust fumes. Yet the presence of oneparticular catalyst in the upper atmosphere posessuch a threat to Earth’s ozone layer that produc-tion of certain chemicals containing that sub-stance has been banned.

HOW IT WORKS

Reactions and Collisions

In a chemical reaction, substances known asreactants interact with one another to create newsubstances, called products. In the present con-text, our concern is not with the reactants andproducts themselves, but with an additional enti-ty, an agent that enables the reaction to move for-ward at faster rates and lower temperatures.

According to the collision model generallyaccepted by chemists, chemical reactions are theresult of collisions between molecules. Collisions

that are sufficiently energetic break the chemicalbonds that hold molecules together; as a result,the atoms in those molecules are free to recom-bine with other atoms to form new molecules.Hastening of a chemical reaction can be pro-duced in one of three ways. If the concentrationsof the reactants are increased, this means thatmore molecules are colliding, and potentiallymore bonds are being broken. Likewise if thetemperature is increased, the speeds of the mole-cules themselves increase, and their collisionspossess more energy.

Energy is an important component in thechemical reaction because a certain threshold,called the activation energy (Ea), must be crossedbefore a reaction can occur. A temperatureincrease raises the energy of the collisions,increasing the likelihood that the activation-energy threshold will be crossed, resulting in thebreaking of molecular bonds.

Catalysts and Catalysis

It is not always feasible or desirable, however, toincrease the concentration of reactants, or thetemperature of the system in which the reactionis to occur. Many of the processes that take placein the human body, for instance,“should” requirehigh temperatures—temperatures too high tosustain human life. But fortunately, our bodiescontain proteins called enzymes, discussed laterin this essay, that facilitate the necessary reactionswithout raising temperatures or increasing theconcentrations of substances.

An enzyme is an example of a catalyst, a sub-stance that speeds up a reaction without partici-pating in it either as a reactant or product. Cata-lysts are thus not consumed in the reaction. The


Catalysts

catalyst does its work—catalysis—by creating adifferent path for the reaction, and though themeans whereby it does this are too complex todiscuss in detail here, the process of catalyst canat least be presented in general terms.

Imagine a graph whose x-axis is labeled“reaction progress,” while the y-axis bears thelegend “energy.” There is some value of y equal tothe normal activation energy, and in the courseof experiencing the molecular collisions that leadto a reaction, the reactants reach this level. In acatalyzed reaction, however, the level of activa-tion energy necessary for the reaction is repre-sented by a lower y-value on the graph. The cat-alyzed substances do not need to have as muchenergy as they do without a catalyst, and there-


fore the reaction can proceed more quickly—without changing the temperature or concentra-tions of reactants.


A Brief History of Catalysis

Long before chemists recognized the existence ofcatalysts, ordinary people had been using theprocess of catalysis for a number of purposes:making soap, for instance, or fermenting wine tocreate vinegar, or leavening bread. Early in thenineteenth century, chemists began to take noteof this phenomenon.

CATALYTIC CONVERTERS EMPLOY A CATALYST TO FACILITATE THE TRANSFORMATION OF POLLUTION-CAUSING EXHAUSTS

TO LESS HARMFUL SUBSTANCES. (Ian Harwood; Ecoscene/Corbis. Reproduced by permission.)


Catalysts

In 1812, Russian chemist Gottlieb Kirchhofwas studying the conversion of starches to sugarin the presence of strong acids when he noticedsomething interesting. When a suspension ofstarch in water was boiled, Kirchhof observed, nochange occurred in the starch. However, when headded a few drops of concentrated acid beforeboiling the suspension (that is, particles of starchsuspended in water), he obtained a very differentresult. This time, the starch broke down to formglucose, a simple sugar, while the acid—whichclearly had facilitated the reaction—underwentno change.

Around the same time, English chemist SirHumphry Davy (1778-1829) noticed that in cer-tain organic reactions, platinum acted to speedalong the reaction without undergoing anychange. Later on, Davy’s star pupil, the greatBritish physicist and chemist Michael Faraday(1791-1867), demonstrated the ability of plat-inum to recombine hydrogen and oxygen thathad been separated by the electrolysis of water.The catalytic properties of platinum later foundapplication in catalytic converters, as we shall see.

AN IMPROVED DEFINITION. In1835, Swedish chemist Jons Berzelius (1779-1848) provided a name to the process Kirchhofand Davy had observed from very different per-spectives: catalysis, derived from the Greek words

kata (“down”) and lyein (“loosen.”) As Berzeliusdefined it, catalysis involved an activity quite dif-ferent from that of an ordinary chemical reac-tion. Catalysis induced decomposition in sub-stances, resulting in the formation of new com-pounds—but without the catalyst itself actuallyentering the compound.

Berzelius’s definition assumed that a catalystmanages to do what it does without changing atall. This was perfectly adequate for describingheterogeneous catalysis, in which the catalyst andthe reactants are in different phases of matter. Inthe platinum-catalyzed reactions that Davy andFaraday observed, for instance, the platinum is asolid, while the reaction itself takes place in agaseous or liquid state. However, homogeneouscatalysis, in which catalyst and reactants are inthe same state, required a different explanation,which English chemist Alexander WilliamWilliamson (1824-1904) provided in an 1852study.

In discussing the reaction observed byKirchhof, of liquid sulfuric acid with starch in anaqueous solution, Williamson was able to showthat the catalyst does break down in the course ofthe reaction. As the reaction takes place, it formsan intermediate compound, but this too is bro-ken down before the reaction ends. The catalystthus emerges in the same form it had at thebeginning of the reaction.

Enzymes: Helpful Catalystsin the Body

In 1833, French physiologist Anselme Payen(1795-1871) isolated a material from malt thataccelerated the conversion of starch to sugar, asfor instance in the brewing of beer. Payen gavethe name “diastase” to this substance, and in1857, the renowned French chemist Louis Pas-teur (1822-1895) suggested that lactic acid fer-mentation is caused by a living organism.

In fact, the catalysts studied by Pasteur arenot themselves separate organisms, as Germanbiochemist Eduard Buchner (1860-1917) showedin 1897. Buchner isolated the catalysts that bringabout the fermentation of alcohol from livingyeast cells—what Payen had called “diastase,” andPasteur “ferments.” Buchner demonstrated thatthese are actually chemical substances, notorganisms. By that time, German physiologistWilly Kahne had suggested the name “enzyme”for these catalysts in living systems.


FRITZ HABER. (Austrian Archives/Corbis. Reproduced by per-

mission.)


A DANGEROUS CATALYST IN

THE ATMOSPHERE. Around the sametime that automakers began rolling out modelsequipped with catalytic converters, scientists andthe general public alike became increasingly con-cerned about another threat to the environment.In the upper atmosphere of Earth are traces ofozone, a triatomic (three-atom) molecular formof oxygen which protects the planet from theSun’s ultraviolet rays. During the latter part ofthe twentieth century, it became apparent that ahole had developed in the ozone layer overAntarctica, and many chemists suspected a cul-prit in chlorofluorocarbons, or CFCs.

CFCs had long been used in refrigerants andair conditioners, and as propellants in aerosolsprays. Because they were nontoxic and noncor-rosive, they worked quite well for such purposes,but the fact that they are chemically unreactivehad an extremely negative side-effect. Instead ofreacting with other substances to form new com-pounds, they linger in Earth’s atmosphere, even-tually drifting to high altitudes, where ultravioletlight decomposes them. The real trouble beginswhen atoms of chlorine, isolated from the CFC,encounter ozone.

Chlorine acts as a catalyst to transform theozone to elemental oxygen, which is not nearly aseffective as ozone for shielding Earth from ultra-violet light. It does so by interacting also withmonatomic, or single-atom oxygen, with which itproduces ClO, or the hypochlorite ion. The endresult of reactions between chlorine, monatomicoxygen, hypochlorite, and ozone is the produc-tion of chlorine, hypochlorite, and diatomic oxy-gen—in other words, no more ozone. It is esti-mated that a single chlorine atom can destroy upto 1 million ozone molecules per second.

Due to concerns about the danger to theozone layer, an international agreement calledthe Montreal Protocol, signed in 1996, bannedthe production of CFCs and the coolant Freonthat contains them. But people still need coolantsfor their homes and cars, and this has led to the creation of substitutes—most notablyhydrochlorofluorocarbons (HCFCs), organiccompounds that do not catalyze ozone.

Other Examples of Catalysts

Catalysts appear in a number of reactions, bothnatural and artificial. For instance, catalysts areused in the industrial production of ammonia,


CatalystsEnzymes are made up of amino acids, whichin turn are constructed from organic compoundscalled proteins. About 20 amino acids make upthe building blocks of the many thousands ofknown enzymes. The beauty of an enzyme is thatit speeds up complex, life-sustaining reactions inthe human body—reactions that would be tooslow at ordinary body temperatures. Rather thanforce the body to undergo harmful increases intemperature, the enzyme facilitates the reactionby opening up a different reaction pathway thatallows a lower activation energy.

One example of an enzyme is cytochrome,which aids the respiratory system by catalyzingthe combination of oxygen with hydrogen with-in the cells. Other enzymes facilitate the conver-sion of food to energy, and make possible a vari-ety of other necessary biological functions.

Because numerous interactions are requiredin their work of catalysis, enzymes are very large,and may have atomic mass figures as high as 1million amu. However, it should be noted thatreactions are catalyzed at very specific loca-tions—called active sites—on an enzyme. Thereactant molecule fits neatly into the active siteon the enzyme, much like a key fitting in a lock;hence the name of this theory, the “lock-and-model.”

Catalysis and the Environment

The exhaust from an automobile contains manysubstances that are harmful to the environment.As a result of increased concerns regarding thepotential damage to the atmosphere, the federalgovernment in the 1970s mandated the adoptionof catalytic converters, devices that employ a cat-alyst to transform pollutants in the exhaust toless harmful substances.

Platinum and palladium are favored materi-als for catalytic converters, though some non-metallic materials, such as ceramics, have beenused as well. In any case, the function of a cat-alytic converter is to convert exhausts throughoxidation-reduction reactions. Nitric oxide isreduced to molecular oxygen and nitrogen; at thesame time, the hydrocarbons in petroleum, alongwith carbon monoxide, are oxidized to form car-bon dioxide and water. Sometimes a reducingagent, such as ammonia, is used to make thereduction process more effective.



ACTIVATION ENERGY: The minimal

energy required to convert reactants into

products, symbolized Ea

AQUEOUS SOLUTIONS: A mixture

of water and a substance that is dis-

solved in it.


a chemical reaction without participat-

ing in it, either as a reactant or product.


reaction.







sions between molecules strong enough to

break bonds in the reactants, resulting in a

reformation of atoms.

HETEROGENEOUS CATALYSIS: A

reaction in which the catalyst and the reac-

tants are in different phases of matter.

HOMOGENEOUS CATALYSIS: A

reaction in which catalyst and reactants are

in the same phase of matter.



reaction.



tion, resulting in a product.








KEY TERMSCatalysts nitric acid (produced from ammonia), sulfuric

acid, and other substances. The ammoniaprocess, developed in 1908 by German chemistFritz Haber (1868-1934), is particularly notewor-thy. Using iron as a catalyst, Haber was able tocombine nitrogen and hydrogen under pressureto form ammonia—one of the world’s mostwidely used chemicals.

Eighteen ninety-seven was a good year forcatalysts. In that year, it was accidentally discov-ered that mercury catalyzes the reaction by whichindigo dye is produced; also in 1897, Frenchchemist Paul Sabatier (1854-1941) found thatnickel catalyzes the production of edible fats.Thanks to Sabatier’s discovery, nickel is used totransform inedible plant oils to margarine andshortening.

Another good year for catalysts—particular-ly those involved in the production of poly-mers—was 1953. That was the year when German chemist Karl Ziegler (1898-1973) dis-covered a resin catalyst for the production ofpolyethylene, which produced a newer, tougherproduct with a much higher melting point thanpolyethylene as it was produced up to that time.Also in 1953, Italian chemist Giulio Natta (1903-1979) adapted Ziegler’s idea, and developed anew type of plastic he called “isotactic” polymers.These could be produced easily, and in abun-dance, through the use of catalysts.

One of the lessons of chemistry, or indeed ofany science, is that there are few things chemistscan do that nature cannot achieve on a far morewondrous scale. No artificial catalyst can com-pete with enzymes, and no use of a catalyst in alaboratory can compare with the grandeur ofthat which takes place on the Sun. As German-American physicist Hans Bethe (1906-) showedin 1938, the reactions of hydrogen that form heli-um on the surface of the Sun are catalyzed bycarbon—the same element, incidentally, foundin all living things on Earth.

WHERE TO LEARN MORE

“Bugs in the News: What the Heck Is an Enzyme?” Univer-sity of Kansas (Web site). <http://falcon.cc.ukans.edu/~jbrown/enzyme.html> (June 9, 2001).

“Catalysis.” University of Idaho Department of Chemistry(Web site). <http://www.chem.uidaho.edu/~honors/rate4.html> (June 9, 2001).


Catalysts“Catalysts” (Web site). <http://edie.cprost.sfu.ca/~rhlogan/catalyst.html> (June 9, 2001).


“Enzymes.” Strategis (Web site). <http://strategis.ic.gc.ca/SSG/tc00048e.html> (June 9, 2001).

“Enzymes: Classification, Structure, Mechanism.” TheHebrew University (Web site). <http://www.md.huji.ac.il/MedChem/Mechanism-Chymotrypsin/> (June9, 2001).


“Ozone Depletion” (Web site). <http://www.energy.rochester.edu/iea/1992/p1/2-3.htm> (June 9, 2001).

“University Chemistry: Chemical Kinetics: Catalysis.” Uni-versity of Alberta Department of Chemistry (Web site).<http://www.chem.ualberta.ca/courses/plambeck/p102/p0217x.htm> (June 9, 2001).





A C IDS AND BASESAcids and Bases

CONCEPTThe name “acid” calls to mind vivid sensoryimages—of tartness, for instance, if the acid inquestion is meant for human consumption, aswith the citric acid in lemons. On the other hand,the thought of laboratory- and industrial-strength substances with scary-sounding names,such as sulfuric acid or hydrofluoric acid, carrieswith it other ideas—of acids that are capable ofdestroying materials, including human flesh. Thename “base,” by contrast, is not widely known inits chemical sense, and even when the older termof “alkali” is used, the sense-impressions pro-duced by the word tend not to be as vivid as thosegenerated by the thought of “acid.” In theirindustrial applications, bases too can be highlypowerful. As with acids, they have many house-hold uses, in substances such as baking soda oroven cleaners. From a taste standpoint, (as any-one who has ever brushed his or her teeth withbaking soda knows), bases are bitter rather thansour. How do we know when something is anacid or a base? Acid-base indicators, such as lit-mus paper and other materials for testing pH,offer a means of judging these qualities in vari-ous substances. However, there are larger struc-tural definitions of the two concepts, whichevolved in three stages during the late nineteenthand early twentieth centuries, that provide amore solid theoretical underpinning to theunderstanding of acids and bases.

HOW IT WORKS

Introduction to Acids and Bases

Prior to the development of atomic and molecu-lar theory in the nineteenth century, followed by

the discovery of subatomic structures in the latenineteenth and early twentieth centuries,chemists could not do much more than makemeasurements and observations. Their defini-tions of substances were purely phenomenologi-cal—that is, the result of experimentation andthe collection of data. From these observations,they could form general rules, but they lackedany means of “seeing” into the atomic andmolecular structures of the chemical world.

The phenomenological distinctions betweenacids and bases, gathered by scientists fromancient times onward, worked well enough formany centuries. The word “acid” comes from theLatin term acidus, or “sour,” and from an earlyperiod, scientists understood that substancessuch as vinegar and lemon juice shared a com-mon acidic quality. Eventually, the phenomeno-logical definition of acids became relativelysophisticated, encompassing such details as thefact that acids produce characteristic colors incertain vegetable dyes, such as those used in mak-ing litmus paper. In addition, chemists realizedthat acids dissolve some metals, releasing hydro-gen in the process.

WHY “BASE” AND NOT “ALKA-

LI”? The word “alkali” comes from the Arabical-qili, which refers to the ashes of the seawortplant. The latter, which typically grows in marshyareas, was often burned to produce soda ash,used in making soap. In contrast to acids, bases—caffeine, for example—have a bitter taste, andmany of them feel slippery to the touch. Theyalso produce characteristic colors in the vegetabledyes of litmus paper, and can be used to promotecertain chemical reactions. Note that todaychemists use the word “base” instead of “alkali,”


Acids andBases


the reason being that the latter term has a nar-rower meaning: all alkalies are bases, but not allbases are alkalies.

Originally, “alkali” referred only to the ashesof burned plants, such as seawort, that containedeither sodium or potassium, and from which theoxides of sodium and potassium could beobtained. Eventually, alkali came to mean the sol-uble hydroxides of the alkali and alkaline earthmetals. This includes sodium hydroxide, theactive ingredient in drain and oven cleaners;magnesium hydroxide, used for instance in milkof magnesia; potassium hydroxide, found insoaps and other substances; and other com-pounds. Broad as this range of substances is, itfails to encompass the wide array of materialsknown today as bases—compounds which reactwith acids to form salts and water.

Toward a Structural Definition

The reaction to form salts and water is, in fact,one of the ways that acids and bases can bedefined. In an aqueous solution, hydrochloricacid and sodium hydroxide react to form sodiumchloride—which, though it is suspended in anaqueous solution, is still common table salt—along with water. The equation for this reactionis HCl(aq) + NaOH(aq) 4H2O + NaCl(aq). Inother words, the sodium (Na) ion in sodiumhydroxide switches places with the hydrogen ionin hydrochloric acid, resulting in the creation ofNaCl (salt) along with water.

But why does this happen? Useful as this def-inition regarding the formation of salts andwater is, it is still not structural—in other words,it does not delve into the molecular structure andbehavior of acids and bases. Credit for the firsttruly structural definition of the difference goesto the Swedish chemist Svante Arrhenius (1859-1927). It was Arrhenius who, in his doctoral dis-sertation in 1884, introduced the concept of anion, an atom possessing an electric charge.

His understanding was particularly impres-sive in light of the fact that it was 13 more yearsbefore the discovery of the electron, the sub-atomic particle responsible for the creation ofions. Atoms have a neutral charge, but when anelectron or electrons depart, the atom becomes apositive ion or cation. Similarly, when an elec-tron or electrons join a previously uncharged

atom, the result is a negative ion or anion. Notonly did the concept of ions greatly influence thefuture of chemistry, but it also provided Arrhe-nius with the key necessary to formulate his dis-tinction between acids and bases.

The Arrhenius Definition

Arrhenius observed that molecules of certaincompounds break into charged particles whenplaced in liquid. This led him to the Arrheniusacid-base theory, which defines an acid as anycompound that produces hydrogen ions (H+)when dissolved in water, and a base as any com-pound that produces hydroxide ions (OH-) whendissolved in water.

This was a good start, but two aspects ofArrhenius’s theory suggested the need for a defi-nition that encompassed more substances. Firstof all, his theory was limited to reactions in aque-ous solutions. Though many acid-base reactionsdo occur when water is the solvent, this is notalways the case.

Second, the Arrhenius definition effectivelylimited acids and bases only to those ionic com-pounds, such as hydrochloric acid or sodiumhydroxide, which produced either hydrogen or

SVANTE ARRHENIUS. (Hulton-Deutsch Collection/Corbis. Repro-



Acids andBases

hydroxide ions. However, ammonia, or NH3, actslike a base in aqueous solutions, even though itdoes not produce the hydroxide ion. The same istrue of other substances, which behave like acidsor bases without conforming to the Arrheniusdefinition.

These shortcomings pointed to the need fora more comprehensive theory, which arrivedwith the formulation of the Brønsted-Lowry definition by English chemist Thomas Lowry(1874-1936) and Danish chemist J. N. Brønsted(1879-1947). Nonetheless, Arrhenius’s theoryrepresented an important first step, and in 1903,he was awarded the Nobel Prize in Chemistryfor his work on the dissociation of moleculesinto ions.

The Brønsted-Lowry Definition

The Brønsted-Lowry acid-base theory defines an

acid as a proton (H+) donor, and a base as a pro-

ton acceptor, in a chemical reaction. Protons are

represented by the symbol H+, and in represent-

ing acids and bases, the symbols HA and A-,

respectively, are used. These symbols indicate

that an acid has a proton it is ready to give away,

while a base, with its negative charge, is ready to

receive the positively charged proton.

Though it is used here to represent a proton,

it should be pointed out that H+ is also the

hydrogen ion—a hydrogen atom that has lost its

sole electron and thus acquired a positive charge.


THE TARTNESS OF LEMONS AND LIMES IS A FUNCTION OF THEIR ACIDITY—SPECIFICALLY THE WAY IN WHICH CITRIC

ACID MOLECULES FIT INTO THE TONGUE’S “SWEET” RECEPTORS. (Orion Press/Corbis. Reproduced by permission.)


Acids andBases

It is thus really nothing more than a lone proton,but this is the one and only case in which anatom and a proton are exactly the same thing. Inan acid-base reaction, a molecule of acid is“donating” a proton, in the form of a hydrogenion. This should not be confused with a far morecomplex process, nuclear fusion, in which anatom gives up a proton to another atom.

AN ACID-BASE REACTION IN

BRØNSTED-LOWRY THEORY. Themost fundamental type of acid-base reaction inBrønsted-Lowry theory can be symbolized thusHA(aq) + H2O(l) 4H3O+(aq) + A-(aq). The firstacid shown—which, like three of the four “play-ers” in this equation, is dissolved in an aqueoussolution—combines with water, which can serveas either an acid or a base. In the present context,it functions as a base.

Water molecules are polar, meaning that thenegative charges tend to congregate on one endof the molecule with the oxygen atom, while thepositive charges remain on the other end with thehydrogen atoms. The Brønsted-Lowry modelemphasizes the role played by water, which pullsthe proton from the acid, resulting in the cre-ation of H3O+, known as the hydronium ion.

The hydronium ion produced here is anexample of a conjugate acid, an acid formedwhen a base accepts a proton. At the same time,the acid has lost its proton, becoming A–, a con-jugate base—that is, the base formed when anacid releases a proton. These two products of thereaction are called a conjugate acid-base pair, aterm that refers to two substances related to oneanother by the donating of a proton.

Brønsted and Lowry’s definition representsan improvement over that of Arrhenius, becauseit includes all Arrhenius acids and bases, as wellas other chemical species not encompassed inArrhenius theory. An example, mentioned earlier, is ammonia. Though it does not produceOH– ions, ammonia does accept a proton from awater molecule, and the reaction between thesetwo (with water this time serving the function ofacid) produces the conjugate acid-base pair ofNH4

+ (an ammonium ion) and OH-. Note thatthe latter, the hydroxide ion, was not produced byammonia, but is the conjugate base that resultedwhen the water molecule lost its H+ atom or proton.

The Lewis Definition

Despite the progress offered to chemists by theBrønsted-Lowry model, it was still limited todescribing compounds that contain hydrogen. AsAmerican chemist Gilbert N. Lewis (1875-1946)recognized, this did not encompass the full rangeof acids and bases; what was needed, instead, wasa definition that did not involve the presence of ahydrogen atom.

Lewis is particularly noted for his work inthe realm of chemical bonding. The bonding ofatoms is the result of activity on the part of thevalence electrons, or the electrons at the “out-side” of the atom. Electrons are arranged in dif-ferent ways, depending on the type of bonding,but they always bond in pairs.

According to the Lewis acid-base theory, anacid is the reactant that accepts an electron pairfrom another reactant in a chemical reaction,while a base is the reactant that donates an electron pair to another reactant. As with theBrønsted-Lowry definition, the Lewis definitionis reaction-dependant, and does not define acompound as an acid or base in its own right.Instead, the manner in which the compoundreacts with another serves to identify it as an acidor base.

AN IMPROVEMENT OVER ITS

PREDECESSORS. The beauty of theLewis definition lies in the fact that it encom-passes all the situations covered by the others—and more. Just as Brønsted-Lowry did not dis-prove Arrhenius, but rather offered a definitionthat covered more substances, Lewis expandedthe range of substances beyond those covered byBrønsted-Lowry. In particular, Lewis theory canbe used to differentiate the acid and base inbond-producing chemical reactions where ionsare not produced, and in which there is no pro-ton donor or acceptor. Thus it represents animprovement over Arrhenius and Brønsted-Lowry respectively.

An example is the reaction of boron trifluo-ride (BF3) with ammonia (NH3), both in the gasphases, to produce boron trifluoride ammoniacomplex (F3BNH3). In this reaction, boron triflu-oride accepts an electron pair and is therefore aLewis acid, while ammonia donates the electronpair and is thus a Lewis base. Though hydrogen isinvolved in this particular reaction, Lewis theoryalso addresses reactions involving no hydrogen.



Acids andBases


pH and Acid-Base Indicators

Though chemists apply the sophisticated struc-tural definitions for acids and bases that we havediscussed, there are also more “hands-on” meth-ods for identifying a particular substance(including complex mixtures) as an acid or base.Many of these make use of the pH scale, devel-oped by Danish chemist Søren Sørensen (1868-1939) in 1909.

