Thursday, January 10, 2013

Science

Chemistry v2i. Semester 1: Introduction

Chapter 1 Lesson 1: Measurement in Chemistry

Chapter 2 Lesson 2: Calculations and Graphing in Chemistry

Chapter 3 Lesson 3: Energy and Matter

Chapter 4 Lesson 4: Atomic Structure

Chapter 5 Lesson 5: Atomic Structure

Chapter 6 Lesson 6: Ionic and Covalent Compounds

Chapter 7 Lesson 7: Chemical Equations

Chapter 8 Lesson 8: Chemical Reactions

Chapter 9 Lesson 9: Calculations With the Mole

Chapter 10 Lesson 10: The Mathematics of Chemical Equations

Chapter 11 Lesson 11: Percent Yield and Review of Stoichiometry

Chapter 12 Lesson 12: Energy Changes in Chemistry

Chapter 13 Lesson 13: Radiant Energy and Quantum Theory

Chapter 14 Lesson 14: Electron Configuration

Chapter 15 Lesson 15: Periodic Trends

Chapter 16 Lesson 16: Designing an Experiment to Solve a Problem/Transition Metals

Chapter 17 Lesson 17: The Periodic Table of Elements

ii. Semester 2: Introduction

Chapter 18 Lesson 18: Kinetic Molecular Model of Gases

Chapter 19 Lesson 19: Gas Laws

Chapter 20 Lesson 20: Chemical and Physical Properties of Matter

Chapter 21 Lesson 21: Solids and Changing Phases

Chapter 22 Lesson 22: Review of Phase Change and Introduction to Solutions

Chapter 23 Lesson 23: Solutions

Chapter 24 Lesson 24: Colligative Properties of Solutions

Chapter 25 Lesson 25: Chemical Equilibrium

Chapter 26 Lesson 26: Le Chatelier's Principle

Chapter 27 Lesson 27: Solubility Equilibrium

Chapter 28 Lesson 28: Acids, Bases, and Salts

Chapter 29 Lesson 29: The pH Scale

Chapter 30 Lesson 30: Redox Reactions

Chapter 31 Lesson 31: Carbon and Its Compounds

Chapter 32 Lesson 32: Classes of Organic Compounds

Chapter 33 Lesson 33: Applications of Biochemistry

Chapter 34 Lesson 34: Review Guide

i.

Semester 1: Introduction

Semester 1 IntroductionWhat in the world isn't chemistry?

Chemistry is a branch of physics that explains how atoms and moleculesinteract. A complete understanding of chemistry will help you understanda host of interesting things: What causes color? How does your bodywork? Why does unpainted iron rust? Why is unpainted aluminumimpervious? How do detergents work? How do lasers and semiconductorswork? Why is water so special? How do the acids in vitamin C, in your carbattery, and in your stomach differ? What makes things burn? Whatstops the fire? You eat, breathe, look at, or work with chemicals everysecond of every day.

In this introductory chemistry course, you will study many principles thathelp explain the world around you. Topics for the first semester includemeasurement, energy and matter, atomic structure, chemical formulasand bonding, chemical reactions, the mole, chemical equations, heat inchemical reactions, electron configurations, the periodic table, andgroups of elements. The lessons point out many real-life applications andexamples of these concepts. In addition, your study of chemistry will bereinforced through a number of laboratory activities.

What will you need for this course?

Textbook: You will not need a textbook for this course. All pertinentinformation will be in each lesson.

Calculator: You will need a calculator that has scientific notation.

Periodic Table: A periodic table is provided for you to print out. Youwill need a copy of the periodic table available when working onsubmissions and exams.

Teacher/Mentor: The laboratory activities are an important part of thiscourse. Many of the activities require assembly of an apparatus orhandling of chemicals. Therefore, it is recommended that you find anadult to help you obtain materials, supervise activities, and aid you whennecessary.

Materials: For the course, the student will need a notebook. It isimportant that you take notes on the lessons. Most of the materialsyou will need for the labs are easily obtained, common items andfrequently, a substitute item will work if you don't have access to theappropriate lab equipment. Ask you teacher or mentor for help withanything you can't find on your own. The following is a list of thenecessary laboratory materials:

l balance accurate to 10 mg (0.01 g) and mass setl penny minted since 1980l nickell dimel quarterl waterl graduated plastic dropping pipet (a medicine dropper may be used)l small cup or plastic containerl 10-mL graduated cylinder

l table saltl sandl iron filingsl sawdustl clear gelatin (Knox gelatin)l balloonsl several 10 cm pieces of scotch tapel threadl popcornl toothpicksl yellow gumdropsl black gumdropsl 1 roll each of three different brands of bathroom tissuel other materials to be determined by the student

Prerequisite: It is recommended that you have a basic knowledge ofalgebra before you attempt this course.

Chapter 1

Lesson 1: Measurement in Chemistry

OBJECTIVES

l