The term pH stands for “potential of hydro-gen,” and the pH scale is a means of determiningthe acidity or alkalinity of a substance. (Though,as noted, the term “alkali” has been replaced by“base,” alkalinity is still used as an adjectival termto indicate the degree to which a substance dis-plays the properties of a base.) There are theoret-ically no limits to the range of the pH scale, butfigures for acidity and alkalinity are usually givenwith numerical values between 0 and 14.

THE MEANING OF pH VALUES.

A rating of 0 on the pH scale indicates a sub-stance that is virtually pure acid, while a 14 rat-ing represents a nearly pure base. A rating of 7indicates a neutral substance. The pH scale islogarithmic, or exponential, meaning that thenumbers represent exponents, and thus anincreased value of 1 represents not a simplearithmetic addition of 1, but an increase of 1power. This, however, needs a little further expla-nation.

The pH scale is actually based on negativelogarithms for the values of H3O+ (the hy-dronium ion) or H+ (protons) in a given sub-stance. The formula is thus pH = -log[H3O+] or -log[H+], and the presence of hydronium ions orprotons is measured according to their concen-tration of moles per liter of solution.

pH VALUES OF VARIOUS SUB-

STANCES. The pH of a virtually pure acid,such as the sulfuric acid in car batteries, is 0, andthis represents 1 mole (mol) of hydronium perliter (l) of solution. Lemon juice has a pH of 2,equal to 10-2 mol/l. Note that the pH value of 2translates to an exponent of -2, which, in thiscase, results in a figure of 0.01 mol/l.

Distilled water, a neutral substance with apH of 7, has a hydronium equivalent of 10-7

mol/l. It is interesting to observe that most of the

fluids in the human body have pH values in theneutral range blood (venous, 7.35; arterial, 7.45);urine (6.0—note the higher presence of acid);and saliva (6.0 to 7.4).

At the alkaline end of the scale is borax, witha pH of 9, while household ammonia has a pHvalue of 11, or 10-11 mol/l. Sodium hydroxide, orlye, an extremely alkaline chemical with a pH of14, has a value equal to 10-14 moles of hydroniumper liter of solution.

LITMUS PAPER AND OTHER

INDICATORS. The most precise pH meas-urements are made with electronic pH meters,which can provide figures accurate to 0.001 pH.However, simpler materials are also used. Bestknown among these is litmus paper (made froman extract of two lichen species), which turnsblue in the presence of bases and red in the pres-ence of acids. The term “litmus test” has becomepart of everyday language, referring to a make-or-break issue—for example, “views on abortionrights became a litmus test for Supreme Courtnominees.”

Litmus is just one of many materials used formaking pH paper, but in each case, the change ofcolor is the result of the neutralization of thesubstance on the paper. For instance, paper coat-ed with phenolphthalein changes from colorlessto pink in a pH range from 8.2 to 10, so it is use-ful for testing materials believed to be moderate-ly alkaline. Extracts from various fruits and veg-etables, including red cabbages, red onions, andothers, are also applied as indicators.

Some Common Acids and Bases

The tables below list a few well-known acids and bases, along with their formulas and a fewapplications

Common Acids

• Acetic acid (CH3COOH): vinegar, acetate

• Acetylsalicylic acid (HOOCC6H4OOCCH3):aspirin

• Ascorbic acid (H2C6H6O6): vitamin C

• Carbonic acid (H2CO3): soft drinks, seltzerwater

• Citric acid (C6H8O7): citrus fruits, artificialflavorings

• Hydrochloric acid (HCl): stomach acid

• Nitric acid (HNO3): fertilizer, explosives

• Sulfuric acid (H2SO4): car batteries



Common Bases

• Aluminum hydroxide (Al[OH]3): antacids,deodorants

• Ammonium hydroxide (NH4OH): glasscleaner

• Calcium hydroxide (Ca[OH]2): causticlime, mortar, plaster

• Magnesium hydroxide (Mg[OH]2): laxa-tives, antacids

• Sodium bicarbonate/sodium hydrogen car-bonate (NaHCO3): baking soda

• Sodium carbonate (Na2CO3): dish detergent

• Sodium hydroxide (NaOH): lye, oven anddrain cleaner

• Sodium hypochlorite (NaClO): bleach

Of course these represent only a few of themany acids and bases that exist. Selected sub-stances listed above are discussed briefly below.

Acids

ACIDS IN THE HUMAN BODY

AND FOODS. As its name suggests, citricacid is found in citrus fruits—particularlylemons, limes, and grapefruits. It is also used as aflavoring agent, preservative, and cleaning agent.Produced commercially from the fermentationof sugar by several species of mold, citric acidcreates a taste that is both tart and sweet. Thetartness, of course, is a function of its acidity, ora manifestation of the fact that it produceshydrogen ions. The sweetness is a more complexbiochemical issue relating to the ways that citricacid molecules fit into the tongue’s “sweet”receptors.

Citric acid plays a role in one famous stom-ach remedy, or antacid. This in itself is interest-ing, since antacids are more generally associatedwith alkaline substances, used for their ability toneutralize stomach acid. The fizz in Alka-Seltzer,however, comes from the reaction of citric acids(which also provide a more pleasant taste) withsodium bicarbonate or baking soda, a base. Thisreaction produces carbon dioxide gas. As a pre-servative, citric acid prevents metal ions fromreacting with, and thus hastening the degrada-tion of, fats in foods. It is also used in the pro-duction of hair rinses and low-pH shampoos andtoothpastes.

The carboxylic acid family of hydrocarbonderivatives includes a wide array of substances—not only citric acids, but amino acids. Amino

acids combine to make up proteins, one of theprincipal components in human muscles, skin,and hair. Carboxylic acids are also applied indus-trially, particularly in the use of fatty acids formaking soaps, detergents, and shampoos.

SULFURIC ACID. There are plenty ofacids found in the human body, includinghydrochloric acid or stomach acid—which, inlarge quantities, causes indigestion, and the needfor neutralization with a base. Nature also pro-duces acids that are toxic to humans, such as sul-furic acid.

Though direct exposure to sulfuric acid isextremely dangerous, the substance has numer-ous applications. Not only is it used in car batter-ies, but sulfuric acid is also a significant compo-nent in the production of fertilizers. On the otherhand, sulfuric acid is damaging to the environ-ment when it appears in the form of acid rain.Among the impurities in coal is sulfur, and thisresults in the production of sulfur dioxide andsulfur trioxide when the coal is burned. Sulfurtrioxide reacts with water in the air, creating sul-furic acid and thus acid rain, which can endangerplant and animal life, as well as corrode metalsand building materials.

Bases

The alkali metal and alkaline earth metal familiesof elements are, as their name suggests, bases. Anumber of substances created by the reaction ofthese metals with nonmetallic elements are takeninternally for the purpose of settling gastric trou-ble or clearing intestinal blockage. For instance,there is magnesium sulfate, better known asEpsom salts, which provide a powerful laxativealso used for ridding the body of poisons.

Aluminum hydroxide is an interesting base,because it has a wide number of applications,including its use in antacids. As such, it reactswith and neutralizes stomach acid, and for thatreason is found in commercial antacids such asDi-Gel™, Gelusil™, and Maalox™. Aluminumhydroxide is also used in water purification, indyeing garments, and in the production of cer-tain kinds of glass. A close relative, aluminumhydroxychloride or Al2(OH)5Cl, appears in manycommercial antiperspirants, and helps to closepores, thus stopping the flow of perspiration.

SODIUM HYDROGEN CAR-

BONATE (BAKING SODA). Bakingsoda, known by chemists both as sodium bicar-


Acids andBases


Acids andBases


bonate and sodium hydrogen carbonate, isanother example of a base with multiple purpos-es. As noted earlier, it is used in Alka-Seltzer™,with the addition of citric acid to improve theflavor; in fact, baking soda alone can perform the

function of an antacid, but the taste is rather un-pleasant.

Baking soda is also used in fighting fires,because at high temperatures it turns into carbondioxide, which smothers flames by obstructing

ACID: A substance that, in its edible

form, is sour to the taste, and in non-edible

forms, is often capable of dissolving met-

als. Acids and bases react to form salts and

water. These are all phenomenological def-

initions, however, in contrast to the three

structural definitions of acids and bases—

the Arrhenius, BrØnsted-Lowry, and Lewis

acid-base theories.

ALKALI: A term referring to the soluble

hydroxides of the alkali and alkaline earth

metals. Once “alkali” was used for the class

of substances that react with acids to form

salts; today, however, the more general

term base is preferred.

ALKALINITY: An adjectival term used

to identify the degree to which a substance

displays the properties of a base.

ANION: The negatively charged ion

that results when an atom gains one or

more electrons. “Anion” is pronounced

“AN-ie-un”.

AQUEOUS SOLUTION: A substance

in which water constitutes the solvent. A

large number of chemical reactions take

place in an aqueous solution.

ARRHENIUS ACID-BASE THEORY:

The first of three structural definitions of

acids and bases. Formulated by Swedish

chemist Svante Arrhenius (1859-1927), the

Arrhenius theory defines acids and bases

according to the ions they produce in an

aqueous solution: an acid produces hydro-

gen ions (H+), and a base hydroxide ions

(OH–).

BASE: A substance that, in its edible

form, is bitter to the taste. Bases tend to be

slippery to the touch, and in reaction with

acids they produce salts and water. Bases

and acids are most properly defined, how-

ever, not in these phenomenological terms,

but by the three structural definitions of

acids and bases—the Arrhenius, BrØnsted-

Lowry, and Lewis acid-base theories.

BRØNSTED-LOWRY ACID-BASE THE-

ORY: The second of three structural

definitions of acids and bases. Formulated

by English chemist Thomas Lowry (1874-

1936) and Danish chemist J. N. BrØnsted

(1879-1947), BrØnsted-Lowry theory

defines an acid as a proton (H+) donor, and

a base as a proton acceptor.

CATION: The positively charged ion

that results when an atom loses one or

more electrons. “Cation” is pronounced

“KAT-ie-un”.





so forth.

CONJUGATE ACID: An acid formed

when a base accepts a proton (H+).

KEY TERMS


Acids andBases

the flow of oxygen to the fire. Of course, bakingsoda is also used in baking, when it is combinedwith a weak acid to make baking powder. Thereaction of the acid and the baking soda pro-duces carbon dioxide, which causes dough and

batters to rise. In a refrigerator or cabinet, bakingsoda can absorb unpleasant odors, and addition-ally, it can be applied as a cleaning product.

SODIUM HYDROXIDE (LYE) .

Another base used for cleaning is sodium


CONJUGATE ACID-BASE PAIR:

The acid and base produced when an acid

donates a single proton to a base. In the

reaction that produces this pair, the acid

and base switch identities. By donating a

proton, the acid becomes a conjugate base,

and by receiving the proton, the base

becomes a conjugate acid.

CONJUGATE BASE: A base formed

when an acid releases a proton.



a net electric charge. There are two types of

ions: anions and cations.




charges.



contain at least one metal and nonmetal

joined by an ionic bond.

LEWIS ACID-BASE THEORY: The

third of three structural definitions of

acids and bases. Formulated by American

chemist Gilbert N. Lewis (1875-1946),

Lewis theory defines an acid as the reactant

that accepts an electron pair from another

reactant in a chemical reaction, and a base

as the reactant that donates an electron

pair to another reactant.

PH SCALE: A logarithmic scale for

determining the acidity or alkalinity of a

substance, from 0 (virtually pure acid) to 7

(neutral) to 14 (virtually pure base).

PHENOMENOLOGICAL: A term

describing scientific definitions based

purely on experimental phenomena. These

convey only part of the picture, however—

primarily, the part a chemist can perceive

either through measurement or through

the senses, such as sight. A structural defi-

nition is therefore usually preferable to a

phenomenological one.



tion, resulting in the creation of a product.

SALTS: Ionic compounds formed by

the reaction between an acid and a base. In

this reaction, one or more of the hydrogen

ions of an acid is replaced with another

positive ion. In addition to producing salts,

acid-base reactions produce water.

SOLUTION: A homogeneous mixture

in which one or more substances (the

solute) is dissolved in one or more other

substances (the solvent)—for example,

sugar dissolved in water.

SOLVENT: A substance that dissolves

another, called a solute, in a solution.

STRUCTURAL: A term describing sci-

entific definitions based on aspects of

molecular structure and behavior rather

than purely phenomenological data.



Acids andBases

hydroxide, known commonly as lye or causticsoda. Unlike baking soda, however, it is not to betaken internally, because it is highly damaging tohuman tissue—particularly the eyes. Lye appearsin drain cleaners, such as Drano™, and ovencleaners, such as Easy-Off™, which make use ofits ability to convert fats to water-soluble soap.

In the process of doing so, however, relative-ly large amounts of lye may generate enough heatto boil the water in a drain, causing the water toshoot upward. For this reason, it is not advisableto stand near a drain being treated with lye. In aclosed oven, this is not a danger, of course; andafter the cleaning process is complete, the con-verted fats (now in the form of soap) can be dis-solved and wiped off with a sponge.

WHERE TO LEARN MORE

“Acids and Bases Frequently Asked Questions.” GeneralChemistry Online (Web site). <http://antoine.fsu.umd.edu/chem/senese/101/acidbase/faq.shtml>(June 7, 2001).

“Acids, Bases, and Salts.” Chemistry Coach (Web site).<http://www.chemistrycoach.com/acids.htm> (June7, 2001).

“Acids, Bases, and Salts.” University of Akron, Departmentof Chemistry (Web site). <http://ull.chemistry.uakron.edu/genobc/Chapter_09/title.html> (June 7,2001).

ChemLab. Danbury, CT: Grolier Educational, 1998.


Haines, Gail Kay. What Makes a Lemon Sour? Illustratedby Janet McCaffery. New York: Morrow, 1977.

Oxlade, Chris. Acids and Bases. Chicago: HeinemannLibrary, 2001.

Patten, J. M. Acids and Bases. Vero Beach, FL: RourkeBook Company, 1995.

Walters, Derek. Chemistry. Illustrated by Denis Bishopand Jim Robins. New York: F. Watts, 1982.

Zumdahl, Steven S. Introductory Chemistry A Founda-tion, 4th ed. Boston: Houghton Mifflin, 2000.




A C ID -BASE REACT IONSAcid-Base Reactions

CONCEPTTo an extent, acids and bases can be defined interms of factors that are apparent to the senses:edible acids taste sour, for instance, while basesare bitter-tasting and slippery to the touch. Thebest way to understand these two types of sub-stances, however, is in terms of their behavior inchemical reactions. Not only do the reactions ofacids and bases result in the creation of salts andwater, but acids and bases can be defined by theways in which they participate in a reaction—forinstance, by donating or accepting electron pairs.The reaction of acids and bases to form waterand salts is called neutralization, and it has a widerange of applications, including the promotionof plant growth in soil and the treatment ofheartburn in the human stomach. Neutralizationalso makes it possible to test substances for theirpH level, a measure of the degree to which thesubstance is acidic or alkaline.

HOW IT WORKS

Phenomenological Defini-tions of Acids and Bases

Before studying the reactions of acids and bases,it is necessary to define exactly what each is. Thisis not as easy as it sounds, and the Acids andBases essay discusses in detail a subject coveredmore briefly here: the arduous task chemistsfaced in developing a workable distinction. Let usstart with the phenomenological differencesbetween the two—that is, aspects relating tothings that can readily be observed without refer-ring to the molecule properties and behaviors ofacids and bases.

Acids are fairly easy to understand on thephenomenological level: the name comes fromthe Latin term acidus, or “sour,” and many soursubstances from daily life—lemons, for instance,or vinegar—are indeed highly acidic. In fact,lemons and most citrus fruits contain citric acid(C6H8O7), while the acidic quality of vinegarcomes from acetic acid (CH3COOH). In addi-tion, acids produce characteristic colors in cer-tain vegetable dyes, such as those used in makinglitmus paper.

The word “base,” as it is used in this context,may be a bit more difficult to appreciate on a sen-sory level. It helps, perhaps, if the older term“alkali” is used, though even so, people tend tothink of alkaline substances primarily in contrastto acids. “Alkali,” which serves to indicate thebasic quality of both the alkali metal and alkalineearth metal families of elements, comes from theArabic al-qili. The latter refers to the ashes of theseawort plant, which usually grows in marshyareas and, in the past, was often burned to pro-duce soda ash for making soap.

The reason chemists of today use the word“base” instead of “alkali” is that the latter termhas a narrower meaning: all alkalies are bases, butnot all bases are alkalies. Originally referring onlyto the ashes of burned plants containing eithersodium or potassium, alkali was eventually usedto designate the soluble hydroxides of the alkaliand alkaline earth metals. Among these are sodi-um hydroxide or lye; magnesium hydroxide(found in milk of magnesia); potassium hydrox-ide, which appears in soaps; and other com-pounds. Because these represent only a few ofthe substances that react with acids in the waysdiscussed in this essay, the term “base” is preferred.


Acid-BaseReactions

THE FORMATION OF SALTS. Aschemistry evolved, and physical scientistsbecame aware of the atomic and molecular sub-structures that make up the material world, theydeveloped more fundamental distinctionsbetween acids and bases. By the early twentiethcentury, chemists had applied structural distinc-tions between acids and bases—that is, defini-tions based on the molecular structures andbehaviors of those substances.

An important intermediary step occurred aschemists came to the conclusion that reactions ofacids and bases form salts and water. Forinstance, in an aqueous solution, hydrochloricacid or HCl(aq) reacts with the base sodiumhydroxide, designated as NaOH(aq), to formsodium chloride, or common table salt(NaCl[aq]) and H2O. What happens is that thesodium (Na) ion (an atom with an electriccharge) in sodium hydroxide switches placeswith the hydrogen ion in hydrochloric acid,resulting in the creation of NaCl and water.

Ions themselves had yet to be defined in1803, when the great Swedish chemist JonsBerzelius (1779-1848) added another piece to thefoundation for a structural definition. Acids andbases, he suggested, have opposite electriccharges. In this, he was about eight decades ahead

of his time: only in 1884 did his countrymanSvante Arrhenius (1859-1927) introduce theconcept of the ion. This, in turn, enabled Arrhe-nius to formulate the first structural distinctionbetween acids and bases.

The Arrhenius Acid-Base Theory

Arrhenius acid-base theory defines the two sub-stances with regard to their behavior in an aque-ous solution: an acid is any compound that pro-duces hydrogen ions (H+), and a base is one thatproduces hydroxide ions (OH–) when dissolvedin water. This occurred, for instance, in the reac-tion discussed above: the hydrochloric acid pro-duced a hydrogen ion, while the sodium hydrox-ide produced a hydroxide ion, and these two ionsbonded to form water.

Though it was a good start, Arrhenius’s the-ory was limited to reactions in aqueous solu-tions. In addition, it confined its definition ofacids and bases only to those ionic compounds,such as hydrochloric acid or sodium hydroxide,that produced either hydrogen or hydroxide ions.But ammonia, or NH3, acts like a base in aqueoussolutions, even though it does not produce thehydroxide ion. These shortcomings pointed tothe need for a more comprehensive theory, whichcame with the formulation of the Brønsted-Lowry definition.

The Brønsted-Lowry Acid-Base Theory

Developed by English chemist Thomas Lowry(1874-1936) and Danish chemist J. N. Brønsted(1879-1947), the Brønsted-Lowry acid-base the-ory defines an acid as a proton (H+) donor, and abase as a proton acceptor, in a chemical reaction.Protons are represented by the symbol H+, acation (positively charged ion) of hydrogen.

Elemental hydrogen, called protium to dis-tinguish it from its isotopes, has just one protonand one electron—no neutrons. Therefore, thehydrogen cation, which has to lose its sole elec-tron to gain a positive charge, is essentially noth-ing but a proton. It is thus at once an atom, anion, and a proton, but the ionization of hydrogenconstitutes the only case in which this is possible.

Thus when the term “proton donor” or “pro-ton acceptor” is used, it does not mean that aproton is splitting off from an atom or joining


GILBERT N. LEWIS.


Acid-BaseReactions

another, as in a nuclear reaction. Rather, when anacid behaves as a proton donor, this means thatthe hydrogen proton/ion/atom is separatingfrom an acidic compound; conversely, when abase acts as a proton acceptor, the positivelycharged hydrogen ion is bonding with the basiccompound.

REACTIONS IN BRØNSTED-

LOWRY ACID-BASE THEORY. Inrepresenting Brønsted-Lowry acids and bases,the symbols HA and A–, respectively, are used.These appear in the equation representing themost fundamental type of Brønsted-Lowry acid-base reaction: HA(aq) + H2O(l) 4H3O+(aq) +A–(aq). The symbols (aq), (l), and 4are explainedin the Chemical Reactions essay. In plain English,this equation states that when an acid in an aque-ous solution reacts with liquid water, the result isthe creation of H3O+, known as the hydroniumion, along with a base. Both products of the reac-tion are dissolved in an aqueous solution.

Because water molecules are polar, the nega-tive charges tend to congregate on one end of themolecule with the oxygen atom, while the posi-tive charges remain on the other end with thehydrogen atoms. The Brønsted-Lowry modelemphasizes the role played by water, which pullsthe proton from the acid, resulting in the cre-ation of the hydronium ion.

The hydronium ion, in this equation, is anexample of a conjugate acid, an acid formedwhen a base accepts a proton. At the same time,the acid has lost its proton, becoming A–, a con-jugate base—that is, the base formed when anacid releases a proton. These two products of thereaction are called a conjugate acid-base pair, aterm that refers to two substances related to oneanother by the donating of a proton.

Brønsted and Lowry’s definition includes allArrhenius acids and bases, as well as other chem-ical species not encompassed in Arrhenius theo-ry. As mentioned earlier, ammonia is a base, yet itdoes not produce OH– ions; however, it doesaccept a proton from a water molecule. Water canserve either as an acid or base; in this instance, itis an acid, and in reaction with ammonia, it pro-duces the conjugate acid-base pair of NH4

+ (anammonium ion) and OH–. Ammonia did notproduce the hydroxide ion here; rather, OH– isthe conjugate base that resulted when the watermolecule lost its H+ atom (i.e., a proton.)

The Lewis Acid-Base Theory

The Brønsted-Lowry model still had its limita-tions, in that it only described compounds con-taining hydrogen. American chemist Gilbert N.Lewis (1875-1946), however, developed a theoryof acids and bases that makes no reference to thepresence of hydrogen. Instead, it relates to some-thing much more fundamental: the fact thatchemical bonding always involves pairs of elec-trons.

Lewis acid-base theory defines an acid as thereactant that accepts an electron pair fromanother reactant in a chemical reaction, while a base is the reactant that donates an electronpair to another reactant. Note that, as with theBrønsted-Lowry definition, the Lewis definitionis reaction-dependant. Instead of defining acompound as an acid or base in its own right, itidentifies these in terms of how the compoundreacts with another.

The Lewis definition encompasses all the sit-uations covered by the others, as well as manyother reactions not described in the theories ofeither Arrhenius or Brønsted-Lowry. In particu-lar, Lewis theory can be used to differentiate theacid and base in chemical reactions where ionsare not produced, something that takes it farbeyond the scope of Arrhenius theory. Also,Lewis theory addresses situations in which thereis no proton donor or acceptor, thus offering animprovement over Brønsted-Lowry.

When boron trifluoride (BF3) and ammonia(NH3), both in the gas phases, react to produceboron trifluoride ammonia complex (F3BNH3),boron trifluoride accepts an electron pair. There-fore, it is a Lewis acid, while ammonia—whichdonates the electron pair—can be defined as aLewis base. This particular reaction involveshydrogen, but since the operative factor in Lewistheory relates to electron pairs and not hydrogen,the theory can be used to address reactions inwhich that element is not present.


Dissociation

Dissociation is the separation of a molecule intoions, and it is a key factor for evaluating the“strength” of acids and bases. The more a sub-stance is prone to dissociation, the better it can



Acid-BaseReactions

conduct an electric current, because the separa-tion of charges provides a “pathway” for the cur-rent’s flow. A substance that dissociates com-pletely, or almost completely, is called a strongelectrolyte, whereas one that dissociates onlyslightly (or not at all) is designated as a weakelectrolyte.

The terms “weak” and “strong” are alsoapplied to acids and bases. For instance, vinegaris a weak acid, because it dissociates only slightly,and therefore conducts little electric current. Bycontrast, hydrochloric acid (HCl) is a strong acid,because it dissociates almost completely intopositively charged hydrogen ions and negativelycharged chlorine ones. Represented symbolically,this is: HCl 4H+ + Cl–.

A REACTION INVOLVING A

STRONG ACID. It may seem a bit back-ward that a strong acid or base is one that “fallsapart,” while the weak one stays together. Tounderstand the difference better, let us return tothe reaction described earlier, in which an acid inaqueous solution reacts with water to produce abase in aqueous solution, along with hydronium:HA(aq) + H2O(l) 4H3O+(aq) + A–(aq). Instead ofusing the generic symbols HA and A–, however, letus substitute hydrochloric acid (HCl) and chlo-ride (Cl–) respectively.

The reaction HCl(aq) + H2O(l) 4H3O+(aq)+ Cl–(aq) is a reversible one, and for that reason,the symbol for chemical equilibrium (!) can beinserted in place of the arrow pointing to theright. In other words, the substances on the rightcan just as easily react, producing the substanceson the left. In this reverse reaction, the reactantsof the forward reaction would become products,and the products of the forward reaction serve asthe reactants.

However, the reaction described here is notperfectly reversible, and in fact the most properchemical symbolism would show a longer arrowpointing to the right, with a shorter arrow point-ing to the left. Due to the presence of a strongelectrolyte, there is more forward “thrust” to thisreaction.

Because it is a strong acid, the hydrogenchloride in solution is not a set of molecules, buta collection of H+ and Cl– ions. In the reaction,the weak Cl– ions to the right side of the equilib-rium symbol exert very little attraction for the H+

ions. Instead of bonding with the chloride, these

hydrogen ions join the water (a stronger base) toform hydronium.

The chloride, incidentally, is the conjugatebase of the hydrochloric acid, and this illustratesanother principal regarding the “strength” ofelectrolytes: a strong acid produces a relativelyweak conjugate base. Likewise, a strong base pro-duces a relatively weak conjugate acid.

THE STRONG ACIDS AND

BASES. There are only a few strong acids andbases, which are listed below:

Strong Acids

• Hydrobromic acid (HBr)

• Hydrochloric acid (HCl)

• Hydroiodic acid (HI)

• Nitric acid (HNO3)

• Perchloric acid (HClO4)

• Sulfuric acid (H2SO4)

Strong Bases

• Barium hydroxide (Ba[OH]2)

• Calcium hydroxide (Ca[OH]2)

• Lithium hydroxide (LiOH)

• Potassium hydroxide (KOH)

• Sodium hydroxide (NaOH)

• Strontium hydroxide (Sr[OH]2)

Virtually all others are weak acids or bases,meaning that only a small percentage of mole-cules in these substances ionize by dissociation.The concentrations of the chemical speciesinvolved in the dissociation of weak acids andbases are mathematically governed by the equi-librium constant K..

Neutralization

Neutralization is the process whereby an acid andbase react with one another to form a salt andwater. The simplest example of this occurs in thereaction discussed earlier, in which hydrochloricacid or HCl(aq) reacts with the base sodiumhydroxide, designated as NaOH(aq), in an aqueous solution. The result is sodium chlo-ride, or common table salt (NaCl[aq]) and H2O.This equation is written thus: HCl(aq) +NaOH(aq) 4NaCl(aq) + H2O.

The human stomach produces hydrochloricacid, commonly known as “stomach acid.” It isgenerated in the digestion process, but when aperson eats something requiring the stomach towork overtime in digesting it—say, a pizza—thestomach may generate excess hydrochloric acid,



and the result is “heartburn.” When this happens,people often take antacids, which contain a basesuch as aluminum hydroxide (Al[OH]3) or mag-nesium hydroxide (Mg[OH]2).

When a person takes an antacid, the reactionleads to the creation of a salt, but not the saltwith which most people are familiar—NaCl. Asshown above, that particular salt is the product ofa reaction between hydrochloric acid and sodiumhydroxide, but a person who ingested sodiumhydroxide (a substance used to unclog drains andclean ovens) would have much worse heartburnthan before! In any case, the antacid reacts withthe stomach acid to produce a salt, as well aswater, and thus the acid is neutralized.

When land formerly used for mining isreclaimed, the acidic water in the area must beneutralized, and the use of calcium oxide (CaO)as a base is one means of doing so. Acidic soil,too, can be neutralized by the introduction ofcalcium carbonate (CaCO3) or limestone, alongwith magnesium carbonate (MgCO3). If soil istoo basic, as for instance in areas where there hasbeen too little precipitation, acid-like substancessuch as calcium sulfate or gypsum (SaSO4) canbe used. In either case, neutralization promotesplant growth.

TITRATION AND pH. One of themost important applications of neutralization isin titration, the use of a chemical reaction todetermine the amount of a chemical substance ina sample of unknown purity. In a typical form ofneutralization titration, a measured amount ofan acid is added to a solution containing anunknown amount of a base. Once enough of theacid has been added to neutralize the base, it ispossible to determine how much base exists inthe solution.

Titration can also be used to measure pH(“power of hydrogen”) level by using an acid-base indicator. The pH scale assigns values rang-ing from 0 (a virtually pure acid) to 14 (a virtu-ally pure base), with 7 indicating a neutral substance. An acid-base indicator such as litmus paper changes color when it neutralizesthe solution.

The transition interval (the pH at which thecolor of an indicator changes) is different for dif-ferent types of indicators, and thus various indi-cators are used to measure substances in specificpH ranges. For instance, methyl red changes

from red to yellow across a pH range of 4.4 to 6.2,so it is most useful for testing a substance sus-pected of being moderately acidic.

BUFFERED SOLUTIONS. Abuffered solution is one that resists a change inpH even when a strong acid or base is added to it.This buffering results from the presence of aweak acid and a strong conjugate base, and it canbe very important to organisms whose cells canendure changes only within a limited range ofpH values. Human blood, for instance, containsbuffering systems, because it needs to be at pHlevels between 7.35 and 7.45.

The carbonic acid-bicarbonate buffer sys-tem is used to control the pH of blood. The mostimportant chemical equilibria (that is, reactionsinvolving chemical equilibrium) for this systemare: H+ + HCO3

– ! H2CO3 ! H20 + CO2. Inother words, the hydrogen ion (H+) reacts withthe hydrogen carbonate ion (HCO3

– ) to producecarbonic acid (H2CO3). The latter is in equi-librium with the first set of reactants, as well aswith water and carbon dioxide in the forwardreaction.

The controls the pH level by changing theconcentration of carbon dioxide by exhalation.In accordance with Le Châtelier’s principle, thisshifts the equilibrium to the right, consuming H+

ions. In normal blood plasma, the concentrationof HCO3

– is about 20 times as great as that ofH2CO3, and this large concentration of hydrogencarbonate gives the buffer a high capacity to neu-tralize additional acid. The buffer has a muchlower capacity to neutralize bases because of themuch smaller concentration of carbonic acid.

Water: Both Acid and Base

Water is an amphoteric substance; in otherwords, it can serve either as an acid or a base.When water experiences ionization, one watermolecule serves as a Brønsted-Lowry acid,donating a proton to another water molecule—the Brønsted-Lowry base. This results in the pro-duction of a hydroxide ion and a hydronium ion:H2O(l) + H2O(l) ! H3O+(aq) + OH–(aq).

This equilibrium equation is actually one inwhich the tendency toward the reverse reaction ismuch greater; therefore the equilibrium symbol,if rendered in its most proper form, would showa much shorter arrow pointing toward the right.In water purified by distillation, the concentra-


Acid-BaseReactions


Acid-BaseReactions


ACID: A substance that in its edible

form is sour to the taste, and in non-edible

forms is often capable of dissolving metals.

Acids and bases react to form salts and

water. These are all phenomenological def-

initions, however, in contrast to the three

structural definitions of acids and bases—

the Arrhenius, Brønsted-Lowry, and Lewis

acid-base theories.

ALKALI: A term referring to the soluble

hydroxides of the alkali and alkaline earth

metals. Once “alkali” was used for the class

of substances that react with acids to form

salts; today, however, the more general

term base is preferred.

ALKALINITY: An adjectival term used

to identify the degree to which a substance

displays the properties of a base.

AMPHOTERIC: A term describing a

substance that can serve either as an acid or

a base. Water is the most significant

amphoteric substance.

AQUEOUS SOLUTION: A substance

in which water constitutes the solvent. A

large number of chemical reactions take

place in an aqueous solution.

ARRHENIUS ACID-BASE THEORY:

The first of three structural definitions of

acids and bases. Formulated by Swedish

chemist Svante Arrhenius (1859-1927), the

Arrhenius theory defines acids and bases

according to the ions they produce in an

aqueous solution an acid produces hydro-

gen ions (H+), and a base hydroxide ions

(OH–).

BASE: A substance that in its edible

form is bitter to the taste. Bases tend to be

slippery to the touch, and in reaction with

acids they produce salts and water. Bases

and acids are most properly defined, how-

ever, not in these phenomenological terms,

but by the three structural definitions of

acids and bases—the Arrhenius, Brønsted-

Lowry, and Lewis acid-base theories.

BASIC: In the context of acids and

bases, the word is the counterpart to

“acidic,” identifying the base-like quality of

a substance.

BRØNSTED-LOWRY ACID-BASE THE-

ORY: The second of three structural

definitions of acids and bases. Formulated

by English chemist Thomas Lowry (1874-

1936) and Danish chemist J. N. Brønsted

(1879-1947), Brønsted-Lowry theory

defines an acid as a proton (H+) donor, and

a base as a proton acceptor.





so forth.

CONJUGATE ACID: An acid formed

when a base accepts a proton (H+).

CONJUGATE ACID-BASE PAIR:

The acid and base produced when an acid

donates a single proton to a base. In the

reaction that produces this pair, the acid

and base switch identities. By donating a

proton, the acid becomes a conjugate base,

and by receiving the proton, the base

becomes a conjugate acid.

CONJUGATE BASE: A base formed

when an acid releases a proton.

DISSOCIATION: The separation of

molecules into ions.

KEY TERMS


Acid-BaseReactions




a net electric charge. There are two types of

ions: anions and cations.




charges.



contain at least one metal and nonmetal

joined by an ionic bond.

LEWIS ACID-BASE THEORY: The

third of three structural definitions of

acids and bases. Formulated by American

chemist Gilbert N. Lewis (1875-1946),

Lewis theory defines an acid as the reactant

that accepts an electron pair from another

reactant in a chemical reaction, and a base

as the reactant that donates an electron

pair to another reactant.

NEUTRALIZATION: The process

whereby an acid and base react with one

another to form a salt and water.

PH SCALE: A logarithmic scale for

determining the acidity or alkalinity of a

substance, from 0 (virtually pure acid) to 7

(neutral) to 14 (virtually pure base).




only convey part of the picture, however—








tion, resulting in the creation of a product.

SALTS: Ionic compounds formed by

the reaction between an acid and a base. In

this reaction, one or more of the hydrogen

ions of an acid is replaced with another

positive ion. In addition to producing salts,

acid-base reactions produce water.



solute) is dissolved in another substance

(the solvent)—for example, sugar dis-

solved in water.

SOLVENT: A substance that dissolves

another, called a solute, in a solution.

STRONG ELECTROLYTE: A substance

highly prone to dissociation. The terms

“strong acid” or “strong base” refers to those

acids or bases which readily dissociate.



molecular structure and behavior rather

than purely phenomenological data.

TITRATION: The use of a chemical

reaction to determine the amount of a

chemical substance in a sample of

unknown purity. Testing pH levels is an

example of titration.

TRANSITION INTERVAL: The pH

level at which the color of an acid-base

indicator changes.

WEAK ELECTROLYTE: A substance

that experiences little or no dissociation.

The terms “weak acid” or “weak base” re-

fer to those acids or bases not prone to

dissociation.



Acid-BaseReactions

tions of hydronium (H3O+) and hydroxide

(OH–) ions are equal. When multiplied by one

another, these yield the constant figure 1.0 • 10–14,

which is the equilibrium constant for water. In

fact, this constant—denoted as Kw—is called the

ion-product constant for water.

Because the product of these two concentra-

tions is always the same, this means that if one of

them goes up, the other one must go down in

order to yield the same constant. This explains

the fact, noted earlier, that water can serve either

as an acid or base—or, if the concentrations of

hydronium and hydroxide ions are equal—as a

neutral substance. In situations where the con-

centration of hydronium is higher, and the

hydroxide concentration automatically decreas-

es, water serves as an acid. Conversely, when the

hydroxide concentration is high, the hydronium

concentration decreases correspondingly, and

the water is a base.

WHERE TO LEARN MORE

“Acids and Bases.” Vision Learning (Web site).<http://www.visionlearning.com/library/science/chemistry-2/CHE2.2-acid_base.htm> (June 7, 2001).

“Acids, Bases, and Chemical Reactions” Open Access Col-lege (Web site). <http://oac.schools.sa.edu.au/8-10science/acids.htm> (June 7, 2001).

“Acids, Bases, pH.” About.com (Web site). <http://chemistry.about.com/science/chemistry/cs/acidsbasesph/ (June 7, 2001).

“Acids, Bases, and Salts” (Web site). <http://edie.cprost.sfu.ca/~rhlogan/ionic_eq.html> (June 7, 2001).

“Junior Part: Acids, Bases, and Salts” (Web site).<http://www.rjclarkson.demon.co.uk/junior/junior4.htm> (June 7, 2001).

Knapp, Brian J. Acids, Bases, and Salts. Danbury, CT:Grolier Educational, 1998.

Moje, Steven W. Cool Chemistry: Great Experiments withSimple Stuff. New York: Sterling Publishing Compa-ny, 1999.

Walters, Derek. Chemistry. Illustrated by Denis Bishopand Jim Robins. New York: F. Watts, 1982.




327


SOLUT IONS ANDM IX TURES

SOLUT IONS ANDM IX TURES

MIXTURES

SOLUT IONS

OSMOS IS

D IST I L LAT ION AND F I LTRAT ION



M I X TURESMixtures

CONCEPTElements and compounds are pure substances,

but much of the material around us—including

air, wood, soil, and even (in most cases) water—

appears in the form of a mixture. Unlike a pure

substance, a mixture has variable composition: in

other words, it cannot be reduced to a single type

of atom or molecule. Mixtures can be either

homogeneous—that is, uniform in appearance,

and having only one phase—or heterogeneous,

meaning that the mixture is separated into vari-

ous regions with differing properties. Most

homogeneous mixtures are also known as solu-

tions, and examples of these include air, coffee,

and even metal alloys. As for heterogeneous mix-

tures, these occur, for instance, when sand or oil

are placed in water. Oil can, however, be dis-

persed in water as an emulsion, with the aid of a

surfactant or emulsifier, and the result is an

almost homogeneous mixture. Another interest-

ing example of a heterogeneous mixture occurs

when colloids are suspended in fluid, whether

that fluid be air or water.

HOW IT WORKS

Mixtures and Pure Substances

A mixture is a substance with a variable compo-

sition, meaning that it is composed of molecules

or atoms of differing types. These may have

intermolecular bonds, or bonds between mole-

cules, but such bonds are not nearly as strong as

those formed when elements bond to form a

compound.

Examples of mixtures include milk, coffee,tea, and soft drinks—indeed, they include virtu-ally any drink except for pure water, which is ararity, because water is highly soluble and tendsto contain minerals and other impurities. Air,too, is a mixture, because it contains oxygen,nitrogen, noble gases, carbon dioxide, and watervapor. Likewise metal alloys such as bronze,brass, and steel are mixtures.

STRUCTURAL AND PHENOME-

NOLOGICAL DISTINCTIONS. Be-fore chemists accepted the atomic theory of mat-ter, it was often difficult to distinguish a mixture,such as air or steel, from pure substances, whichhave the same composition throughout. Puresubstances are either elements such as oxygen oriron, or compounds—for example, carbon diox-ide or iron oxide. An element is made up of onlyone type of atom, while a compound can bereduced to one type of bonded chemicalspecies—usually either a molecule or a forma-tion of ions.

A distinction between mixtures and puresubstances based on atomic and molecular struc-tures is called, fittingly enough, a structural defi-nition. Prior to the development of atomic theo-ry by English chemist John Dalton (1766-1844),and the subsequent identification of moleculesby Italian physicist Amedeo Avogadro (1776-1856), chemists defined compounds and mix-tures on the basis of phenomenological data garnered through observation and measure-ment. This could and did lead to incorrect suppositions.

PROUST AND THE DEBATE

OVER CONSTANT COMPOSITION.

Around the time of Dalton and Avogadro, French


Mixtures


chemist Antoine Lavoisier (1743-1794) correctlyidentified an element as a substance that couldnot be broken down into a simpler substance.Later, his countryman Joseph-Louis Proust(1754-1826) clarified the definition of a com-pound, stating it in much the same terms that wehave used: the composition of a compound is thesame throughout. But with the very rudimentaryknowledge of atoms and molecules that chemistsof the early nineteenth century possessed, thedistinctions between compounds and mixturesremained elusive.

Proust’s contemporary and intellectual rivalClaude Louis Berthollet (1748-1822) observedthat when metals such as iron are heated, theyform oxides in which the percentage of oxygenincreases with temperature. If constant composi-tion is a fact, he challenged Proust, then how isthis possible? Proust maintained that the propor-tions of elements in a compound are always thesame, a hypothesis Berthollet countered byobserving that metals can form variable alloys.Bronze, for instance, is an alloy of tin and copperat a ratio of about 25:75, but if the ratio werechanged to, say, 30:70, it would still be calledbronze.

At the time, the equipment—both in termsof knowledge and tools for analysis—simply did

not exist whereby Berthollet’s assertions could besuccessfully proven wrong. Indeed, Berthollet, astudent of Lavoisier who contributed to progressin the study of acids and reactants, was no anti-scientific villain: he was simply responding to theevidence as he saw it. Only with the discovery ofsubatomic particles in the late nineteenth andearly twentieth centuries, discoveries thatchanged the entire character of chemistry as adiscipline, did it truly become possible to answerBerthollet’s challenges.

Chemists now understand that the oxideformed on the surface of a metal is a compoundseparate from the metal itself, and that alloys(discussed later in this essay) are not compoundsat all: they are mixtures.

Mixtures and Compounds

With our modern knowledge of atomic andmolecular properties, we can more fully distin-guish between a compound and a mixture. First,a mixture can exist in virtually any proportionamong its constituent parts, whereas a com-pound has a definite and constant composition.Coffee, whether weak or strong, is still coffee, butif a second oxygen atom chemically bonds to theoxygen in carbon monoxide (CO), the resulting

THE WONDROUS INVENTION KNOWN AS SOAP IS A MIXTURE WITH SURFACTANT QUALITIES. (Chad Weckler/Corbis. Reproduced

by permission.)


Mixtures


compound is not simply “strong carbon monox-ide” or even “oxygenated carbon monoxide”: it iscarbon dioxide (CO2), an entirely different sub-stance. The difference is enormous: carbon diox-ide is a part of the interactions between plant andanimal life that sustains the natural environment,whereas carbon monoxide is a toxic gas pro-duced, for instance, by automobiles.

Second, the parts of a mixture retain most oftheir properties when they join together, but ele-ments lose their elemental characteristics oncethey form a compound. Sugar is still sweet,regardless of the substance into which it is dis-solved. In coffee or another drink, it may nolonger be dry and granular, though it could bereturned to that state by evaporating the coffee.In any case, dryness and granularity are physicalproperties, whereas sweetness is a chemical one.On the other hand, when elemental carbonbonds with hydrogen and oxygen to form sugaritself, the resulting compound is nothing like theelements from which it is made.

Third, a mixture can usually be created sim-ply by the physical act of stirring items together,and/or by applying heat. As discussed below, rel-atively high temperatures are sometimes neededto dissolve sugar in tea, for instance; likewise,alloys are usually formed by heating metals to

much higher temperatures. On the other hand,the formation of a compound involves a chemi-cal reaction, a vastly more complex interactionthan the one required to create a mixture. Car-bon, a black solid found in coal, can be mixedwith hydrogen and oxygen—both colorless andodorless gases—and the resulting mess would besomething; but it would not be sugar. Only thechemical reaction between the three elements, inspecific proportions and under specific condi-tions, results in their chemical bonding to formtable sugar, known to chemists as sucrose.

Homogeneous and Heteroge-neous Mixtures

HOMOGENEOUS MIXTURES

AND SOLUTIONS. As we have seen, thecomposition of a mixture is variable. Coffee, forinstance, can be weak or strong, and milk can be“whole” or low-fat. Beer can be light or dark incolor, as well as “light” in terms of calorie con-tent; furthermore, its alcohol content can be var-ied, as with all alcoholic drinks. Though theirmolecular composition is variable, each of themixtures described here is the same throughout:in other words, there is no difference betweenone region and another in a glass of beer. Such amixture is described as homogeneous.

BRONZE—AN ALLOY OF TIN AND COPPER—WAS ONE OF THE MOST IMPORTANT INVENTIONS IN HUMAN HISTORY. THIS

STONE CARVING IS FROM THE PERIOD KNOWN AS THE BRONZE AGE. (Kevin Schafer/Corbis. Reproduced by permission.)


Mixtures A homogeneous mixture is one that is thesame throughout, but this is not the same as say-ing that a compound has a constant composi-tion. Rather, when we say that a homogeneousmixture is the same in every region, we are speak-ing phenomenologically rather than structurally.From the standpoint of ordinary observation, aglass of milk is uniform, but as we shall see, milkis actually a type of emulsion, in which one sub-stance is dispersed in another. There are no milk“molecules,” and structural analysis would revealthat milk does not have a constant molecularcomposition.

Coffee is an example of a solution, a specifictype of homogeneous mixture. Actually, mosthomogeneous mixtures can be considered solu-tions, but since a solution is properly defined as ahomogeneous mixture in which one or moresubstances is dissolved in another substance, a50:50 homogeneous mixture is technically not asolution. When particles are perfectly dissolvedin a solution, the composition is uniform, butagain, this is not a compound, because that com-position can always be varied.

HETEROGENEOUS MIXTURES.

Anyone who has ever made iced tea has observedthat it is easy to sweeten when it is hot, becausetemperature affects the ability of a solvent suchas water—in this case, water with particles fromtea leaves filtered into it—to dissolve a solute,such as sugar. Cold tea, on the other hand, ismuch harder to sweeten with ordinary sugar.Usually what happens is that, instead of obtain-ing a homogeneous mixture, the result is hetero-geneous.

Whereas every region within a homoge-neous mixture is more or less the same as everyother region, heterogeneous mixtures containregions that differ from one another. In theexample we have used, a glass of cold tea withundissolved sugar at the bottom is a heteroge-neous mixture: the tea at the top is unsweetenedor even bitter, whereas at the bottom, there is anoverly sweet sludge of tea and sugar.

FROM HOMOGENEOUS TO

HETEROGENEOUS AND BACK

AGAIN. The distinction between homoge-neous and heterogeneous solutions only serves tofurther highlight the difference between com-pounds and mixtures. Whereas the compositionof a compound is definite and quantitative (somany atoms of element x bonded with so many

atoms of element y in such-and-such a struc-ture), there is less of a sharp dividing linebetween homogeneous and heterogeneous solutions.

When too much of a solute is added to thesolvent, the solvent becomes saturated, and is nolonger truly homogeneous. Such is the case whenthe sugar drifts to the bottom of the tea, or whencoffee grounds form in the bottom of a coffee potcontaining an otherwise uniform mixture of cof-fee and water.

Milk comes to us as a homogeneous sub-stance, thanks to the process of homogenization,but when it comes out of the cow, it is a moreheterogeneous mixture of milk fat and water. Infact, milk is an example of an emulsion, the resultof a process whereby two substances that wouldotherwise form a heterogeneous mixture aremade to form a homogeneous one.


Colloids

In 1827, Scottish naturalist Robert Brown (1773-1858) was studying pollen grains under a micro-scope when he noticed that the grains underwenta curious zigzagging motion in the water. At first,he assumed that the motion had a biologicalexplanation—in other words, that it resultedfrom life processes within the pollen—but later,he discovered that even pollen from long-deadplants behaved in the same way.

What Brown had observed was the suspen-sion of a colloid—a particle intermediate in sizebetween a molecule and a speck of dust—in thewater. When placed in water or another fluid(both liquids and gases are called “fluids” in thephysical sciences), a colloid forms a mixture sim-ilar both to a solution and a suspension. As witha solution, the composition remains homoge-neous—the particles never settle to the bottomof the container. At the same time, however, theparticles remain suspended rather than dis-solved, rendering a cloudy appearance to the dispersion.

Almost everyone, especially as a child, hasbeen fascinated by the colloidal dispersion ofdust particles in a beam of sunlight. They seem tobe continually in motion, as indeed they are, andthis movement is called “Brownian motion” in



Mixtureshonor of the man who first observed it. YetBrown did not understand what he was seeing;only later did scientists recognize that Brownianmotion is the result of movement on the part ofmolecules in a fluid. Even though the moleculesare much smaller than the colloid, there areenough of them producing enough collisions tocause the colloid to be in constant motion.

Another remarkable aspect of dust particlesfloating in sunlight is the way that they cause acolumn of sunlight to become visible. This phe-nomenon, called the Tyndall effect after Englishphysicist John Tyndall (1820-1893), makes itseem as though we are actually seeing a beam oflight. In fact, what we are seeing is the reflectionof light off the colloidal dispersion. Anotherexample of a colloidal dispersion occurs when apuff of smoke hangs in the air; furthermore, aswe shall see, milk is a substance made up of col-loids—in this case, particles of fat and proteinsuspended in water.

Emulsions

Miscibility is a qualitative term identifying therelative ability of two substances to dissolve inone another. Generally, water and water-basedsubstances have high miscibility with regard toone another, as do oil and oil-based substanceswith one another. Oil and water, on the otherhand (or substances based in either) have verylow miscibility.

The reason is that water molecules are polar,meaning that the positive electric charge is in oneregion of the molecule, while the negative chargeis in another region. Oil, on the other hand, isnonpolar, meaning that charges are more evenlydistributed throughout the molecule. Trying tomix oil and water is thus rather like attempting touse a magnet to attract a nonmagnetic metalsuch as gold.

HOW SURFACTANTS AID THE

EMULSION PROCESS. An emulsion isa mixture of two immiscible liquids such as oiland water—we will use these simple terms, withthe idea that these stand for all oil- and water-soluble substances—by dispersing microscopicdroplets of one liquid in another. In order toachieve this dispersion, there needs to be a “mid-dle man,” known as an emulsifier or surfactant.The surfactant, made up of molecules that areboth water- and oil-soluble, acts as an agent forjoining other substances in an emulsion.

All liquids have a certain degree of surfacetension. This is what causes water on a hard sur-face to bead up instead of lying flat. (Mercury hasan extremely high surface tension.) Surfactantsbreak down this surface tension, enabling themarriage of two formerly immiscible liquids.

In an emulsion, millions of surfactants sur-round the dispersed droplets (known as theinternal phase), shielding them from the otherliquid (the external phase). Supposing oil is the internal phase, then the oil-soluble end of thesurfactant points toward the oils, while the water-soluble side joins with the water of theexternal phase.

PHYSICAL CHARACTERISTICS

OF AN EMULSION. The resulting emul-sion has physical (as opposed to chemical) prop-erties different from those of the substances thatmake it up. Water and most oils are transparent,whereas emulsions are typically opaque and mayhave a lustrous, pearly gleam. Oil and water flowfreely, but emulsions can be made to form thick,non-flowing creams.

Emulsions are inherently unstable: they are,as it were, only temporarily homogeneous. Agood example of this is an oil-and-vinegar saladdressing, which has to be shaken before putting iton salad; otherwise, what comes out of the bottleis likely to be mainly oil, which floats to the topof the water-based vinegar. It is characteristic ofemulsions that eventually gravity pulls themapart, so that the lighter substances float to thetop. This process may take only a few minutes, aswith the salad dressing; on the other hand, it cantake thousands of years.

APPLICATIONS. Surfactants them-selves are often used in laundry or dish detergent,because most stains on plates or clothes are oil-based, whereas the detergent itself is applied tothe clothes in an aqueous solution. As for emul-sions, these are found in a wide variety of prod-ucts, from cosmetics to paint to milk.

In the pharmaceutical industry, for instance,emulsions are used as a means of delivering theactive ingredients of drugs, many of which arenot water soluble, in a medium that is. Pharma-ceutical manufacturers also use emulsions tomake medicines more palatable (easier to take);to control dosage and improve effectiveness; andto improve the usefulness of topical drugs such asointments, which must contend with both oilyand water-soluble substances on the skin.



Mixtures The oils and other conditioning agents usedin hair-care products would leave hair limp andsticky if applied by themselves. Instead, theseproducts are applied in emulsions that dilute theoils in other materials. Emulsions are also used incolor photography, as well as in the productionof pesticides and fungicides.

Paints and inks, too, are typically emulsifiedsubstances, and these sometimes come in theform of dispersions akin to those of colloids. Aswe have seen, in a dispersion fine particles ofsolids are suspended in a liquid. Surfactants areused to bond the particles to the liquid, thoughpaint must be shaken, usually with the aid of ahigh-speed machine, before it is consistentenough to apply.

As mentioned earlier, milk is a form ofemulsion that contains particles of milk fat orcream dispersed in water. Light strikes the micro-scopic fat and protein colloids, resulting in thefamiliar white color. Other examples of emulsi-fied foods include not only salad dressings, butgravies and various sauces, peanut butter, icecream, and other items. Not only does emulsionaffect the physical properties of foods, but it mayenhance the taste by coating the tongue withemulsified oils.

Soap

The wondrous invention called soap, which hasmade the world a cleaner place, exhibits several ofthe properties we have discussed. It is a mixturewith surfactant qualities, and in powdered form,it can form something close to a true solutionwith water. Discovered probably by the Phoeni-cians around 600 B.C., it was made in ancienttimes by boiling animal fat (tallow) or vegetableoils with an alkali or base of wood ashes.

This was a costly process, and for many cen-turies, only the wealthy could afford soap—a factthat contributed to the notorious lack of person-al hygiene that prevailed among Europeans of theMiddle Ages. In fact, the situation did not changeuntil late in the eighteenth century, when thework of two French chemists yielded a more eco-nomical manufacturing method.

In 1790, Nicholas Leblanc (1742-1806)developed a process for creating sodium hydrox-ide (called caustic soda at the time) from sodiumchloride, or ordinary table salt. This made for aninexpensive base or alkali, with which soap man-ufacturers could combine natural fats and oils.

Then in 1823, Michel Eugène Chevreul (1786-1889) showed that fat is a compound of glycerolwith organic acids, a breakthrough that led to theuse of fatty acids in producing soaps.

THE CHEMISTRY OF SOAP.

Soap as it is made today comes from a salt of analkali metal, such as sodium or potassium, com-bined with a mixture of carboxylic acids (“fattyacids”). These carboxylic acids are the result of areaction called saponification, involving triglyc-erides and a base, such as sodium hydroxide.Saponification breaks the triglycerides into theircomponent fatty acids; then, the base neutralizesthese to salts. This reaction produces glycerin, ahydrocarbon bonded to a hydroxyl (-OH) group.(Hydrocarbons are discussed in the essays onOrganic Chemistry; Polymers.)

In general, the formula for soap is RCOOX,where X stands for the alkali metals, and R for thehydrocarbon chain. Because it is a salt (meaningthat it is formed from the reaction of an acid witha base), soap partially separates into its compo-nent ions in water. The active ion is RCOO-,whose two ends behave in different fashions,making it a surfactant. The hydrocarbon end (R-) is said to be lipophilic, or “oil-loving,” whichis appropriate, since most oils are hydrocarbons.On the other hand, the carboxylate end (-COO-)is hydrophilic, or “water-loving.” As a result, soapcan dissolve in water, but can also clean greasystains.

When soap is mixed with water, it does notform a true homogeneous mixture or solution dueto the presence of hydrocarbons. These attract oneanother, forming spherical aggregates calledmicelles. The lipophilic “tails” of the hydrocarbonsare turned toward the interior of the micelle, whilethe hydrophilic heads remain facing toward thewater that forms the external phase.

Alloys

We have primarily discussed liquid mixtures, butin fact mixtures can be gaseous or even solid—asfor example in the case of an alloy, a mixture oftwo or more metals. Alloys are usually created bymelting the component metals, then mixingthem together in specific proportions, but analloy can also be created by bonding metal powders.

The structure of an elemental metal includestight “electron sea” bonds, which are discussed inthe essay on Metals. These, combined with the



Mixtures


ALLOY: A mixture of two or more

metals.

AQUEOUS: A term describing a solu-

tion in which a substance or substances are

dissolved in water.

ATOM: The smallest particle of an

element.





so forth.






changed chemically.

DISPERSION: A term describing the

distribution of particles in a fluid.


only one kind of atom. Unlike compounds,

elements cannot be chemically broken

down into other substances.

EMULSIFIER: A surfactant.

EMULSION: A mixture of two immis-

cible liquids such as oil and water, made by

dispersing microscopic droplets of one liq-

uid in another. In creating an emulsion,

surfactants act as bridges between the two

liquids.

FLUID: A substance with the ability to

flow. In physical sciences such as chemistry

and physics, “fluid” refers both to liquids

and gases.


ing a mixture that is not the same through-

out; rather, it has various regions possess-

ing different properties. An example of a

mixture is sand in a container of water.

Rather than dissolving to form a homoge-

neous mixture, as sugar would, the sand

sinks to the bottom.


a mixture that is the same throughout, as

for example when sugar is fully dissolved

in water. A solution is a homogenous

mixture.

IMMISCIBLE: Possessing a negligible

value of miscibility.




MISCIBILITY: A qualitative term iden-

tifying the relative ability of two substances

to dissolve in one another.

MIXTURE: A substance with a variable

composition, meaning that it is com-

posed of molecules or atoms of differing

types. Compare with pure substance or

compound.


ly but not always representing more



molecules.




only convey part of the picture, however—






KEY TERMS


Mixtures


metal’s crystalline structure, create a situation in

which internal bonding is very strong, but

nondirectional. As a result, most metals form

alloys, but again, this is not the same as a com-

pound: the elemental metals retain their identity

in these metal composites, forming characteristic

fibers, beads, and other shapes.

SOME FAMOUS ALLOYS. One of

the most important alloys in early human histo-

ry was a 25:75 mixture of tin and copper that

gave its name to an entire stage of technological

development: the Bronze Age (c. 3300-1200 B.C.).

This combination formed a metal much stronger

than either copper or tin, and bronze remained

dominant until the discovery of new iron-smelt-

ing methods ushered in the Iron Age in about

1200 B.C.

Another important alloy of the ancient

world was brass, which is about one-third zinc to

two-thirds copper. Introduced in the period c.

1400-c. 1200 B.C, brass was one of the metals that

defined the world in which the Bible was written.

Biblical references to it abound, as in the famous

passage from I Corinthians 13 (read at many

weddings), in which the Apostle Paul proclaims

that “without love, I am as sounding brass.” Some

biblical scholars, however, maintain that some of

the references to brass in the Old Testament are

actually mistranslations of “bronze.”

Tin, copper, and antimony form pewter, asoft mixture that can be molded when cold andbeaten regularly without turning brittle. Thoughused in Roman times, it became most popular inEngland from the fourteenth to the eighteenthcenturies, when it offered a cheap substitute forsilver in plates, cups, candelabra, and pitchers.The pewter turned out by colonial Americanmetalsmiths is still admired for its beauty andfunctionality.

As for iron, it appears in nature as an impureore, but even when purified, it is typically alloyedwith other elements. Among the well-knownforms of iron are cast iron, a variety of mixturescontaining carbon and/or silicon; wrought iron,which contains small amounts of various otherelements, such as nickel, cobalt, copper, chromi-um, and molybdenum; and steel. Steel is a mix-ture of iron with manganese and chromium,purified with a blast of hot air. Steel is also some-times alloyed with aluminum in a one-third/two-thirds mixture to make duraluminum, developedfor the superstructures of German zeppelins dur-ing World War I.

WHERE TO LEARN MORE

“All About Colloids.” Synthashield (Web site).<http://www.synthashield.net/vault/colloids.html>(June 6, 2001).

Carona, Phillip B. Magic Mixtures: Alloys and Plastics.Englewood Cliffs, NJ: Prentice Hall, 1963.

PURE SUBSTANCE: A substance that

has the same composition throughout.

Pure substances are either elements or

compounds, and should not be confused

with mixtures. A homogeneous mixture is

not the same as a pure substance, because

although it has an unvarying composition,

it cannot be reduced to a single type of

atom or molecule.


in which one or more substances is dis-

solved in another substances—for exam-

ple, sugar dissolved in water.



molecular structure rather than purely

phenomenological data.

SURFACTANT: A substance made up

of molecules that are both water- and oil-

soluble, which acts as an agent for joining

other substances in an emulsion.

SUSPENSION: A term that refers to a

mixture in which solid particles are sus-

pended in a fluid.



MixturesChemLab. Danbury, CT: Grolier Educational, 1998.

“Corrosion of Alloys.” Corrosion Doctors (Web site).<http://www.corrosion-doctors.org/MatSelect/corralloys.htm> (June 6, 2001).

“Emulsions.” Pharmweb (Web site). <http://pharmweb.usc.edu/phar306/Handouts/emulsions/> (June 6,2001).


Knapp, Brian J. Elements, Compounds, and Mixtures.Danbury, CT: Grolier Educational, 1998.

Maton, Anthea. Exploring Physical Science. Upper SaddleRiver, NJ: Prentice Hall, 1997.

Patten, J. M. Elements, Compounds, and Mixtures. VeroBeach, FL: Rourke Book Company, 1995.

The Soap and Detergent Association (Web site).<http://www.sdahq.org/> (June 6, 2001).




SOLUT IONSSolutions

CONCEPTWe are most accustomed to thinking of solutionsas mixtures of a substance dissolved in water, butin fact the meaning of the term is broader thanthat. Certainly there is a special place in chem-istry for solutions in which water—”the univer-sal solvent”—provides the solvent medium. Thisis also true in daily life. Coffee, tea, soft drinks,and even water itself (since it seldom appears inpure form) are solutions, but the meaning of theterm is not limited to solutions involving water.Indeed, solutions do not have to be liquid; theycan be gaseous or solid as well. One of the mostimportant solutions in the world, in fact, is theair we breathe, a combination of nitrogen, oxy-gen, noble gases, carbon dioxide, and watervapor. Nonetheless, aqueous or water-based sub-stances are the focal point of study where solu-tions are concerned, and reactions that take placein an aqueous solution provide an importantarea of study.

HOW IT WORKS

Mixtures

Solutions belong to the category of substancesidentified as mixtures, substances with variablecomposition. This may seem a bit confusing,since among the defining characteristics of asolution are its uniformity and consistency. Butthis definition of a solution involves externalqualities—the things we can observe with oursenses—whereas a mixture can be distinguishedfrom a pure substance not only in terms of exter-nal characteristics, but also because of its internaland structural properties.

At one time, chemists had a hard time dif-ferentiating between mixtures and pure sub-stances (which are either elements or com-pounds), primarily because they had little or noknowledge of those “internal and structuralproperties.” Externally, a mixture and compoundcan be difficult to distinguish, but the “variablecomposition” mentioned above is a key factor.Coffee can be almost as weak as water, or sostrong it seems to galvanize the throat as it goesdown—yet in both cases, though its compositionhas varied, it is still coffee.

By contrast, if a compound is altered by thevariation of just one atom in its molecular struc-ture, it becomes something else entirely. When anoxygen atom bonds with a carbon and oxygenatom, carbon monoxide—a poisonous gas pro-duced by the burning of fossil fuels—is turnedinto carbon dioxide, an essential component ofthe interaction between plants and animals in thenatural environment.

MIXTURES VS. COMPOUNDS:

THREE DISTINCTIONS. Air (a mix-ture) can be distinguished from a compound,such as carbon dioxide, in three ways. First, amixture can exist in virtually any proportionsamong its constituent parts, whereas a com-pound has a definite and constant composition.Air can be oxygen-rich or oxygen-poor, but if wevary the numbers of oxygen atoms chemicallybonded to form a compound, the result is anentirely different substance.

Second, the parts of a mixture retain most oftheir properties when they join together, but ele-ments—once they form a compound—lose theirelemental characteristics. If pieces of carbon werefloating in the air, their character would be quite


Solutions


different from that of the carbon in carbon diox-ide. Carbon, in its natural form, is much like coalor graphite (graphite is pure carbon, and coal animpure form of the element), but when a carbonatom combines with an oxygen atom, the twoform a gas. Pieces of carbon in the air, on theother hand, would be solid, as carbon itself is.

Third, a mixture can usually be created sim-ply by the physical act of stirring items together,and/or by applying heat. Certainly this is true ofmany solutions, such as brewed coffee. On theother hand, a compound such as carbon dioxidecan only be formed by the chemical reaction ofatoms in specific proportions and at specifictemperatures. Some such reactions consumeheat, whereas others produce it. These interac-tions of heat are highly complex; on the otherhand, in making coffee, heat is simply producedby an external source (the coffee maker) andtransferred to the coffee in the brewing process.

HETEROGENEOUS AND HOMO-

GENEOUS MIXTURES. There are twobasic types of mixtures. One of these occurs, forinstance, when sand is added to a beaker ofwater: the sand sinks to the bottom, and the com-position of the sand-water mixture cannot be

said to be the same throughout. It is thus a het-erogeneous mixture. The same is true of cold teawhen table sugar (sucrose) is added to it: thesugar drifts to the bottom, and as a result, thefirst sip of tea from the glass is not nearly as sweetas the last.

Wood is a mixture that typically appears inheterogeneous form, a fact that can be easily con-firmed by studying wood paneling. There areknots, for instance—areas where branches oncegrew, and where the concentrations of sap arehigher. Striations of color run through theboards of paneling, indicating complex varia-tions in the composition of the wood. Much thesame is true of soil, with a given sample varyingwidely depending on a huge array of factors: theconcentrations of sand, rock, or decayed vegeta-tion, for instance, or the activities of earthwormsand other organisms in processing the soil.

In a homogeneous mixture, by contrast,there is no difference between one area of themixture and another. If coffee has been brewedproperly, and there are no grounds at the bottomof the pot, it is homogeneous. If sweetener isadded to tea while it is hot, it too should yield ahomogeneous mixture.

WATER AND OIL DROPLETS EXPOSED TO POLARIZED LIGHT. WATER AND OIL HAVE VERY LITTLE MISCIBILITY WITH

REGARD TO ONE ANOTHER, MEANING THAT THEY TEND TO REMAIN SEPARATE WHEN COMBINED. (Science Pictures

Limited/Corbis. Reproduced by permission.)


SolutionsMISCIBILITY AND SOLUBILI-

TY. Miscibility is a term describing the relativeability of two substances to dissolve in oneanother. Generally, water and water-based sub-stances have high miscibility with regard to oneanother, as do oil and oil-based substances. Oiland water, on the other hand (or substancesbased in either) have very low miscibility. Thereasons for this will be discussed below.

Miscibility is a qualitative term like “fast” or“slow”; on the other hand, solubility—the maxi-mum amount of a substance that dissolves in agiven amount of solvent at a specific tempera-ture—can be either qualitative or quantitative. Ina general sense, the word is qualitative, referringto the property of being soluble, or able to dis-solve in a solution.

At some points in this essay, “solubility” willbe used in this qualitative sense; however, withinthe realm of chemistry, it also has a quantitativemeaning. In other words, instead of involving animprecise quality such as “fast” or “slow,” solubil-ity in this sense relates to precise quantities morealong the lines of “100 MPH (160.9 km/h).” Sol-ubility is usually expressed in grams (g) (0.0022lb) of solute per 100 g (0.22 lb) of solvent.

OTHER QUANTITATIVE TERMS.

Another quantitative means of describing a solu-tion is in terms of mass percent: the mass of thesolute divided by the mass of the solution, andmultiplied by 100%. If there are 25 g of solute ina solution of 200 g, for instance, 25 would bedivided by 200 and multiplied by 100% to yield a12.5% figure of solution composition.

Molarity also provides a quantitative meansof showing the concentration of solute to solu-tion. Whereas mass percent is, as its name indi-cates, a comparison of the mass of the solute tothat of the solvent, molarity shows the amount ofsolute in a given volume of solution. This ismeasured in moles of solute per liter of solution,abbreviated mol/l.


Saturation and Dilution

The quantitative terms for describing solubilitythat we have reviewed are useful to a profession-al chemist, or to anyone performing work in alaboratory. For the most part, however, qualita-


Solutions and Solubility

Virtually all varieties of homogeneous mixtures

can be classified as a solution—that is, a homo-

geneous mixture in which one or more sub-

stances is dissolved in another substance. If two

or more substances were combined in exactly

equal proportions, this would technically not

be a solution; hence the distinction between

solutions and the larger class of homogeneous

mixtures.

A solution is made up of two parts: the

solute, the substance or substances that are dis-

solved; and the solvent, the substance that dis-

solves the solute. The solvent typically deter-

mines the physical state of the solution: in other

words, if the solvent is a liquid, the solution will

most likely be a liquid. Likewise if the solvent is a

solid such as a metal, melted at high tempera-

tures to form a metallic solution called an alloy,

then when the solution is cooled to its normal

temperature, it will be solid again.

WATER AND ETHANOL ARE COMPLETELY MISCIBLE,WHETHER THE SOLUTION BE 99% WATER AND 1%ETHANOL, 50% OF BOTH, OR 99% ETHANOL AND 1%WATER. SHOWN HERE IS A GLASS OF BEER—AN

EXAMPLE OF A WATER-ETHANOL SOLUTION. (Owen

Franken/Corbis. Reproduced by permission.)


Solutionstive expressions will suffice for the present dis-cussion. Among the most useful qualitative termsis saturation.

When it is possible to dissolve solute in a sol-vent, the solution is said to be unsaturated—rather like a sponge that has not been filled tocapacity with liquid. By contrast, a solution thatcontains as much solute as it can at a given tem-perature is like a sponge that has been filled withall the liquid it can possibly contain. It can nolonger absorb more of the liquid; rather, it canonly push liquid particles along, and the spongeis said to be saturated. In the same way, if tea ishot, it is easier to introduce sugar to it, but whenthe temperature is low, sugar will not dissolve aseasily.

SATURATION AND TEMPERA-

TURE. With tea and many other substances,higher temperatures mean that a greater amountof solute can be added before the solution is fullysaturated. The reason for this relationshipbetween saturation and temperature, accordingto generally accepted theory, is that heated sol-vent particles move more quickly than cold ones,and as a result, create more space into which thesolvent can fit. Indeed, “space” is a prerequisitefor a solution: the molecules of solute need tofind a “hole” between the molecules of solventinto which they can fit. Thus the molecules in asolution can be compared to a packed crowd: if acrowd is suddenly dispersed, it is easier to walkthrough it.

Not all substances respond to rises in tem-perature in the manner described, however. As arule, gases (with the exception of helium) aremore soluble at lower temperatures than at high-er ones. Higher temperatures mean an increasein kinetic energy, with more molecules collidingwith one another, and this makes it easier for the gases to escape the liquid in which they aredissolved.

Carbonated soft drinks get their “fizz” fromcarbon dioxide gas dissolved, along with sugarand other flavorings, in a solution of water. Whenthe soft-drink container is opened, much of thecarbon dioxide quickly departs, while a certainproportion stays dissolved in the cola. Theremaining carbon dioxide will eventually escapeas well, but it will do so more quickly if left in awarm environment rather than a cold one, suchas the inside of a refrigerator.

DILUTION AND CONCENTRA-

TION. Another useful set of qualitative termsfor solutions also relates to the relative amount ofsolute that has been dissolved. But whereas satu-ration is an absolute condition at a certain tem-perature (in other words, the solution is definite-ly saturated at such-and-such a temperature),concentration and dilution are much more rela-tive terms.

When a solution contains a relatively smallamount of solute, it is said to be dilute; on theother hand, a solution with a relatively largeamount of solute is said to be concentrated. Forinstance, hydrogen peroxide (H2O2) as sold overthe counter in drug stores, to be used in homes asa disinfectant of bleaching agent, is a dilute 3%solution in water. In higher concentrations,hydrogen peroxide can cause burns to humanskin, and is used to power rockets.

Chemists often keep on hand in their labo-ratories concentrated versions of various fre-quently used solutions. Maintaining them inconcentrated form saves space; when more of asolution is needed, the chemist dilutes it in wateraccording to desired molarity values. To do so,the chemist determines how much water needs tobe added to a particular “stock solution” (as theseconcentrated solutions in a laboratory are called)to create a solution of a particular concentration.In the dilution with water, the number of molesof the solute does not change; rather, the particlesof solute are dispersed over a larger volume.

“Like Dissolves Like”

In the molecules of a compound, the atoms areheld together by powerful attractions. The mole-cules within a compound, however, also have anintermolecular attraction, but this is much weak-er than the interatomic bonds in a molecule. It isthus quite possible for the molecules in a solu-tion to exert a greater attractive force than theone holding the molecules of solute togetherwith one another. When this happens, the solutedissolves.

A useful rule of thumb when studying solu-tions is “like dissolves like.” In other words, waterdissolves water-based solutes, whereas oil dis-solves oily substances. Or to put it another way,water and water-based substances are misciblewith one another, as are oil and oil-based sub-stances. Water and oil together, however, arelargely immiscible.



SolutionsPOLAR WATER, NONPOLAR

OIL. Water and oil and immiscible due to theirrespective molecule structures, and hence theirinherent characteristics of intermolecular bond-ing. Water molecules are polar, meaning that thepositive electric charge is at one end of the mole-cule, while the negative charge is at the other end.Oil, on the other hand, is nonpolar—charges are more evenly distributed throughout the molecule.

The oxygen atom in a water molecule has amuch greater electronegativity than the hydro-gen atoms, so it pulls negatively charged elec-trons to its end of the molecule. Conversely, mostoils are composed of carbon and hydrogen,which have similar values of electronegativity. Asa result, positive and negative charges are distrib-uted evenly throughout the oil molecule. Where-as a water molecule is like a magnet with a northand south pole, an oil molecule is like a nonmag-netic metal. Thus, as most people know, “oil andwater don’t mix.”

The immiscible quality of oil and water inrelation to one another explains a number ofphenomena from the everyday world. It is easyenough to wash mud from your hands underrunning water, because the water and the mud(which, of course, contains water) are highly mis-cible. But motor oil or oil-based paint will leave afilm on the hands, as these substances bond tooily molecules in the skin itself, and no amountof water will wash the film off. To remove it, oneneeds mineral spirits, paint thinner, or soapmanufactured to bond with the oils.

Everyone has had the experience of eatingspicy foods and then trying to cool down themouth with cold water—and just about everyonehas discovered that this does not work. This isbecause most spicy substances contain emulsi-fied oils that coat the tongue, and these oils arenot attracted to water. A much better “solution”(in more ways than one) is milk, and in particu-lar, whole milk. Though a great deal of milk iscomposed of water, it also contains tiny dropletsof fats and proteins that join with the oils in spicysubstances, thus cooling down the mouth.

Crossing the Oil/Water Barrier

The active ingredient in a dry-cleaning chemicalusually has nonpolar molecules, because thetoughest stains are made by oils and greases. But

what about ordinary soap, which we use in waterto wash both water- and oil-based substancesfrom our bodies? Though, as noted earlier, cer-tain kinds of oil stains require special cleaningsolutions, ordinary soap is fine for washing offthe natural oils secreted by the skin. Likewise, itcan wash off small concentrations of oil thatcome from the environment and attach to theskin—for instance, from working over a deepfryer in a fast-food restaurant.

How does the soap manage to “connect”both with oils and water? Likewise, how doesmilk—which is water-soluble and capable ofdilution in water—wash away the oils associatedwith spicy food? To answer these questions, weneed to briefly consider a subject examined inmore depth within the Mixtures essay: emul-sions, or mixtures of two immiscible liquids.

EMULSIFIERS AND SURFAC-

TANTS. The dispersion of two substances inan emulsion is achieved through the use of anemulsifier or surfactant. Made up of moleculesthat are both water- and oil-soluble, an emulsifi-er or surfactant acts as an agent for joining othersubstances in an emulsion. The two words arevirtually synonymous, but “emulsifier” is usedtypically in reference to foods, whereas “surfac-tant” most often refers to an ingredient in deter-gents and related products.

In an emulsion, millions of surfactants sur-round the dispersed droplets of solute, known asthe internal phase, shielding them from the sol-vent, or external phase. Surfactants themselvesare often used in laundry or dish detergent,because most stains on plates or clothes are oil-based, whereas the detergent itself is applied tothe clothes in a water-based external phase. Theemulsifiers in milk help to bond oily particles ofmilk fat (cream) and protein to the externalphase of water that comprises the majority ofmilk’s volume.

As for soap, it is a mixture of an acid and abase—specifically, carboxylic acids joined with abase such as sodium hydroxide. Carboxylic acids,chains of hydrogen and carbon atoms, are justsome of many hydrocarbons that form the chem-ical backbone of a vast array of organic sub-stances. Because it is a salt, meaning that it isformed from the reaction of an acid with a base,soap partially separates into its component ionsin water. The active ion is RCOO- (R stands for



Solutionsthe hydrocarbon chain), whose two ends behavein different fashions, making it a surfactant.

The hydrocarbon end (R-) is said to belipophilic, or “oil-loving”; on the other hand, thecarboxylate end (-COO-) is hydrophilic, or“water-loving.” As a result, soap can dissolve inwater, but can also clean greasy stains. When soapis mixed with water, it does not form a true solu-tion, due to the presence of the hydrocarbons,which attract one another to form sphericalaggregates called micelles. The lipophilic “tails”of the hydrocarbons are turned toward the inte-rior of the micelle, while the hydrophilic “heads”remain facing toward the water that forms theexternal phase.

ETHANOL. When a hydrocarbon joinswith other substances in one of the hydrocarbonfunctional groups, these form numerous hydro-carbon derivatives. Among the hydrocarbonderivatives is alcohol, and within this grouping isethyl alcohol or ethanol, which includes a hydro-carbon bonded to the –OH functional group.The formula for ethanol is C2H6O, though this issometimes rendered as C2H5OH to show that thealcohol functional group (–OH) is bonded to thehydroxyl radical (C2H5).

Ethanol, the same alcohol found in alcoholicbeverages, is obviously miscible with water. Beer,after all, is mostly water, and if ethanol and waterwere not miscible, it would be impossible to mixalcohol with water or a water-based substance tomake Scotch and water or a Bloody Mary (vodkaand tomato juice.) In fact, water and ethanol arecompletely miscible, whether the solution be99% water and 1% ethanol, 50% of both, or 99%ethanol and 1% water.

How can this be, given the fact that ethanolis a hydrocarbon derivative—in other words, acousin of petroleum? The addition of the term“derivative” gives us a clue: note that ethanolcontains oxygen, just as water does. As in water,the oxygen and hydrogen form a polar bond, giv-ing that portion of the ethanol molecule a highaffinity for water. Molecules of water and thealcohol functional group are joined through anintermolecular force called hydrogen bonding.

Aqueous Solutions

The term aqueous refers to water, and becausewater has extraordinary solvent qualities (dis-cussed in the Osmosis essay) it serves as themedium for numerous solutions. Furthermore,

reactions in aqueous solutions—though we canonly touch on them briefly here—constitute alarge and significant body of reactions studied bychemists.

Most of the solutions we have discussed upto this point are aqueous. We have referred, atleast in an external or phenomenological way, tothe means by which sugar dissolves in an aque-ous solution, such as tea. Now let us consider,from a molecular or structural standpoint, themeans by which salt does so as well.

SALTWATER. Salt is formed of posi-tively charged sodium ions, firmly bonded tonegatively charged chlorine ions. To break theseionic bonds and dissolve the sodium chloride,water must exert a strong attraction. However,salt is not really composed of molecules but sim-ply repeating series of sodium and chlorineatoms joined together in a simple face-centeredcubic lattice which has each each sodium ion(Na+) surrounded by six chlorine ions (Cl–); atthe same time, each chlorine ion is surroundedby six sodium ions.

In dissolving the salt, water molecules sur-round ions that make up the salt, holding themin place through the electrical attraction ofopposite charges. Due to the differences in electronegativity that give the water molecule itshigh polarity, the hydrogen atoms are more positively charged, and thus these attract the negatively charged chlorine ion. Similarly,the negatively charged oxygen end of the watermolecule attracts the positively charged sodiumion. As a result, the salt is “surrounded” or dissolved.

PLASMA AND AQUEOUS-SO-

LUTION REACTIONS IN THE

BODY. Whereas saltwater is an aqueous solu-tion that can eventually kill someone who drinksit, plasma is an essential component of the lifeprocess—and in fact, it contains a small quantityof sodium chloride. Not to be confused with thephase of matter also called plasma, this plasma isthe liquid portion of blood. Blood itself is about55% plasma, with red and white blood cells andplatelets suspended in it.

Plasma is in turn approximately 90% water,with a variety of other substances suspended ordissolved in it. Prominent among these sub-stances are proteins, of which there are about 60in plasma, and these serve numerous importantfunctions—for instance, as antibodies. Plasma



Solutions


also contains ions, which prevent red blood cellsfrom taking up excess water in osmosis. Promi-nent among these ions are those of salt, or sodi-um chloride. Plasma also transports nutrientssuch as amino acids, glucose, and fatty acids, aswell as waste products such as urea and uric acid,which it passes on to the kidneys for excretion.

Much of the activity that sustains life in thebody of a living organism can be characterized asa chemical reaction in an aqueous solution.When we breathe in oxygen, it is taken to thelungs and fed into the bloodstream, where it

associates with the iron-containing hemoglobin

in red blood cells and is transported to the

organs. In the stomach, various aqueous-solu-

tion reactions process food, turning part of it

into fuel that the blood carries to the cells, where

the oxygen engages in complex reactions with

nutrients.

A Miscellany of Solutions

Aqueous-solution reactions can lead to the for-

mation of a solid, as when a solution of potassi-

AQUEOUS: A term describing a solu-

tion in which a substance or substances are

dissolved in water.






changed chemically.

CONCENTRATED: A qualitative term

describing a solution with a relatively large

ratio of solute to solvent.

DILUTE: A qualitative term describing

a solution with a relatively small ratio of

solute to solvent.

EMULSIFIER: A surfactant. (Usually

the term “emulsifier” is applied to foods,

and “surfactant” to detergents and other

inedible products.)

EMULSION: A mixture of two immis-

cible liquids, such as oil and water, made by

dispersing microscopic droplets of one liq-

uid in another. In creating an emulsion,

surfactants act as bridges between the two

liquids.




ing different properties. One example of

this is sand in a container of water. Rather

than dissolving to form a homogeneous

mixture as sugar does, the sand sinks to the

bottom.





mixture.

IMMISCIBLE: Possessing a negligible

value of miscibility.

MASS PERCENT: A quantitative

means of measuring solubility in terms of

the mass of the solute divided by the mass

of the solution, multiplied by 100%.

MISCIBILITY: A term identifying the

relative ability of two substances to dissolve

in one another. Miscibility is qualitative,

like “fast” or “slow”; solubility, on the other

hand, can be a quantitative term.


composition, composed of molecules or

atoms of differing types. Compare with

pure substance.

KEY TERMS


Solutions


um chromate (K2CrO4) is added to an aqueoussolution of barium nitrate (Ba[NO3]2 to formsolid barium chromate (BaCrO4) and a solutionof potassium nitrate (KNO3). This reaction isdescribed in the essay on Chemical Reactions.

We have primarily discussed liquid solu-tions, and in particular aqueous solutions. Itshould be stressed, however, that solutions canalso exist in the gaseous or solid phases. The airwe breathe is a solution, not a compound: inother words, there is no such thing as an “airmolecule.” Instead, it is made up of diatomic ele-

ments (those in which two atoms join to form amolecule of a single element); monatomic ele-ments (those elements that exist as single atoms);one element in a triatomic molecule; and twocompounds.

The “solvent” in air is nitrogen, a diatomicelement that accounts for 78% of Earth’s atmos-phere. Oxygen, also diatomic, constitutes anadditional 21%. Argon, which like all noble gasesis monatomic, ranks a distant third, with 0.93%.The remaining 0.07% is made up of traces ofother noble gases; the two compounds men-

MOLARITY: A quantitative means of

showing the concentration of solute to

solution. This is measured in moles of

solute per liter of solution, abbreviated

mol/L.

PURE SUBSTANCE: A substance—

either an element or compound—that has

an unvarying composition. This means

that by changing the proportions of atoms,

the result would be an entirely different

substance. Compare with mixture.

QUALITATIVE: Involving a compari-

son between qualities that are not precisely

defined, such as “fast” and “slow” or

“warm” and “cold.”

QUANTITATIVE: Involving a compari-

son between precise quantities—for

instance, 10 lb vs. 100 lb, or 50 MPH vs. 120

MPH.

SATURATED: A qualitative term

describing a solution that contains as

much solute as it can dissolve at a given

temperature.

SOLUBILITY: In a broad, qualitative

sense, solubility refers to the property of

being soluble. Chemists, however, usually

apply the word in a quantitative sense, to

indicate the maximum amount of a sub-

stance that dissolves in a given amount of

solvent at a specific temperature. Solubility

is usually expressed in grams of solute per

100 g of solvent.

SOLUBLE: Capable of dissolving in a

solvent.

SOLUTE: The substance or substances

that are dissolved in a solvent to form a

solution.



solute) is dissolved in a solvent—for exam-


SOLVENT: The substance or sub-

stances that dissolve a solute to form a

solution.

SURFACTANT: A substance made up

of molecules that are both water- and oil-

soluble, which acts as an agent for joining

other substances in an emulsion.


solution that is capable of dissolving

additional solute, if that solute is intro-

duced to it.



Solutions tioned, carbon dioxide and water (in vaporform); and, high in the atmosphere, the triatom-ic form of oxygen known as ozone (O3).

The most significant solid solutions arealloys of metals, discussed in the essay on Mix-tures, as well as in essays on various metal fami-lies, particularly the Transition Metals. Somewell-known alloys include bronze (three-quar-ters copper, one-quarter tin); brass (two-thirdscopper, one-third zinc); pewter (a mixture of tinand copper with traces of antimony); andnumerous alloys of iron—particularly steel—aswell as alloys involving other metals.

WHERE TO LEARN MORE

“Aqueous Solutions” (Web site). <http://www.tannerm.com/aqueous.htm> (June 6, 2001).

Gibson, Gary. Making Things Change. Brookfield, CT:Copper Beech Books, 1995.


“Introduction to Solutions” (Web site). <http://edie.cprost.sfu.ca/~rhlogan/solintro.html> (June 6, 2001).

Knapp, Brian J. Air and Water Chemistry. Illustrated byDavid Woodroffe. Danbury, CT: Grolier Educational,1998.

Seller, Mick. Elements, Mixtures, and Reactions. NewYork: Shooting Star Press, 1995.

“Solubility.” Spark Notes (Web site). <http://www.sparknotes.com/chemistry/solutions/solubility/terms.html> (June 6, 2001).

“Solutions and Solubility” (Web site). <http://educ.queensu.ca/~science/main/concept/chem/c10/c101amo1.htm> (June 6, 2001).

Watson, Philip. Liquid Magic. Illustrated by ElizabethWood and Ronald Fenton. New York: Lothrop, Lee,and Shepard Books, 1982.





OSMOS ISOsmosis

CONCEPTThe term osmosis describes the movement of asolvent through a semipermeable membranefrom a less concentrated solution to a more con-centrated one. Water is sometimes called “theperfect solvent,” and living tissue (for example, ahuman being’s cell walls) is the best example of asemipermeable membrane. Osmosis has a num-ber of life-preserving functions: it assists plantsin receiving water, it helps in the preservation offruit and meat, and is even used in kidney dialy-sis. In addition, osmosis can be reversed toremove salt and other impurities from water.

HOW IT WORKSIf you were to insert a hollow tube of a certaindiameter into a beaker of water, the water wouldrise inside the tube and reach the same level asthe water outside it. But suppose you sealed thebottom end of the tube with a semipermeablemembrane, then half-filled the tube with saltwater and again inserted it into the beaker. Overa period of time, the relative levels of the saltwater in the tube and the regular water in thebeaker would change, with the fresh water grad-ually rising into the beaker.

This is osmosis at work; however, beforeinvestigating the process, it is necessary to under-stand at least three terms. A solvent is a liquidcapable of dissolving or dispersing one or moreother substances. A solute is the substance that isdissolved, and a solution is the resulting mixtureof solvent and solute. Hence, when you mix apacket of sugar into a cup of hot coffee, the cof-fee—which is mostly water—acts as a solvent forthe sugar, a solute, and the resulting sweetenedcoffee is a solution. (Indeed, people who need a

cup of coffee in the morning might say that it isa “solution” in more ways than one!) The relativeamount of solute in the solution determineswhether it can be described as more or less con-centrated.

Water and Oil: MolecularDifferences

In the illustrations involving the beaker and thehollow tube, water plays one of its leading roles,as a solvent. It is possible to use a number ofother solvents for osmosis, but most of the onesthat will be discussed here are water-based sub-stances. In fact, virtually everything people drinkis either made with water as its central compo-nent (soft drinks, coffee, tea, beer and spirits), orcomes from a water-based plant or animal lifeform (fruit juices, wine, milk.) Then of coursethere is water itself, still the world’s most populardrink.

By contrast, people are likely to drink an oilyproduct only in extreme circumstances: forinstance, to relieve constipation, holistic-healthpractitioners often recommend a mixture ofolive oil and other compounds for this purpose.Oil, unlike water, has a tendency to pass straightthrough a person’s system, without largeamounts of it being absorbed through osmosis.In fact, oil and water differ significantly at themolecular level.

Water is the best example of a polar mole-cule, sometimes called a dipole. As everyoneknows, water is a name for the chemical H2O, inwhich two relatively small hydrogen atoms bondwith a large atom of oxygen. You can visualize awater molecule by imagining oxygen as a basket-ball with hydrogen as two baseballs fused to the


Osmosis


basketball’s surface. Bonded together as they are,the oxygen tends to pull electrons from thehydrogen atoms, giving it a slight negative chargeand the hydrogen a slight positive charge.

As a result, one end of a water molecule hasa positive electrical charge, and the other end anegative charge. This in turn causes the positiveend of one molecule to attract the negative sideof its neighbor, and vice versa. Though the elec-tromagnetic force is weak, even in relative terms,it is enough to bond water molecules tightly toone another.

By contrast, oily substances—whether theoil is animal-, vegetable-, or petroleum-based—are typically nonpolar, meaning that the positiveand negative charges are distributed evenly

across the surface of the molecule. Hence, thebond between oil molecules is much less tightthan for water molecules. Clean motor oil in acar’s crankshaft behaves as though it were madeof millions of tiny ball-bearings, each rollingthrough the engine without sticking. Water, onthe other hand, has a tendency to stick to sur-faces, since its molecules are so tightly bonded toone another.

This tight bond gives water highly unusualproperties compared to other substances close toits molecular weight. Among these are its highboiling point, its surprisingly low density whenfrozen, and the characteristics that make osmosispossible. Thanks to its intermolecular structure,water is not only an ideal solvent, but its closely

A SEMIPERMEABLE MEMBRANE PERMITS ONLY CERTAIN COMPONENTS OF SOLUTIONS, USUALLY THE SOLVENT, TO PASS

THROUGH.


Osmosis


packed structure enables easy movement, as, forinstance, from an area of low concentration to anarea of high concentration.

In the beaker illustration, the “pure” water isalmost pure solvent. (Actually, because of its sol-vent qualities, water seldom appears in a purestate unless one distills it: even water flowingthrough a “pure” mountain stream carries allsorts of impurities, including microscopic parti-cles of the rocks over which it flows.) In any case,the fact that the water in the beaker is almostpure makes it easy for it to flow through thesemipermeable membrane in the bottom of thetube. By contrast, the solute particles in the salt-water solution have a much harder time passingthrough, and are much more likely to block theopenings in the membrane. As a result, the move-ment is all in one direction: water in the beakermoves through the membrane, and into the tube.

A few points of clarification are in orderhere. A semipermeable membrane is anythingwith a structure somewhere between that of, say,plastic on the one hand and cotton on the other.Were the tube in the beaker covered with Saranwrap, for instance, no water would pass through.On the other hand, if one used a piece of cottonin the bottom of the tube, the water would passstraight through without osmosis taking place. Incontrast to cotton, Gore-tex is a fabric containing

a very thin layer of plastic with billions of tinypores which let water vapor flow out withoutallowing liquid water to seep in. This accountsfor the popularity of Gore-tex for outdoorgear—it keeps a person dry without holding intheir sweat. So Gore-tex would work well as asemipermeable membrane.

Also, it is important to consider the possibil-ities of what can happen during the process ofosmosis. If the tube were filled with pure salt, orsalt with only a little water in it, osmosis wouldreach a point and then stop due to osmotic pres-sure within the substance. Osmotic pressureresults when a relatively concentrated substancetakes in a solvent, thus increasing its pressureuntil it reaches a point at which the solution willnot allow any more solvent to enter.


Cell Behavior and Salt Water

Cells in the human body and in the bodies of allliving things behave like microscopic bags ofsolution housed in a semipermeable membrane.The health and indeed the very survival of a per-son, animal, or plant depends on the ability of

OSMOSIS EQUALIZES CONCENTRATION.


Osmosis


the cells to maintain their concentration ofsolutes.

Two illustrations involving salt waterdemonstrate how osmosis can produce disas-trous effects in living things. If you put a carrot insalty water, the salt water will “draw” the waterfrom inside the carrot—which, like the humanbody and most other forms of life, is mostlymade up of water. Within a few hours, the carrotwill be limp, its cells shriveled.

Worse still is the process that occurs when aperson drinks salt water. The body can handle alittle bit, but if you were to consume nothing butsalt water for a period of a few days, as in the caseof being stranded on the proverbial desert island,the osmotic pressure would begin drawing waterfrom other parts of your body. Since a humanbody ranges from 60% water (in an adult male)to 85% in a baby, there would be a great deal ofwater available—but just as clearly, water is theessential ingredient in the human body. If youcontinued to ingest salt water, you would eventu-ally experience dehydration and die.

How, then, do fish and other forms ofmarine life survive in a salt-water environment?In most cases, a creature whose natural habitat isthe ocean has a much higher solute concentra-tion in its cells than does a land animal. Hence,

for them, salt water is an isotonic solution, or onethat has the same concentration of solute—andhence the same osmotic pressure—as in theirown cells.

Osmosis in Plants

Plants depend on osmosis to move water fromtheir roots to their leaves. The further toward theedge or the top of the plant, the greater the soluteconcentration, which creates a difference inosmotic pressure. This is known as osmoticpotential, which draws water upward. In addi-tion, osmosis protects leaves against losing waterthrough evaporation.

Crucial to the operation of osmosis in plantsare “guard cells,” specialized cells dispersed alongthe surface of the leaves. Each pair of guard cellssurrounds a stoma, or pore, controlling its abili-ty to open and thus release moisture.

In some situations, external stimuli such assunlight may cause the guard cells to draw inpotassium from other cells. This leads to anincrease in osmotic potential: the guard cellbecomes like a person who has eaten a dry bis-cuit, and is now desperate for a drink of water towash it down. As a result of its increased osmot-ic potential, the guard cell eventually takes on

IN THIS 1966 PHOTO, A GIRL IS UNDERGOING KIDNEY DIALYSIS, A CRUCIAL MODERN MEDICAL APPLICATION THAT

RELIES ON OSMOSIS. (Bettmann/Corbis. Reproduced by permission.)


Osmosiswater through osmosis. The guard cells thenswell with water, opening the stomata andincreasing the rate of gas exchange throughthem. The outcome of this action is an increasein the rate of photosynthesis and plant growth.

When there is a water shortage, however,other cells transmit signals to the guard cells thatcause them to release their potassium. Thisdecreases their osmotic potential, and waterpasses out of the guard cells to the thirsty cellsaround them. At the same time, the resultantshrinkage in the guard cells closes the stomata,decreasing the rate at which water transpiresthrough them and preventing the plant fromwilting.

Osmosis and Medicine

Osmosis has several implications where medicalcare is concerned, particularly in the case of thestorage of vitally important red blood cells. Theseare normally kept in a plasma solution which isisotonic to the cells when it contains specific pro-portions of salts and proteins. However, if redblood cells are placed in a hypotonic solution, orone with a lower solute concentration than in thecells themselves, this can be highly detrimental.

Hence water, a life-giving and life-preservingsubstance in most cases, is a killer in this context.If red blood cells were stored in pure water,osmosis would draw the water into the cells,causing them to swell and eventually burst. Sim-ilarly, if the cells were placed in a solution with ahigher solute concentration, or hypertonic solu-tion, osmosis would draw water out of the cellsuntil they shriveled.

In fact, the plasma solution used by mosthospitals for storing red blood cells is slightlyhypertonic relative to the cells, to prevent themfrom drawing in water and bursting. Physiciansuse a similar solution when injecting a drugintravenously into a patient. The active ingredi-ent of the drug has to be suspended in some kindof medium, but water would be detrimental forthe reasons discussed above, so instead the doc-tor uses a saline solution that is slightly hyper-tonic to the patient’s red blood cells.

One vital process closely linked to osmosis isdialysis, which is critical to the survival of manyvictims of kidney diseases. Dialysis is the processby which an artificial kidney machine removeswaste products from a patients’ blood—per-forming the role of a healthy, normally function-

ing kidney. The openings in the dialyzing mem-brane are such that not only water, but salts andother waste dissolved in the blood, pass throughto a surrounding tank of distilled water. The redblood cells, on the other hand, are too large toenter the dialyzing membrane, so they return tothe patient’s body.

Preserving Fruits and Meats

Osmosis is also used for preserving fruits andmeats, though the process is quite different forthe two. In the case of fruit, osmosis is used todehydrate it, whereas in the preservation of meat,osmosis draws salt into it, thus preventing theintrusion of bacteria.

Most fruits are about 75% water, and thismakes them highly susceptible to spoilage. Topreserve fruit, it must be dehydrated, which—asin the case of the salt in the meat— presents bac-teria with a less-than-hospitable environment.Over the years, people have tried a variety ofmethods for drying fruit, but most of these havea tendency to shrink and harden the fruit. Thereason for this is that most drying methods, suchas heat from the Sun, are relatively quick anddrastic; osmosis, on the other hand, is slower,more moderate—and closer to the behavior ofnature.

Osmotic dehydration techniques, in fact,result in fruit that can be stored longer than fruitdehydrated by other methods. This in turn makesit possible to provide consumers with a widervariety of fruit throughout the year. Also, thefruit itself tends to maintain more of its flavorand nutritional qualities while keeping outmicroorganisms.

Because osmosis alone can only removeabout 50% of the water in most ripe fruits, how-ever, the dehydration process involves a second-ary method as well. First the fruit is blanched, orplaced briefly in scalding water to stop enzymat-ic action. Next it is subjected to osmotic dehy-dration by dipping it in, or spreading it with, aspecially made variety of syrup whose sugardraws out the water in the fruit. After this, airdrying or vacuum drying completes the process.The resulting product is ready to eat; can be pre-served on a shelf under most climatic conditions;and may even be powdered for making confec-tionery items.

Whereas osmotic dehydration of fruit is cur-rently used in many parts of the world, the salt-



Osmosis

curing of meat in brine is largely a thing of thepast, due to the introduction of refrigeration.Many poorer families, even in the industrializedworld, however, remained without electricitylong after it spread throughout most of Europeand North America. John Steinbeck’s Grapes ofWrath (1939) offers a memorable scene in whicha contemporary family, the Joads, kill and cure apig before leaving Oklahoma for California. Anda Web site for Walton Feed, an Idaho companyspecializing in dehydrated foods, offers reminis-cences by Canadians whose families were stillsalt-curing meats in the middle of the twentiethcentury. Verla Cress of southern Alberta, forinstance, offered a recipe from which the follow-ing details are drawn.


First a barrel is filled with a solution con-taining 2 gal (7.57 l) of hot water and 8 oz (.2268kg) of salt, or 32 parts hot water to one part salt,as well as a small quantity of vinegar. The pig orcow, which would have just been slaughtered,should then be cut up into what Cress called“ham-sized pieces (about 10-15 lb [5-7 kg])each.” The pieces are then soaked in the brinebarrel for six days, after which the meat isremoved, dried, “and put... in flour or gunnysacks to keep the flies away. Then hang it up in acool dry place to dry. It will keep like this for per-haps six weeks if stored in a cool place during theSummer. Of course, it will keep much longer inthe Winter. If it goes bad, you’ll know it!”

Cress offered another method, one still usedon ham today. Instead of salt, sugar is used in amixture of 32 oz (.94 l) to 3 gal (11.36 l) of water.After being removed, the meat is smoked—thatis, exposed to smoke from a typically aromaticwood such as hickory, in an enclosed barn—forthree days. Smoking the meat tends to make itlast much longer: four months in the summer,according to Cress.

The Walton Feeds Web page included anoth-er brine-curing recipe, this one used by thewomen of the Stirling, Alberta, Church of theLatter-Day Saints in 1973. Also included werereminiscences by Glenn Adamson (born 1915):“...When we butchered a pig, Dad filled a wood-en 45-gal (170.34 l) barrel with salt brine. We cutup the pig into maybe eight pieces and put it inthe brine barrel. The pork soaked in the barrelfor several days, then the meat was taken out, andthe water was thrown away.... In the hot summerdays after they [the pieces of meat] had dried,they were put in the root cellar to keep them cool.The meat was good for eating two or threemonths this way.”

For thousands of years, people used salt tocure and preserve meat: for instance, the sailingships that first came to the New World carried onboard barrels full of cured meat, which fed sailorson the voyage over. Meat was not the only type offood preserved through the use of salt or brine,which is hypertonic—and thus lethal—to bacte-ria cells. Among other items packed in brine werefish, olives, and vegetables.

Even today, some types of canned fish cometo the consumer still packed in brine, as do olives.Another method that survives is the use of

HYPERTONIC: Of higher osmotic

pressure.

HYPOTONIC: Of lower osmotic

pressure.

ISOTONIC: Of equal osmotic pressure.

OSMOSIS: The movement of a solvent

through a semipermeable membrane from

a less concentrated solution to a more con-

centrated one.

OSMOTIC POTENTIAL: A difference

in osmotic pressure that draws water from

an area of less osmotic pressure to an area

of greater osmotic pressure.

OSMOTIC PRESSURE: The pressure

that builds in a substance as it experiences

osmosis, and eventually stops that process.

SOLVENT: A liquid capable of dis-

solving or dispersing one or more other

substances.

SOLUTE: A substance capable of being

dissolved in a solvent.

SOLUTION: A mixture of solvent and

solute.

KEY TERMS


Osmosissugar—which can be just as effective as salt forkeeping out bacteria—to preserve fruit in jam.

Reverse Osmosis

Given the many ways osmosis is used for pre-serving food, not to mention its many interac-tions with water, it should not be surprising todiscover that osmosis can also be used for desali-nation, or turning salt water into drinking water.Actually, it is not osmosis, strictly speaking, butrather reverse osmosis that turns salt water fromthe ocean—97% of Earth’s water supply—intowater that can be used for bathing, agriculture,and in some cases even drinking.

When you mix a teaspoon of sugar into acup of coffee, as mentioned in an earlier illustra-tion, this is a non-reversible process. Short ofsome highly complicated undertaking—forinstance, using ultrasonic sound waves—it would be impossible to separate solute and solvent.

Osmosis, on the other hand, can be reversed.This is done by using a controlled external pres-sure of approximately 60 atmospheres, an atmos-phere being equal to the air pressure at sealevel—14.7 pounds-per-square-inch (1.013 �

105 Pa.) In reverse osmosis, this pressure isapplied to the area of higher solute concentra-tion—in this case, the seawater. As a result, thepressure in the seawater pushes water moleculesinto a reservoir of pure water.

If performed by someone with a few rudi-mentary tools and a knowledge of how to pro-vide just the right amount of pressure, it is possi-ble that reverse osmosis could save the life of ashipwreck victim stranded in a location withouta fresh water supply. On the other hand, a personin such a situation may be able to absorb suffi-cient water from fruits and plant life, as TomHanks’s character did in the 2001 film Cast Away.

Companies such as Reverse Osmosis Sys-tems in Atlanta, Georgia, offer a small unit forhome or business use, which actually performsthe reverse-osmosis process on a small scale. Theunit makes use of a process called crossflow,which continually cleans the semipermeablemembrane of impurities that have been removedfrom the water. A small pump provides the pres-sure necessary to push the water through themembrane. In addition to an under-the-sinkmodel, a reverse osmosis water cooler is alsoavailable.

Not only is reverse osmosis used for makingwater safe, it is also applied to metals in a varietyof capacities, not least of which is its use in treat-ing wastewater from electroplating. But there areother metallurgical methods of reverse osmosisthat have little to do with water treatment: metalfinishing, as well as recycling of metals andchemicals. These processes are highly complicat-ed, but they involve the same principle of remov-ing impurities that governs reverse osmosis.

WHERE TO LEARN MORE

Francis, Frederick J., editor-in-chief. Encyclopedia of FoodScience and Technology. New York: Wiley, 2000.

Gardner, Robert. Science Project Ideas About KitchenChemistry. Berkeley, N.J.: Enslow Publishers, 2002.

Laschish, Uri. “Osmosis, Reverse Osmosis, and OsmoticPressure: What They Are” (Web site). <http://members.tripod.com/~urila/> (February 20, 2001).

“Lesson 5: Osmosis” (Web site). <http://www.biologylessons.sdsu.edu/classes/lab5/semnet/> (Feb-ruary 20, 2001).

Rosenfeld, Sam. Science Experiments with Water. Illus-trated by John J. Floherty, Jr. Irvington-on-Hudson,NY: Harvey House, 1965.

“Salt-Curing Meat in Brine.” Walton Feed (Web site).<http://waltonfeed.com/old/brine.html> (February20, 2001).




D I S T I L L AT ION AND F I LTR AT IONDistillation and Filtration

CONCEPTWhen most people think of chemistry, they thinkabout joining substances together. Certainly, thebonding of elements to form compoundsthrough chemical reactions is an integral compo-nent of the chemist’s study; but chemists are alsoconcerned with the separation of substances.Some forms of separation, in which compoundsare returned to their elemental form, or in whichatoms split off from molecules to yield a com-pound and a separated element, are complexphenomena that require chemical reactions. Butchemists also use simpler physical methods,distillation and filtration, to separate mixtures.Distillation can be used to purify water, makealcohols, or separate petroleum into variouscomponents that range from natural gas to gaso-line to tar. Filtration also makes possible the sep-aration of gases or liquids from solids, andsewage treatment—a form of filtration—sepa-rates liquids, solids, and gases through a series ofchemical and physical processes.

HOW IT WORKS

Mixtures: Homogeneous andHeterogeneous

In contrast to a pure substance, a class thatincludes only elements and compounds, mix-tures constitute a much broader category. Usual-ly, a mixture consists of many compounds mixedtogether, but it is possible to have a mixture (airis an example) in which elemental substances arecombined with compounds. Elements alone canbe combined in a mixture, as when copper andzinc are alloyed to make brass. It is also possibleto have a mixture of mixtures, as for example

when milk (a particular type of mixture called anemulsion) is added to coffee.

It is sometimes difficult, without studyingthe chemical aspects involved, to recognize thedifference between a mixture and a pure sub-stance. A mixture, however, can be defined as asubstance with a variable composition, meaningthat if more of one component is added, it doesnot change the essential character of the mixture.People unconsciously recognize this when theyjoke, “Would you like a little coffee with yourmilk?”

This comment, made when someone puts agreat quantity of milk into a cup of coffee,implies that we generally regard coffee and milkas a solution in which the coffee dissolves themilk. But even if someone made a cup of coffeethat was half milk, so that the resulting substancehad a vanilla-like color, we would still call it cof-fee. If the components of a pure substance arealtered, however, it becomes something entirelydifferent.

Two hydrogen atoms bonded to an oxygenatom create water, but when two hydrogen atomsbond to two oxygen atoms, this is hydrogen per-oxide, an altogether different substance. Water atordinary temperatures does not bubble, whereashydrogen peroxide does. Hydrogen peroxideboils at a temperature more than 1.5 times theboiling point of water, and potatoes boiled inhydrogen peroxide would be nothing like pota-toes boiled in water. Eating them, in fact, couldvery well be fatal.

HOMOGENEOUS VS. HETERO-

GENEOUS. Unlike a compound or element,a mixture can never be broken down to a singletype of molecule or atom, yet there are mixtures


Distillationand

Filtration

that appear to be the same throughout. These arecalled homogeneous mixtures, in which theproperties and characteristics are the samethroughout the mixture. In a heterogeneous mix-ture, on the other hand, the composition is notthe same throughout, and in fact the mixture isseparated into regions with varying properties.

Chemists of the past sometimes had difficul-ty distinguishing between homogeneous mix-tures—virtually all of which can be described assolutions—and pure substances. Air, for exam-ple, appears to be a pure substance; yet there is nosuch thing as an “air molecule.” Instead, air iscomposed primarily of nitrogen and oxygen,which appear in molecular rather than atomicform, along with separated atoms of noble gasesand two important compounds: carbon dioxideand water. (Water exists in air as a vapor.)

There is less probability of confusing a het-erogeneous mixture with a pure substance,because one of the defining aspects of a hetero-geneous substance is the fact that it is clearly amixture. When sand is added to a container ofwater, and the sand sinks to the bottom, it isobvious that there is no “sand-and-water” mole-cule pervading the entire mixture. Clearly, theupper portion of the container is mostly water,and the lower portion mostly sand.

Separating Mixtures

There are two basic processes for separating mix-tures, distillation and filtration. In general, theseare applied for the separation of homogeneousand heterogeneous mixtures, respectively. Distil-lation is the use of heat to separate the compo-nents of a liquid and/or gas, while filtration is theseparation of solids from a fluid (either a gas or aliquid) by allowing the fluid to pass through a filter.

In a solution such as salt water, there are twocomponents: the solvent (the water) and thesolute (the salt). These can be separated by distil-lation in a laboratory using a burner placedunder a beaker containing the salt water. As thewater is heated, it passes out of the beaker in theform of steam, and travels through a tube cooledby a continual flow of cold water. Inside the tube,the steam condenses to form liquid water, whichpasses into a second beaker. Eventually, all of thesolvent will be distilled from the first beaker,leaving behind the salt that constituted the soluteof the original solution.

Distillation may seem like a chemicalprocess, but in fact it is purely physical, becausethe composition of neither compound—salt orwater—has been changed. As for filtration, it isclearly a physical process, just as heterogeneousmixtures are more obviously mixtures and notcompounds. Suppose we wanted to separate theheterogeneous mixture we described earlier, ofsand in water: all we would need would be amesh screen through which the liquid could bestrained, leaving behind the sand.

Despite the apparent simplicity of filtration,it can become rather complicated, as we shall see,in the treatment of sewage to turn it into waterthat poses no threat to the environment. Further-more, it should be noted that sometimes theseprocesses—distillation and filtration—are usedin tandem. Suppose we had a mixture of sandand salt water taken from the beach, and wewanted to separate the mixture. The first stepwould involve the simpler process of filtration,separating the sand from the salt water. Thiswould be followed by the separation of purewater and salt through the distillation process.


Distillation

In distillation, the more volatile component ofthe mixture—that is, the part that is more easilyvaporized—is separated from the less volatileportion. With regard to the illustration usedabove, of separating water from salt, clearly thewater is the more volatile portion: its boilingpoint is much, much lower than that of salt, andthe heat required to vaporize the salt is so greatthat the water would be long since vaporized bythe time the salt was affected.

Once the volatile portion has been vapor-ized, or turned into distillate, it experiences con-densation—the cooling of a gas to form a liquid,after which the liquid or condensate is collected.Note that in all these stages of distillation, heat isthe agent of separation. This is especially impor-tant in the process of fractional distillation, dis-cussed below.

The process of removing water vapor fromsalt water, described earlier, is an example ofwhat is called “single-stage differential distilla-tion.” Sometimes, however—especially when a




high level of purity in the resulting condensate isdesired—it may be necessary to subject the con-densate to multiple processes of distillation toremove additional impurities. This process iscalled rectification.

PURIFIED WATER AND ETHA-

NOL. If distillation can be used to separatewater from salt, obviously it can also be used toseparate water from microbes and other impuri-ties through a more detailed process of rectifica-tion. Distilled water, purchased for drinking andother purposes, is just one of the more commonapplications of the distillation process. Anotherwell-known use of this process is the productionof ethyl alcohol, made famous (or perhaps infa-mous) by stories of bootleggers producing

homemade whiskey from “stills” or distilleries inthe woods.

Ethanol, the alcohol in alcoholic beverages,is produced by the fermentation of glucose, asugar found in various natural substances. Thechoice of substances is a function of the desiredproduct: grapes for wine; barley and hops forbeer; juniper berries for gin; potatoes for vodka,and so on. Glucose reacts with yeast, fermentingto yield ethanol, and this reaction is catalyzed byenzymes in the yeast. The reaction continuesuntil the alcohol content reaches a point equal toabout 13%.

After that point, the yeast enzymes can nolonger survive, and this is where the productionof beer or wine usually ends. Fortified wine or

A WATER FILTRATION PLANT IN CONTRA COSTA COUNTY, CALIFORNIA. (James A. Sugar/Corbis. Reproduced by permission.)

Distillationand Filtration



malt liquor—variations of wine and beer, respec-

tively, that are more than 13% alcohol, or 26

proof—are exceptions. These two products are

the result of mixtures between undistilled beer

and wine and products of distillation having a

higher alcohol content.

Distillation constitutes the second stage in

the production of alcohol. Often, the fermenta-

tion products are subjected to a multistage recti-

fication process, which yields a mixture of up to

95% ethanol and 5% water. This mixture is called

an azeotrope, meaning that it will not change

composition when distilled further. Usually the

production of whiskey or other liquor stops well

below the point of yielding an azeotrope, but in

some states, brands of grain alcohol at 190 proof(95%) are sold.

INDUSTRIAL DISTILLATION.

In contrast to ethanol, methanol or “wood alco-hol” is a highly toxic substance whose ingestioncan lead to blindness or death. It is widely used inindustry for applications in adhesives, plastics,and other products, and was once produced bythe distillation of wood. In this old productionmethod, wood was heated in the absence of air,such that it did not burn, but rather decomposedinto a number of chemicals, a fraction of whichwere methanol and other alcohols.

The distillation of wood alcohol was aban-doned for a number of reasons, not least ofwhich was the environmental concern: thou-

FEW INDUSTRIES EMPLOY DISTILLATION TO A GREATER DEGREE THAN DOES THE PETROLEUM INDUSTRY. SHOWN HERE

IS AN OIL REFINERY. (Mark L. Stephenson/Corbis. Reproduced by permission.)

Distillationand

Filtration




sands of trees had to be cut down for use in aninefficient process yielding only small portions ofthe desired product. Today, methanol is pro-duced from synthesis gas, itself a product of coalgasification; nonetheless, numerous industrialprocesses use distillation.

Distillation is employed, for instance, to sep-arate the hydrocarbons benzene and toluene, aswell as acetone from acetic acid. In nuclear powerplants that require quantities of the hydrogenisotope deuterium, the lighter protium isotope,or plain hydrogen, is separated from the heavierdeuterium by means of distillation. Few indus-tries, however, employ distillation to a greaterdegree than the petroleum industry.

Through a process known as fractional dis-tillation, oil companies separate hydrocarbonsfrom petroleum, allowing those of lower molec-ular mass to boil off and be collected first. Forinstance, at temperatures below 96.8°F (36°C),natural gas separates from petroleum. Other sub-stances, including petroleum ether and naphtha,separate before the petroleum reaches the156.2°F-165.2°F (69°C-74°C) range, at whichpoint gasoline separates.

Fractional distillation continues, with theseparation of other substances, all the way up to959°F (515°C), above which tar becomes the lastitem to be separated. Petroleum and petroleum-product companies account for a large portion ofthe more than 40,000 distillation units in opera-tion across the country. About 95% of all indus-trial separation processes involve distillation,which consumes large amounts of energy; forthat reason, ever more efficient forms of distilla-tion are continually being researched.

Filtration

In the simplest form of filtration, described ear-lier, a suspension (solid particles floating in a liq-uid) is allowed to pass through filter paper sup-ported in a glass funnel. The filter paper traps thesolid particles, allowing a clear solution (the fil-trate) to pass through the funnel and into areceiving container.

There are two basic purposes for filtration:either to capture the solid material suspended inthe fluid, or to clarify the fluid in which the solidis suspended. An example of the former occurs,for instance, when panning for gold: the water isallowed to pass through a sieve, leaving behindrocks that—the prospector hopes—contain gold.

An example of the latter occurs in sewage treat-ment. Here the solid left behind is feces—asundesirable as gold is desirable.

Filtration can further be classified as eithergaseous or liquid, depending on which of the flu-ids constitutes the filtrate. Furthermore, the forcethat moves the fluid through the filter—whethergravity, a vacuum, or pressure—is another defin-ing factor. The type of filter used can also serve todistinguish varieties of filtration.

LIQUID FILTRATION. In liquid fil-tration, such as that applied in sewage treatment,a liquid can be pulled through the filter by grav-itational force, as in the examples given. On theother hand, some sort of applied pressure, or apressure differential created by the existence of avacuum, can force the liquid through the filter.

Water filtration for purification purposes isoften performed by means of gravitation. Usual-ly, water is allowed to run down through a thicklayer of granular material such as sand, gravel, orcharcoal, which removes impurities. These layersmay be several feet thick, in which case the filteris known as a deep-bed filter. Water purificationplants may include fine particles of charcoal,known as activated carbon, in the deep-bed filterto absorb unpleasant-smelling gases.

For separating smaller volumes of solution,a positive-pressure system—one that uses exter-nal pressure to push the liquid through the fil-ter—may be used. Fluid may be introducedunder pressure at one end of a horizontal tank,then forced through a series of vertical platescovered with thick filtering cloths that collectsolids. In a vacuum filter, a vacuum (an area vir-tually devoid of matter, including air) is createdbeneath the solution to be filtered, and as a result,atmospheric pressure pushes the liquid throughthe filter.

When the liquid is virtually freed of impuri-ties, it may be passed through a second variety offilter, known as a clarifying filter. This type of fil-ter is constructed of extremely fine-mesh wiresor fibers designed to remove the smallest parti-cles suspended in the liquid. Clarifying filtersmay also use diatomaceous earth, a finely pow-dered material produced from the decay ofmarine organisms.

GAS FILTRATION. Gas filtration isused in a common appliance, the vacuum clean-er, which passes a stream of dust-filled air


Distillationand

Filtration

through a filtering bag inside the machine. The

bag traps solid particles, while allowing clean air

to pass back out into the room. This is essential-

ly the same principle applied in air filters and

even air conditioning and heating systems,

which, in addition to regulating temperature,

also remove dust, pollen, and other impurities

from the air.

Various industries use gas filtration, not only

to purify the products released into the atmos-

phere, but also to filter the air in the workplace,

as for instance in a factory room where fine air-

borne particles are a hazard of production. The

gases from power plants that burn coal and oil

are often purified by passing them through filter-

ing systems that collect particles.


AZEOTROPE: A solution that has been

subjected through distillation and is still

not fully separated, but cannot be separat-

ed any further by means of distillation. An

example is ethanol, which cannot be dis-

tilled beyond the point at which it reaches

a 95% concentration with water.

CONDENSATE: The liquid product

that results from distillation.

CONDENSATION: The cooling of a

gas to form a liquid or condensate.

DISTILLATE: The vapor collected from

the volatile material or materials in distilla-

tion. This vapor is then subjected to con-

densation.

DISTILLATION: The use of heat to

separate the components of a liquid and/or

gas. Generally, distillation is used to sepa-

rate mixtures that are homogeneous.

FILTRATE: The liquid or gas separated

from a solid by means of filtration.

FILTRATION: The separation of solids

from a fluid (either gas or liquid) by allow-

ing the fluid to pass through a filter. Gen-

erally, filtration is used to separate mix-

tures that are heterogeneous.




ing different properties. An example is

sand in a container of water.





mixture.


composition, composed of molecules or

atoms of differing types. Compare with

pure substance.

PURE SUBSTANCE: A substance—

either an element or compound—that has

an unvarying composition. This means

that by changing the proportions of atoms,

the result would be an entirely different

substance. Compare with mixture.

SOLUTE: The substance or substances

that are dissolved in a solvent to form a

solution.


in which one or more substances is dis-

solved in an another substance—for exam-


SOLVENT: The substance that dissolves

a solute to form a solution.

SUSPENSION: A term that refers to a

mixture in which solid particles are sus-

pended in a fluid.

VOLATILE: Easily vaporized.

KEY TERMS



Sewage Treatment

Sewage treatment is a complex, multistageprocess whereby waste water is freed of harmfulcontents and rendered safe, so that it can bereturned to the environment. This is a criticalpart of modern life, because when people con-gregate in cities, they generate huge amounts ofwastes that contain disease-causing bacteria,viruses, and other microorganisms. Lack ofproper waste management in ancient andmedieval times led to devastating plagues incities such as Athens, Rome, and Constantinople.

Indeed, one of the factors contributing tothe end of the Athenian golden age of the fifthcentury B.C. was a plague, caused by lack of prop-er sewage-disposal methods, that felled many ofthe city’s greatest figures. The horrors created byimproper waste management reached theirapogee in 1347-51, when microorganisms car-ried by rats brought about the Black Death, inwhich Europe’s population was reduced by one-third.

AEROBIC AND ANAEROBIC

DECAY. Small amounts of sewage can be dealtwith through aerobic (oxygen-consuming)decay, in which microorganisms such as bacteriaand fungi process the toxins in human waste.With larger amounts of waste, however, oxygen isdepleted and the decay mechanism must beanaerobic, or carried out in the absence ofoxygen.

For people who are not connected to amunicipal sewage system, and therefore mustrely on a septic tank for waste disposal, wasteproducts pass through a tank in which they aresubjected to anaerobic decay by varieties ofmicroorganisms that do not require oxygen tosurvive. From there, the waste goes through thedrain field, a network of pipes that allow thewater to seep into a deep-bed filter. In the drainfield, the waste is subjected to aerobic decay byoxygen-dependant microorganisms before iteither filters through the drain pipes into theground, or is evaporated.

MUNICIPAL WASTE TREAT-

MENT. Municipal waste treatment is a muchmore involved process, and will only be

described in the most general terms here. Essen-tially, large-scale sewage treatment requires theseparation of larger particles first, followed bythe separation of smaller particles, as the processcontinues. Along the way, the sewage is chemical-ly treated as well. The separation of smaller solidparticles is aided by aluminum sulfate, whichcauses these particles to settle to the bottommore rapidly.

Through continual processing by filtration,water is eventually clarified of biosolids (that is,feces), but it is still far from being safe. At thatpoint, all the removed biosolids are dried andincinerated; or they may be subjected to addi-tional processes to create fertilizer. The water, oreffluent, is further clarified by filtration in adeep-bed filter, and additional air may be intro-duced to the water to facilitate aerobic decay—for instance, by using algae.

Through a long series of such processes, thewater is rendered safe to be released back into theenvironment. However, wastes containingextremely high levels of toxins may require addi-tional or different forms of treatment.

WHERE TO LEARN MORE

“Case Study: Petroleum Distillation” (Web site).<http://www.pafko.com/history/h_distill.html>(June 6, 2001).

“Classification of Matter” (Web site). <http://www.cards.anoka.k12.mn.us/projects/html>_98/smith/smith2.ht ml (June 6, 2001).

“Fractional Distillation” (Web site). <http://www.geoci-ties.com/chemforum/fdistillation.htm> (June 6,2001).

“Introduction to Distillation” (Web site). <http://lorien.ncl.ac.uk/ming/distil/distil0.htm> (June 6, 2001).

Kellert, Stephen B. and Matthew Black. The Encyclopediaof the Environment. New York: Franklin Watts, 1999.


Seuling, Barbara. Drip! Drop! How Water Gets to YourTap. Illustrated by Nancy Tobin. New York: HolidayHouse, 2000.

“Sewage Treatment” (Web site). <http://www.dcs.exeter.ac.uk/water/sewage.htm> (June 6, 2001).




361


ORGAN IC CHEM ISTRY

ORGAN IC CHEM ISTRY

ORGAN IC CHEM ISTRY

POLYMERS



ORGAN IC CHEM ISTRYOrganic Chemistry

CONCEPTThere was once a time when chemists thought“organic” referred only to things that were living,and that life was the result of a spiritual “lifeforce.” While there is nothing wrong with view-ing life as having a spiritual component, spiritualmatters are simply outside the realm of science,and to mix up the two is as silly as using mathe-matics to explain love (or vice versa). In fact, the“life force” has a name: carbon, the commondenominator in all living things. Not everythingthat has carbon is living, nor are all the areasstudied in organic chemistry—the branch ofchemistry devoted to the study of carbon and itscompounds—always concerned with livingthings. Organic chemistry addresses an array ofsubjects as vast as the number of possible com-pounds that can be made by strings of carbonatoms. We can thank organic chemistry for muchof what makes life easier in the modern age: fuelfor cars, for instance, or the plastics found inmany of the products used in an average day.

HOW IT WORKS

Introduction to Carbon

As the element essential to all of life, and hencethe basis for a vast field of study, carbon isaddressed in its own essay. The Carbon essay, inaddition to examining the chemical properties ofcarbon (discussed below), approaches a numberof subjects, such as the allotropes of carbon.These include three crystalline forms (graphite,diamond, and buckminsterfullerene), as well asamorphous carbon. In addition, two oxides ofcarbon—carbon dioxide and carbon monox-ide—are important, in the case of the former, to

the natural carbon cycle, and in the case of thelatter, to industry. Both also pose environmentaldangers.

The purpose of this summary of the carbonessay is to provide a hint of the complexitiesinvolved with this sixth element on the periodictable, the 14th most abundant element on Earth.In the human body, carbon is second only to oxy-gen in abundance, and accounts for 18% of thebody’s mass. Capable of combining in seeminglyendless ways, carbon, along with hydrogen, is atthe center of huge families of compounds. Theseare the hydrocarbons, present in deposits of fos-sil fuels: natural gas, petroleum, and coal.

A PROPENSITY FOR LIMIT -

LESS BONDING. Carbon has a valenceelectron configuration of 2s22p2. Likewise all themembers of Group 4 on the periodic table(Group 14 in the IUPAC version of the table)—sometimes known as the “carbon family”—haveconfigurations of ns2np2, where n is the numberof the period or row the element occupies on thetable. There are two elements noted for their abil-ity to form long strings of atoms and seeminglyendless varieties of molecules: one is carbon, andthe other is silicon, directly below it on the peri-odic table.

Just as carbon is at the center of a vast net-work of organic compounds, silicon holds thesame function in the inorganic realm. It is foundin virtually all types of rocks, except the calciumcarbonates—which, as their name implies, con-tain carbon. In terms of known elemental mass,silicon is second only to oxygen in abundance onEarth. Silicon atoms are about one and a halftimes as large as those of carbon; thus not evensilicon can compete with carbon’s ability to form


OrganicChemistry


an almost limitless array of molecules in variousshapes and sizes, and with various chemicalproperties.

Electronegativity

Carbon is further distinguished by its high valueof electronegativity, the relative ability of anatom to attract valence electrons. To mention afew basic aspects of chemical bonding, developedat considerably greater length in the ChemicalBonding essay, if two atoms have an electriccharge and thus are ions, they form strong ionicbonds. Ionic bonding occurs when a metal bondswith a nonmetal. The other principal type ofbond is a covalent bond, in which two unchargedatoms share eight valence electrons. If the elec-tronegativity values of the two elements involvedare equal, then they share the electrons equally;but if one element has a higher electronegativityvalue, the electrons will be more drawn to thatelement.

The electronegativity of carbon is certainlynot the highest on the periodic table. That dis-tinction belongs to fluorine, with an electronega-tivity value of 4.0, which makes it the most reac-tive of all elements. Fluorine is at the head ofGroup 7, the halogens, all of which are highly

reactive and most of which have high electroneg-

ativity values. If one ignores the noble gases,

which are virtually unreactive and occupy the

extreme right-hand side of the periodic table,

electronegativity values are highest in the upper

right-hand side of the table—the location of flu-

orine—and lowest in the lower left. In other

words, the value increases with group or column

number (again, leaving out the noble gases in

Group 8), and decreases with period or row

number.

With an electronegativity of 2.5, carbon ties

with sulfur and iodine (a halogen) for sixth place,

behind only fluorine; oxygen (3.5); nitrogen and

chlorine (3.0); and bromine (2.8). Thus its elec-

tronegativity is high, without being too high.

Fluorine is not likely to form the long chains for

which is carbon is known, simply because its

electronegativity is so high, it overpowers other

elements with which it comes into contact. In

addition, with four valence electrons, carbon is

ideally suited to find other elements (or other

carbon atoms) for forming covalent bonds

according to the octet rule, whereby most ele-

ments bond so that they have eight valence

electrons.

“I JUST WANT TO SAY ONE WORD TO YOU...PLASTICS.” SIGNIFYING THAT PLASTICS WERE THE WAVE OF THE FUTURE,THESE WORDS WERE UTTERED TO DUSTIN HOFFMAN’S CHARACTER IN THE 1967 FILM THE GRADUATE. (The Kobal

Collection. Reproduced by permission.)


OrganicChemistry

Carbon’s Multiple Bonds

Carbon—with its four valence electrons—hap-pens to be tetravalent, or capable of bonding tofour other atoms at once. It is not necessarily thecase that an element has the ability to bond withas many other elements as it has valence elec-trons; in fact, this is rarely the case. Additionally,carbon is capable of forming not only a singlebond, with one pair of shared valence electrons,but a double bond (two pairs) or even a triplebond (three pairs.)

Another special property of carbon is itsability to bond in long chains that constitutestrings of carbons and other atoms. Further-more, though sometimes carbon forms a typicalmolecule (for example, carbon dioxide, or CO2, isjust one carbon atom with two oxygens), it is alsocapable of forming “molecules” that are reallynot molecules in the way that the word is typical-ly used in chemistry. Graphite, for instance, isjust a series of “sheets” of carbon atoms bondedtightly in a hexagonal, or six-sided, pattern, whilea diamond is simply a huge “molecule” com-posed of carbon atoms strung together by cova-lent bonds.

Organic Chemistry

Organic chemistry is the study of carbon, itscompounds, and their properties. The onlymajor carbon compounds considered inorganicare carbonates (for instance, calcium carbonate,alluded to above, which is one of the major formsof mineral on Earth) and oxides, such as carbondioxide and carbon monoxide. This leaves a hugearray of compounds to be studied, as we shall see.

The term “organic” in everyday languageconnotes “living,” but organic chemistry isinvolved with plenty of compounds not part ofliving organisms: petroleum, for instance, is anorganic compound that ultimately comes fromthe decayed bodies of organisms that once werealive. It should be stressed that organic com-pounds do not have to be produced by livingthings, or even by things that once were alive;they can be produced artificially in a laboratory.

The breakthrough event in organic chem-istry came in 1828, when German chemistFriedrich Wöhler (1800-1882) heated a sampleof ammonium cyanate (NH4OCN) and convert-ed it to urea (H2N-CO-NH2). Ammonium cyan-ite is an inorganic material, whereas urea, a wasteproduct in the urine of mammals, is an organicone. “Without benefit of a kidney, a bladder, or a


dog,” as Wöhler later said, he had managed to transform an inorganic substance into anorganic one.

Ammonium cyanate and urea are isomers:substances having the same formula, but possess-ing different chemical properties. Thus they haveexactly the same numbers and proportions ofatoms, yet these atoms are arranged in differentways. In urea, the carbon forms an organic chain,and in ammonium cyanate, it does not. Thus, toreduce the specifics of organic chemistry evenfurther, this discipline can be said to constitutethe study of carbon chains, and ways to rearrangethem to create new substances.


Organic Chemistry and Modern Life

At first glance, the term “organic chemistry”might sound like something removed from

IT IS VIRTUALLY IMPOSSIBLE FOR A PERSON IN MODERN

AMERICA TO SPEND AN ENTIRE DAY WITHOUT COMING

INTO CONTACT WITH AT LEAST ONE, AND MORE LIKELY

DOZENS, OF PLASTIC PRODUCTS, ALL MADE POSSIBLE

BY ORGANIC CHEMISTRY. HERE, A FASHION MODEL

PARADES A DRESS MADE OF PVC. (Reuters NewMedia

Inc./Corbis. Reproduced by permission.)


OrganicChemistry


everyday life, but this could not be further from

the truth. The reality of the role played by organ-

ic chemistry in modern existence is summed up

in a famous advertising slogan used by E. I. du

Pont de Nemours and Company (usually

referred to as “du Pont”): “Better Things for Bet-

ter Living Through Chemistry.”

Often rendered simply as “Better Living

Through Chemistry,” the advertising campaign

made its debut in 1938, just as du Pont intro-

duced a revolutionary product of organic chem-

istry: nylon, the creation of a brilliant young

chemist named Wallace Carothers (1896-1937).

Nylon, an example of a polymer (discussed

below), started a revolution in plastics that was

still unfolding three decades later, in 1967. That

was the year of the film The Graduate, which

included a famous interchange between the char-

acter of Benjamin Braddock (Dustin Hoffman)

and an adult named Mr. McGuire (Walter

Brooke):

• Mr. McGuire: I just want to say one word to

you... just one word.

• Benjamin Braddock: Yes, sir.

• Mr. McGuire: Are you listening?

• Benjamin Braddock: Yes, sir, I am.

• Mr. McGuire: Plastics.

The meaning of this interchange was thatplastics were the wave of the future, and that anintelligent young man such as Ben should investhis energies in this promising new field. Instead,Ben puts his attention into other things, quiteremoved from “plastics,” and much of the plotrevolves around his revolt against what he per-ceives as the “plastic” (that is, artificial) characterof modern life.

In this way, The Graduate spoke for a wholegeneration that had become ambivalent concern-ing “better living through chemistry,” a phrasethat eventually was perceived as ironic in view ofconcerns about the environment and the manyartificial products that make up modern life.Responding to this ambivalence, du Pontdropped the slogan in the late 1970s; yet the real-ity is that people truly do enjoy “better livingthrough chemistry”—particularly organic chemistry.

APPLICATIONS OF ORGANIC

CHEMISTRY. What would the world be likewithout the fruits of organic chemistry? First, itwould be necessary to take away all the variousforms of rubber, vitamins, cloth, and paper madefrom organically based compounds. Aspirins andall types of other drugs; preservatives that keepfood from spoiling; perfumes and toiletries; dyes

ANOTHER BYPRODUCT OF ORGANIC CHEMISTRY: PETROLEUM JELLY. (Laura Dwight/Corbis. Reproduced by permission.)


OrganicChemistry

and flavorings—all these things would have to goas well.

Synthetic fibers such as nylon—used ineverything from toothbrushes to parachutes—would be out of the picture if it were not for theenormous progress made by organic chemistry.The same is true of plastics or polymers in gen-eral, which have literally hundreds upon hun-dreds of applications. Indeed, it is virtuallyimpossible for a person in twenty-first centuryAmerica to spend an entire day without cominginto contact with at least one, and more likelydozens, of plastic products. Car parts, toys, com-puter housings, Velcro fasteners, PVC (polyvinylchloride) plumbing pipes, and many more fix-tures of modern life are all made possible by plas-tics and polymers.

Then there is the vast array of petrochemi-cals that power modern civilization. Best-knownamong these is gasoline, but there is also coal, stillone of the most significant fuels used in electri-cal power plants, as well as natural gas and vari-ous other forms of oil used either directly orindirectly in providing heat, light, and electricpower to homes. But the influence of petrochem-icals extends far beyond their applications forfuel. For instance, the roofing materials and tarthat (quite literally) keep a roof over people’sheads, protecting them from sun and rain, are theproduct of petrochemicals—and ultimately, oforganic chemistry.

Hydrocarbons

Carbon, together with other elements, forms somany millions of organic compounds that evenintroductory textbooks on organic chemistryconsist of many hundreds of pages. Fortunately,it is possible to classify broad groupings oforganic compounds. The largest and most signif-icant is that class of organic compounds knownas hydrocarbons—chemical compounds whosemolecules are made up of nothing but carbonand hydrogen atoms.

Every molecule in a hydrocarbon is builtupon a “skeleton” composed of carbon atoms,either in closed rings or in long chains. Thechains may be straight or branched, but in eachcase—rings or chains, straight chains orbranched ones—the carbon bonds not used intying the carbon atoms together are taken up byhydrogen atoms.

Theoretically, there is no limit to the numberof possible hydrocarbons. Not only does carbonform itself into apparently limitless molecularshapes, but hydrogen is a particularly good part-ner. It has the smallest atom of any element onthe periodic table, and therefore it can bond toone of carbon’s valence electrons without gettingin the way of the other three.

There are two basic varieties of hydrocar-bon, distinguished by shape: aliphatic and aro-matic. The first of these forms straight orbranched chains, as well as rings, while the sec-ond forms only benzene rings, discussed below.Within the aliphatic hydrocarbons are three vari-eties: those that form single bonds (alkanes),double bonds (alkenes), and triple bonds(alkynes.)

ALKANES. The alkanes are also knownas saturated hydrocarbons, because all the bondsnot used to make the skeleton itself are filled totheir capacity (that is, saturated) with hydrogenatoms. The formula for any alkane is CnH2n+2,where n is the number of carbon atoms. In thecase of a linear, unbranched alkane, every carbonatom has two hydrogen atoms attached, but the two end carbon atoms each have an extrahydrogen.

What follows are the names and formulasfor the first eight normal, or unbranched, alka-nes. Note that the first four of these receivedcommon names before their structures wereknown; from C5 onward, however, they weregiven names with Greek roots indicating thenumber of carbon atoms (e.g., octane, a refer-ence to “eight.”)

• Methane (CH4)

• Ethane (C2H6)

• Propane (C3H8)

• Butane (C4H10)

• Pentane (C5H12)

• Hexane (C6H14)

• Heptane (C7H16)

• Octane (C8H18)

The reader will undoubtedly notice a num-ber of familiar names on this list. The first four,being the lowest in molecular mass, are gases atroom temperature, while the heavier ones areoily liquids. Alkanes even heavier than those onthis list tend to be waxy solids, an example beingparaffin wax, for making candles. It should benoted that from butane on up, the alkanes havenumerous structural isomers, depending on



OrganicChemistry

whether they are straight or branched, and theseisomers have differing chemical properties.

Branched alkanes are named by indicatingthe branch attached to the principal chain.Branches, known as substituents, are named bytaking the name of an alkane and replacing thesuffix with yl—for example, methyl, ethyl, and soon. The general term for an alkane which func-tions as a substituent is alkyl.

Cycloalkanes are alkanes joined in a closedloop to form a ring-shaped molecule. They arenamed by using the names above, with cyclo- asa prefix. These start with propane, or rathercyclopropane, which has the minimum numberof carbon atoms to form a closed shape: threeatoms, forming a triangle.

ALKENES AND ALKYNES. Thenames of the alkenes, hydrocarbons that containone or more double bonds per molecule, are par-allel to those of the alkanes, but the family end-ing is -ene. Likewise they have a common formu-la: CnH2n. Both alkenes and alkynes, discussedbelow, are unsaturated—in other words, some ofthe carbon atoms in them are free to form otherbonds. Alkenes with more than one double bondare referred to as being polyunsaturated.

As with the alkenes, the names of alkynes(hydrocarbons containing one or more triplebonds per molecule) are parallel to those of thealkanes, only with the replacement of the suffix -yne in place of -ane. The formula for alkenes isCnH2n-2. Among the members of this group areacetylene, or C2H2, used for welding steel. Plasticpolystyrene is another important product fromthis division of the hydrocarbon family.

AROMATIC HYDROCARBONS.

Aromatic hydrocarbons, despite their name, donot necessarily have distinctive smells. In fact thename is a traditional one, and today these com-pounds are defined by the fact that they havebenzene rings in the middle. Benzene has a for-mula C6H6, and a benzene ring is usually repre-sented as a hexagon (the six carbon atoms andtheir attached hydrogen atoms) surrounding acircle, which represents all the bonding electronsas though they were everywhere in the moleculeat once.

In this group are products such as naphtha-lene, toluene, and dimethyl benzene. These lasttwo are used as solvents, as well as in the synthe-sis of drugs, dyes, and plastics. One of the morefamous (or infamous) products in this part of

the vast hydrocarbon network is trinitrotoluene,or TNT. Naphthalene is derived from coal tar,and used in the synthesis of other compounds.A crystalline solid with a powerful odor, it isfound in mothballs and various deodorant-disinfectants.

PETROCHEMICALS. As for petro-chemicals, these are simply derivatives of petrole-um, itself a mixture of alkanes with somealkenes, as well as aromatic hydrocarbons.Through a process known as fractional distilla-tion, the petrochemicals of the lowest molecularmass boil off first, and those having higher massseparate at higher temperatures.

Among the products derived from the frac-tional distillation of petroleum are the following,listed from the lowest temperature range (that is,the first material to be separated) to the highest:natural gas; petroleum ether, a solvent; naphtha,a solvent (used for example in paint thinner);gasoline; kerosene; fuel for heating and dieselfuel; lubricating oils; petroleum jelly; paraffinwax; and pitch, or tar. A host of other organicchemicals, including various drugs, plastics,paints, adhesives, fibers, detergents, syntheticrubber, and agricultural chemicals, owe theirexistence to petrochemicals.

Obviously, petroleum is not just for makinggasoline, though of course this is the first prod-uct people think of when they hear the word“petroleum.” Not all hydrocarbons in gasolineare desirable. Straight-chain or normal heptane,for instance, does not fire smoothly in an inter-nal-combustion engine, and therefore disruptsthe engine’s rhythm. For this reason, it is given arating of zero on a scale of desirability, whileoctane has a rating of 100. This is why gas sta-tions list octane ratings at the pump: the higherthe presence of octane, the better the gas is forone’s automobile.

Hydrocarbon Derivatives

With carbon and hydrogen as the backbone, thehydrocarbons are capable of forming a vast arrayof hydrocarbon derivatives by combining withother elements. These other elements arearranged in functional groups—an atom orgroup of atoms whose presence identifies a spe-cific family of compounds. Below we will brieflydiscuss some of the principal hydrocarbon deriv-



OrganicChemistry

atives, which are basically hydrocarbons with theaddition of other molecules or single atoms.

Alcohols are oxygen-hydrogen moleculeswedded to hydrocarbons. The two most impor-tant commercial types of alcohol are methanol,or wood alcohol; and ethanol, which is found inalcoholic beverages, such as beer, wine, andliquor. Though methanol is still known as “woodalcohol,” it is no longer obtained by heatingwood, but rather by the industrial hydrogena-tion of carbon monoxide. Used in adhesives,fibers, and plastics, it can also be applied as a fuel.Ethanol, too, can be burned in an internal-combustion engine, when combined with gaso-line to make gasohol. Another significant alcoholis cholesterol, found in most living organisms.Though biochemically important, cholesterolcan pose a risk to human health.

Aldehydes and ketones both involve a dou-ble-bonded carbon-oxygen molecule, known as acarbonyl group. In a ketone, the carbonyl groupbonds to two hydrocarbons, while in an alde-hyde, the carbonyl group is always at the end of ahydrocarbon chain. Therefore, instead of twohydrocarbons, there is always a hydrocarbon andat least one other hydrogen bonded to the carbonatom in the carbonyl. One prominent example ofa ketone is acetone, used in nail polish remover.Aldehydes often appear in nature—for instance,as vanillin, which gives vanilla beans their pleas-ing aroma. The ketones carvone and camphorimpart the characteristic flavors of spearmintleaves and caraway seeds.

CARBOXYLIC ACIDS AND ES-

TERS. Carboxylic acids all have in commonwhat is known as a carboxyl group, designated bythe symbol -COOH. This consists of a carbonatom with a double bond to an oxygen atom, anda single bond to another oxygen atom that is, inturn, wedded to a hydrogen. All carboxylic acidscan be generally symbolized by RCOOH, with Ras the standard designation of any hydrocarbon.Lactic acid, generated by the human body, is acarboxylic acid: when a person overexerts, themuscles generate lactic acid, resulting in a feelingof fatigue until the body converts the acid towater and carbon dioxide. Another example of acarboxylic acid is butyric acid, responsible in partfor the smells of rancid butter and human sweat.

When a carboxylic acid reacts with an alco-hol, it forms an ester. An ester has a structuresimilar to that described for a carboxylic acid,

with a few key differences. In addition to itsbonds (one double, one single) with the oxygenatoms, the carbon atom is also attached to ahydrocarbon, which comes from the carboxylicacid. Furthermore, the single-bonded oxygenatom is attached not to a hydrogen, but to a sec-ond hydrocarbon, this one from the alcohol. Onewell-known ester is acetylsalicylic acid—betterknown as aspirin. Esters, which are a key factor inthe aroma of various types of fruit, are oftennoted for their pleasant smell.

Polymers

Polymers are long, stringy molecules made ofsmaller molecules called monomers. They appearin nature, but thanks to Carothers—a tragic fig-ure, who committed suicide a year before Nylonmade its public debut—as well as other scientistsand inventors, synthetic polymers are a funda-mental part of daily life.

The structure of even the simplest polymer,polyethylene, is far too complicated to discuss inordinary language, but must be represented bychemical symbolism. Indeed, polymers are a sub-ject unto themselves, but it is worth noting herejust how many products used today involve poly-mers in some form or another.

Polyethylene, for instance, is the plastic usedin garbage bags, electrical insulation, bottles, anda host of other applications. A variation on poly-ethylene is Teflon, used not only in nonstickcookware, but also in a number of other devices,such as bearings for low-temperature use. Poly-mers of various kinds are found in siding forhouses, tire tread, toys, carpets and fabrics, and a variety of other products far too lengthy to enumerate.

WHERE TO LEARN MORE

Blashfield, Jean F. Carbon. Austin, TX: Raintree Steck-Vaughn, 1999.

“Carbon.” Xrefer (Web site). <http://www.xrefer.com/entry/639742> (May 30, 2001).

Chemistry Help Online for Students (Web site). <http://members.tripod.com/chemistryhelp/> May 30,2001).


Loudon, G. Marc. Organic Chemistry. Menlo Park, CA:Benjamin/Cummings, 1988.



OrganicChemistry


ALKANES: Hydrocarbons that form

single bonds. Alkanes are also called satu-

rated hydrocarbons.

ALKENES: Hydrocarbons that form

double bonds.

ALKYL: A general term for an alkane

that functions as a substituent.

ALKYNES: Hydrocarbons that form

triple bonds.



ular structure.

AMORPHOUS: Having no definite

structure.




CRYSTALLINE: A term describing a

type of solid in which the constituent parts

have a simple and definite geometric

arrangement repeated in all directions.




single bonds and triple bonds.


ability of an atom to attract valence elec-

trons.

FUNCTIONAL GROUPS: An atom or

group of atoms whose presence identifies a

specific family of compounds.




HYDROCARBON DERIVATIVES: Fam-

ilies of compounds formed by the joining

of hydrocarbons with various functional

groups.

ISOMERS: Substances having the same

chemical formula, but that are different

chemically due to disparities in the






electrons.

ORGANIC CHEMISTRY: The study of

carbon, its compounds, and their proper-

ties. (Some carbon-containing com-

pounds, most notably oxides and carbon-

ates, are not considered organic.)

SATURATED: A term describing a

hydrocarbon in which each carbon is

already bound to four other atoms. Alka-

nes are saturated hydrocarbons.

SINGLE BOND: A form of bonding in

which two atoms share one pair of valence

electrons. Carbon is also capable of double

bonds and triple bonds.

SUBSTITUENTS: Branches of alka-

nes, named by taking the name of an alka-

ne and replacing the suffix with yl—for

example, methyl, ethyl, and so on.

TETRAVALENT: Capable of bonding

to four other elements.

TRIPLE BOND: A form of bonding in

which two atoms share three pairs of


single bonds and double bonds.


hydrocarbon in which the carbons

involved in a multiple bond (a double

bond or triple bond) are free to bond with

other atoms. Alkenes and alkynes are both

unsaturated.





KEY TERMS


OrganicChemistry

“Organic Chemistry” (Web site). <http://edie.cprost.sfu.ca/~rhlogan/organic.html> (May 30, 2001).

“Organic Chemistry.” Frostburg State University Chem-istry Helper (Web site). <http://www.chemhelper.com/> (May 30, 2001).

Sparrow, Giles. Carbon. New York: Benchmark Books, 1999.





POLYMERSPolymers

CONCEPTFormed from hydrocarbons, hydrocarbon deriv-atives, or sometimes from silicon, polymers arethe basis not only for numerous natural materi-als, but also for most of the synthetic plastics thatone encounters every day. Polymers consist ofextremely large, chain-like molecules that are, inturn, made up of numerous smaller, repeatingunits called monomers. Chains of polymers canbe compared to paper clips linked together inlong strands, and sometimes cross-linked toform even more durable chains. Polymers can becomposed of more than one type of monomer,and they can be altered in other ways. Likewisethey are created by two different chemicalprocesses, and thus are divided into addition andcondensation polymers. Among the naturalpolymers are wool, hair, silk, rubber, and sand,while the many synthetic polymers includenylon, synthetic rubber, Teflon, Formica, Dacron,and so forth. It is very difficult to spend a daywithout encountering a natural polymer—evenif hair is removed from the list—but in the twenty-first century, it is probably even harder toavoid synthetic polymers, which have collectivelyrevolutionized human existence.

HOW IT WORKS

Polymers of Silicon and Carbon

Polymers can be defined as large, typically chain-like molecules composed of numerous smaller,repeating units known as monomers. There arenumerous varieties of monomers, and sincethese can be combined in different ways to formpolymers, there are even more of the latter.

The name “polymer” does not, in itself,define the materials that polymers contain. Ahandful of polymers, such as natural sand or synthetic silicone oils and rubbers, are builtaround silicon. However, the vast majority ofpolymers center around an element that occupiesa position just above silicon on the periodic table:carbon.

The similarities between these two are sogreat, in fact, that some chemists speak of Group4 (Group 14 in the IUPAC system) on the peri-odic table as the “carbon family.” Both carbonand silicon have the ability to form long chains ofatoms that include bonds with other elements.The heavier elements of this “family,” however(most notably lead), are made of atoms too big toform the vast array of chains and compounds forwhich silicon and carbon are noted.

Indeed, not even silicon—though it is at thecenter of an enormous range of inorganic com-pounds—can compete with carbon in its abilityto form arrangements of atoms in various shapesand sizes, and hence to participate in an almostlimitless array of compounds. The reason, inlarge part, is that carbon atoms are much smallerthan those of silicon, and thus can bond to oneanother and still leave room for other bonds.

Carbon is such an important element thatan entire essay in this book is devoted to it, whilea second essay discusses organic chemistry, thestudy of compounds containing carbon. In thepresent context, there will be occasional refer-ences to non-carbon (that is, silicon) polymers,but the majority of our attention will be devotedto hydrocarbon and hydrocarbon-derivativepolymers, which most of us know simply as“plastics.”


Polymers


Organic Chemistry

As explained in the essay on Organic Chemistry,chemists once defined the term “organic” asrelating only to living organisms; the materialsthat make them up; materials derived from them;and substances that come from formerly livingorganisms. This definition, which more or lessrepresents the everyday meaning of “organic,”includes a huge array of life forms and materials:humans, all other animals, insects, plants,microorganisms, and viruses; all substances thatmake up their structures (for example, blood,DNA, and proteins); all products that come fromthem (a list diverse enough to encompass every-thing from urine to honey); and all materialsderived from the bodies of organisms that wereonce alive (paper, for instance, or fossil fuels).

As broad as this definition is, it is not broadenough to represent all the substances addressedby organic chemistry—the study of carbon, itscompounds, and their properties. All living oronce-living things do contain carbon; however,organic chemistry is also concerned with carbon-containing materials—for instance, the syntheticplastics we will discuss in this essay—that havenever been part of a living organism.

It should be noted that while organic chem-istry involves only materials that contain carbon,

carbon itself is found in other compounds notconsidered organic: oxides such as carbon diox-ide and monoxide, as well as carbonates, mostnotably calcium carbonate or limestone. In otherwords, as broad as the meaning of “organic” is, itstill does not encompass all substances contain-ing carbon.

Hydrocarbons

As for hydrocarbons, these are chemical com-pounds whose molecules are made up of nothingbut carbon and hydrogen atoms. Every moleculein a hydrocarbon is built upon a “skeleton” ofcarbon atoms, either in closed rings or in longchains, which are sometimes straight and some-times branched.

Theoretically, there is no limit to the numberof possible hydrocarbons: not only does carbonform itself into seemingly limitless molecularshapes, but hydrogen is a particularly good part-ner. It is the smallest atom of any element on theperiodic table, and therefore it can bond to oneof carbon’s four valence electrons without gettingin the way of the other three.

There are many, many varieties of hydrocar-bon, classified generally as aliphatic hydrocar-bons (alkanes, alkenes, and alkynes) and aromat-ic hydrocarbons, the latter being those that con-

COTTON IS AN EXAMPLE OF A NATURAL POLYMER. (Carl Corey/Corbis. Reproduced by permission.)


Polymers


tain a benzene ring. By means of a basic alter-

ation in the shape or structure of a hydrocarbon,

it is possible to create new varieties. Thus, as

noted above, the number of possible hydrocar-

bons is essentially unlimited.

Certain hydrocarbons are particularly use-

ful, one example being petroleum, a term that

refers to a wide array of hydrocarbons. Among

these is an alkane that goes by the name of octane

(C8H18), a preferred ingredient in gasoline.

Hydrocarbons can be combined with various

functional groups (an atom or group of atoms

whose presence identifies a specific family of

compounds) to form hydrocarbon derivatives

such as alcohols and esters.


Types of Polymers and Polymerization

Many polymers exist in nature. Among these aresilk, cotton, starch, sand, and asbestos, as well asthe incredibly complex polymers known as RNA(ribonucleic acid) and DNA (deoxyribonucleicacid), which hold genetic codes. The polymersdiscussed in this essay, however, are primarily ofthe synthetic kind. Artificial polymers includesuch plastics (defined below) as polyethylene,styrofoam, and Saran wrap; fibers such as nylon,Dacron (polyester), and rayon; and other materi-als such as Formica, Teflon, and PVC pipe.

MODERN APPLIANCES CONTAIN NUMEROUS EXAMPLES OF SYNTHETIC POLYMERS, FROM THE FLOORING TO THE COUN-TERTOPS TO VIRTUALLY ALL APPLIANCES. (Scott Roper/Corbis. Reproduced by permission.)


PolymersAs noted earlier, most polymers are formedfrom monomers either of hydrocarbon or hydro-carbon derivatives. The most basic syntheticmonomer is ethylene (C2H4), a name whose -eneending identifies it as an alkene, a hydrocarbonformed by double bonds between carbon atoms.Another alkene hydrocarbon monomer is buta-diene, whose formula is C4H6. This is an exampleof the fact that the formula of a compound doesnot tell the whole story: on paper, the differencebetween these two appears to be merely a matterof two extra atoms each of carbon and hydro-gen. In fact, butadiene’s structure is much morecomplex.

Still more complex is styrene, which includesa benzene ring. Several other monomers involveother elements: chloride, in vinyl chloride; nitro-gen, in acrylonitrile; and fluorine, in tetrafluo-roethylene. It is not necessary, in the present con-text, to keep track of all of these substances,which in any case represent just some of themore prominent among a wide variety of syn-thetic monomers. A good high-school or collegechemistry textbook (either general chemistry ororganic chemistry) should provide structuralrepresentations of these common monomers.Such representations will show, for instance, thevast differences between purely hydrocarbonmonomers such as ethylene, propylene, styrene,and butadiene.

When combined into polymers, themonomers above form the basis for a variety ofuseful and familiar products. Once the carbondouble bonds in tetrafluoroethylene (C2F4) arebroken, they form the polymer known as Teflon,used in the coatings of cooking utensils, as well asin electrical insulation and bearings. Vinyl chlo-ride breaks its double bonds to form polyvinylchloride, better known as PVC, a material usedfor everything from plumbing pipe to toys toSaran wrap. Styrene, after breaking its doublebonds, forms polystyrene, used in containers andthermal insulation.

POLYMERIZATION. Note that sever-al times in the preceding paragraph, there was areference to the breaking of carbon doublebonds. This is often part of one variety of poly-merization, the process whereby monomers jointo form polymers. If monomers of a single typejoin, the resulting polymer is called a homopoly-mer, but if the polymer consists of more than onetype of monomer, it is known as a copolymer.This joining may take place by one of two


processes. The first of these, addition polymer-ization, is fairly simple: monomers add them-selves to one another, usually breaking doublebonds in the process. This results in the creationof a polymer and no other products.

Much more complex is the process known ascondensation polymerization, in which a smallmolecule called a dimer is formed as monomersjoin. The specifics are too complicated to discussin any detail, but a few things can be said hereabout condensation polymerization. Themonomers in condensation polymerization mustbe bifunctional, meaning that they have a func-tional group at each end. When characteristicstructures at the ends of the monomers react toone another by forming a bond, they create adimer, which splits off from the polymer. Theproducts of condensation polymerization arethus not only the polymer itself, but also a dimer,which may be water, hydrochloric acid (HCl), orsome other substance.

A Plastic World

A DAY IN THE LIFE. Though “plas-tic” has a number of meanings in everyday life,

A SYNTHETIC POLYMER KNOWN AS KEVLAR IS USED IN

THE CONSTRUCTION OF BULLETPROOF VESTS FOR LAW-ENFORCEMENT OFFICERS. (Anna Clopet/Corbis. Reproduced by

permission.)


Polymers and in society at large (as we shall see), the scien-tific definition is much more specific. Plastics arematerials, usually organic, that can be caused toflow under certain conditions of heat and pres-sure, and thus to assume a desired shape whenthe pressure and temperature conditions arewithdrawn. Most plastics are made of polymers.

Every day, a person comes into contact withdozens, if not hundreds, of plastics and poly-mers. Consider a day in the life of a hypotheticalteenage girl. She gets up in the morning, brushesher teeth with a toothbrush made of nylon, thenopens a shower door—which is likely to be plas-tic rather than glass—and steps into a moldedplastic shower or bathtub. When she gets out ofthe shower, she dries off with a towel containinga polymer such as rayon, perhaps while standingon tile that contains plastics, or polymers.

She puts on makeup (containing polymers)that comes in plastic containers, and later blow-dries her hair with a handheld hair dryer made ofinsulated plastic. Her clothes, too, are likely tocontain synthetic materials made of polymers.When she goes to the kitchen for breakfast, shewill almost certainly walk on flooring with a plas-tic coating. The countertops may be of formica, acondensation polymer, while it is likely that vir-tually every appliance in the room will containplastic. If she opens the refrigerator to get out amilk container, it too will be made of plastic, orof paper with a thin plastic coating. Much of thepackaging on the food she eats, as well as sand-wich bags and containers for storing food, is alsomade of plastic.

And so it goes throughout the day. Thephone she uses to call a friend, the computer shesits at to check her e-mail, and the stereo in herroom all contain electrical components housedin plastic. If she goes to the gym, she may workout in Gore-tex, a fabric containing a very thinlayer of plastic with billions of tiny pores, so thatit lets through water vapor (that is, perspiration)without allowing the passage of liquid water. Onthe way to the health club, she will ride in a carthat contains numerous plastic molds in thesteering wheel and dashboard. If she plays a com-pact disc—itself a thin wafer of plastic coatedwith metal—she will pull it out of a plastic jewelcase. Finally, at night, chances are she will sleep insheets, and with a pillow, containing syntheticpolymers.

A SILENT REVOLUTION. Thescenario described above—a world surroundedby polymers, plastics, and synthetic materials—represents a very recent phenomenon. “Beforethe 1930s,” wrote John Steele Gordon in an arti-cle about plastics for American Heritage, “almosteverything people saw or handled was made ofmaterials that had been around since ancienttimes: wood, stone, metal, and animal and plantfibers.” All of that changed in the era just beforeWorld War II, thanks in large part to a brilliantyoung American chemist named WallaceCarothers (1896-1937).

By developing nylon for E. I. du Pont deNemours and Company (known simply as“DuPont” or “du Pont”), Carothers and his col-leagues virtually laid the foundation for modernpolymer chemistry—a field that employs morechemists than any other. These men created whatGordon called a “materials revolution” by intro-ducing the world to polymers and plastics, whichare typically made of polymers.

Yet as Gordon went on to note, “It has beena curiously silent revolution.... When we think ofthe scientific triumphs of [the twentieth centu-ry], we think of nuclear physics, medicine, spaceexploration, and the computer. But all thesedevelopments would have been much impeded,in some cases impossible, without... plastics. Andyet ‘plastic’ remains, as often as not, a term ofopprobrium.”

AMBIVALENCE TOWARD PLAS-

TICS. Gordon was alluding to a cultural atti-tude discussed in the essay on Organic Chem-istry: the association of plastics, a physical mate-rial developed by chemical processes, with thecondition—spiritual, moral, and intellectual—ofbeing “plastic” or inauthentic. This was symbol-ized in a famous piece of dialogue about plasticsfrom the 1967 movie The Graduate, in which anonplussed Ben Braddock (Dustin Hoffman) lis-tens as one of his parents’ friends advises him toinvest his future in plastics. As Gordon noted,“however intergenerationally challenged thathalf-drunk friend of Dustin Hoffman’s parentsmay have been... he was right about the impor-tance of the materials revolution in the twentiethcentury.”

One aspect of society’s ambivalence overplastics relates to very genuine concerns aboutthe environment. Most synthetic polymers aremade from petroleum, a nonrenewable resource;



Polymersbut this is not the greatest environmental dangerthat plastics present. Most plastics are notbiodegradable: though made of organic materi-als, they do not contain materials that willdecompose and eventually return to the ground.Nor is there anything in plastics to attractmicroorganisms, which, by assisting in thedecomposition of organic materials, help to facil-itate the balance of decay and regeneration nec-essary for life on Earth.

Efforts are underway among organicchemists in the research laboratories of corpora-tions and other institutions to developbiodegradable plastics that will speed up thedecomposition of materials in the polymers—aprocess that normally takes decades. Until suchreplacement polymers are developed, however,the most environmentally friendly solution tothe problem of plastics is recycling. Today onlyabout 1% of plastics are recycled, while the restgoes into waste dumps, where they account for30% of the volume of trash.

Long before environmental concerns cameto the forefront, however, people had begunalmost to fear plastics as a depersonalizing aspectof modern life. It seemed that in a given day, aperson touched fewer and fewer things that camedirectly from the natural environment: the“wood, stone, metal, and animal and plantfibers” to which Gordon alluded. Plastics seemedto have made human life emptier; yet the truth ofthe matter—including the fact that plastics addmore than they take away from the landscape ofour world—is much more complex.

The Plastics Revolution

Though the introduction of plastics is typicallyassociated with the twentieth century, in fact the“materials revolution” surrounding plasticsbegan in 1865. That was the year when Englishchemist Alexander Parkes (1813-1890) producedthe first plastic material, celluloid. Parkes couldhave become a rich man from his invention, buthe was not a successful marketer. Instead, theman who enjoyed the first commercial success inplastics was—not surprisingly—an American,inventor John Wesley Hyatt (1837-1920).

Responding to a contest in which a billiard-ball manufacturer offered $10,000 to anyone whocould create a substitute for ivory, which wasextremely costly, Hyatt turned to Parkes’s cellu-loid. Actually, Parkes had given his creation—

developed from cellulose, a substance found inthe cell walls of plants—a much less appealingname, “Parkesine.” Hyatt, who used celluloid tomake smooth, hard, round billiard balls (therebywinning the contest) took out a patent for theprocess involved in making the material he haddubbed “Celluloid,” with a capital C.

Though the Celluloid made by Hyatt’sprocess was flammable (as was Parkesine), itproved highly successful as a product when heintroduced it in 1869. He marketed it successful-ly for use in items such as combs and baby rattles,and Celluloid sales received a powerful boostafter photography pioneer George Eastman(1854-1932) chose the material for use in thedevelopment of film. Eventually, Celluloid wouldbe applied in motion-picture film, and eventoday, the adjective “celluloid” is sometimes usedin relation to the movies. Actually, Celluloid(which can be explosive in large quantities) wasphased out in favor of “safety film,” or celluloseacetate, beginning in 1924.

Two important developments in the creationof synthetic polymers occurred at the turn of thecentury. One was the development of Galalith, anivory-like substance made from formaldehydeand milk, by German chemist Adolf Spitteler. Aneven more important innovation happened in1907, when Belgian-American chemist LeoBaekeland (1863-1944) introduced Bakelite. Thelatter, created in a reaction between phenol andformaldehyde, was a hard, black plastic thatproved an excellent insulator. It soon foundapplication in making telephones and householdappliances, and by the 1920s, chemists had fig-ured out how to add pigments to Bakelite, thusintroducing the public to colored plastics.

SYNTHETIC RUBBER. Through-out these developments, chemists had only avague understanding of polymers, but by the1930s, they had come to accept the model ofpolymers as large, flexible, chain-like molecules.One of the most promising figures in the emerg-ing field of polymer chemistry was Carothers,who in 1926 left a teaching post at Harvard Uni-versity to accept a position as director of thepolymer research laboratory at DuPont.

Among the first problems Carothers tackledwas the development of synthetic rubber. Natur-al rubber had been known for many centurieswhen English chemist Joseph Priestley (1733-1804) gave it its name because he used it to rub



Polymers out pencil marks. In 1839, American inventorCharles Goodyear (1800-1860) accidentally dis-covered a method for making rubber moredurable, after he spilled a mixture of rubber andsulfur onto a hot stove. Rather than melting, therubber bonded with the sulfur to form a muchstronger but still elastic product, and Goodyearsoon patented this process under the name vulcanization.

Natural rubber, nonetheless, had manyundesirable properties, and hence DuPont putCarothers to the task of developing a substitute.The result was neoprene, which he created byadding a chlorine atom to an acetylene deriva-tive. Neoprene was stronger, more durable, andless likely to become brittle in cold weather thannatural rubber. It would later prove an enormousboost to the Allied war effort, after the Japaneseseized the rubber plantations of Southeast Asia in 1941.

NYLON. Had neoprene, which Carothersdeveloped in 1931, been the extent of his achieve-ments, he would still be remembered by sciencehistorians. However, his greatest creation still layahead of him. Studying the properties of silk, hebecame convinced that he could develop a moredurable compound that could replicate the prop-erties of silk at a much lower cost.

Carothers was not alone in his efforts, asGordon showed in his account of events at theDuPont laboratories:

One day, an assistant, Julian Hill, noticedthat when he stuck a glass stirring rod into agooey mass at the bottom of a beaker theresearchers had been investigating, he could drawout threads from it, the polymers forming spon-taneously as he pulled. When Carothers wasabsent one day, Hill and his colleagues decided tosee how far they could go with pulling threadsout of goo by having one man hold the beakerwhile another ran down the hall with the glassrod. A very long, silk-like thread was produced.

Realizing what they had on their hands,DuPont devoted $27 million to the researchefforts of Carothers and his associates at the lab,and in 1937, Carothers presented his boss withthe results, saying “Here is your synthetic textilefabric.” DuPont introduced the material, nylon,to the American public the following year withone of the most famous advertising campaigns ofall time: “Better Things for Better LivingThrough Chemistry.”

The product got an additional boostthrough exposure at the 1939 World’s Fair.When DuPont put 4,000 pairs of nylon stock-ings on the market, they sold in a matter ofhours. A few months later, four million pairssold in New York City in a single day. Womenstood in line to buy stockings of nylon, a muchbetter (and less expensive) material for thatpurpose than silk—but they did not have longto enjoy it. During World War II, all nylon wentinto making war materials such as parachutes,and nylon did not become commercially avail-able again until 1946.

As Gordon noted, Carothers would surelyhave won the Nobel Prize in chemistry for hiswork—“but Nobel prizes go only to living recip-ients....” Carothers had married in 1936, and byearly 1937, his wife Helen was pregnant. (Pre-sumably, he was unaware of the fact that he wasabout to become a father.) Though highly enthu-siastic about his work, Carothers was always shyand withdrawn, and in Gordon’s words, “he hadfew outlets other than work.” He was, however, atalented singer, as was his closest sibling, Isobel, aradio celebrity. Her death in January 1937 senthim into a bout of depression, and on April 29,he killed himself with a dose of cyanide. Sevenmonths later, on November 27, Helen gave birthto a daughter, Jane.

HOW PLASTICS HAVE EN-

HANCED LIFE. Despite his tragic end,Carothers had brought much good to the worldby sparking enormous interest in polymerresearch and plastics. Over the years that fol-lowed, polymer chemists developed numerousproducts that had applications in a wide varietyof areas. Some, such as polyester—a copolymerof terephthalic acid and ethylene—seemed to fitthe idea of “plastics” as ugly, inauthentic, andeven dehumanizing. During the 1970s, clothes ofpolyester became fashionable, but by the early1980s, there was a public backlash against syn-thetics, and in favor of natural materials.

Yet even as the public rejected synthetic fab-rics for everyday wear, Gore-tex and other syn-thetics became popular for outdoor and workoutclothing. At the same time, the polyester thatmany regarded as grotesque when used in cloth-ing was applied in making safer beverage bottles.The American Plastics Council dramatized thisin a 1990s commercial that showed a few secondsin the life of a mother. Her child takes a soft-



Polymers


drink bottle out of the refrigerator and drops it,

and the mother cringes at what she thinks she is

about to see next: glass shattering around her

child. But she is remembering the way things

were when she was a child, when soft drinks still

came in glass bottles: instead, the plastic bottle

bounces harmlessly.

Of course, such dramatizations may seem abit self-serving to critics of plastic, but the factremains that plastics enhance—and in somecases even preserve—life. Kevlar, for instance,enhances life when it is used in making canoesfor recreation; when used to make a bulletproofvest, it can save the life of a law-enforcement offi-cer. Mylar, a form of polyester, enhances life

ADDITION POLYMERIZATION: A

form of polymerization in which

monomers having at least one double bond

or triple bond simply add to one another,

forming a polymer and no other products.

Compare to condensation polymerization.

ALKENES: Hydrocarbons that contain

double bonds.

COPOLYMER: A polymer composed of

more than one type of monomer.

DIMER: A molecule formed by the

joining of two monomers.



valence electrons. Carbon is noted for its

ability to form double bonds, as for

instance in many hydrocarbons.

FUNCTIONAL GROUPS: An atom or

group of atoms whose presence identifies a

specific family of compounds. When

combined with hydrocarbons, various

functional groups form hydrocarbon

derivatives.

HOMOPOLYMER: A polymer that

consists of only one type of monomer.




HYDROCARBON DERIVATIVES:

Families of compounds formed by the

joining of hydrocarbons with various func-

tional groups.

MONOMERS: Small, individual sub-

units, often built of hydrocarbons, that join

together to form polymers.

ORGANIC: A term referring to any

compound that contains carbon, except

for oxides such as carbon dioxide, or car-

bonates such as calcium carbonate (i.e.,

limestone).

ORGANIC CHEMISTRY: The study

of carbon, its compounds, and their

properties.

PLASTICS: Materials, usually organic,

that can be caused to flow under certain

conditions of heat and pressure, and thus

to assume a desired shape when the pres-

sure and temperature conditions are with-

drawn. Plastics are usually made up of

polymers.

POLYMERIZATION: The process

whereby monomers join to form polymers.

POLYMERS: Large, typically chain-like

molecules composed of numerous smaller,

repeating units known as monomers.





KEY TERMS


Polymers when used to make a durable child’s balloon—but this highly nonreactive material also saveslives when it is applied to make replacementhuman blood vessels, or even replacement skinfor burn victims.

Recycling

As mentioned above, plastics—for all their bene-fits—do pose a genuine environmental threat,due to the fact that the polymers break downmuch more slowly than materials from livingorganisms. Hence the need not only to developbiodegradable plastics, but also to work on moreeffective means of recycling.

One of the challenges in the recycling arenais the fact that plastics come in a variety ofgrades. Different catalysts are used to make poly-mers that possess different properties, with vary-ing sizes of molecules, and in chains that may belinear, branched, or cross-linked. Long chains of10,000 or more monomers can be packed closelyto form a hard, tough plastic known as high-density polyethylene or HDPE, used for bottlescontaining milk, soft drinks, liquid soap, andother products. On the other hand, shorter,branched chains of about 500 ethylenemonomers each produce a much less dense plas-tic, low-density polyethylene or LDPE. This isused for plastic food or garment bags, spray bot-tles, and so forth. There are other grades of plas-tic as well.

In some forms of recycling, plastics of allvarieties are melted down together to yield acheap, low-grade product known as “plastic lum-ber,” used in materials such as landscaping tim-bers, or in making park benches. In order toachieve higher-grade recycled plastics, the mate-rials need to be separated, and to facilitate this,recycling codes have been developed. Many plas-tic materials sold today are stamped with a recy-cling code number between 1 and 6, identifyingspecific varieties of plastic. These can be melted

or ground according to type at recycling centers,and reprocessed to make more plastics of thesame grade.

To meet the environmental challenges posedby plastics, polymer chemists continue toresearch new methods of recycling, and of usingrecycled plastic. One impediment to recycling,however, is the fact that most state and local gov-ernments do not make it convenient, for instanceby arranging trash pickup for items that havebeen separated into plastic, paper, and glassproducts. Though ideally private recycling cen-ters would be preferable to government-operatedrecycling, few private companies have the finan-cial resources to make recycling of plastics andother materials practical.

WHERE TO LEARN MORE

Bortz, Alfred B. Superstuff!: Materials That Have ChangedOur Lives. New York: Franklin Watts, 1990.


Galas, Judith C. Plastics: Molding the Past, Shaping theFuture. San Diego, CA: Lucent Books, 1995.

Gordon, John Steele. “Plastics, Chance, and the PreparedMind.” American Heritage, July-August 1998, p. 18.

Mebane, Robert C. and Thomas R. Rybolt. Plastics andPolymers. Illustrated by Anni Matsick. New York:Twenty-First Century Books, 1995.

Plastics.com (Web site). <http://www.plastics.com> (June5, 2001).

“Plastics 101.” Plastics Resource (Web site).<http://www2.plasticsresource.com/plastics_101/>(June 5, 2001).

“Polymers.” University of Illinois, Urbana/Champaign,Materials Science and Technology Department (Website). <http://matse1.mse.uiuc.edu/~tw/polymers/polymers.html> (June 5, 2001).

“Polymers and Liquid Crystals.” Case Western ReserveUniversity (Web site). <http://abalone.cwru.edu/>(June 5, 2001).


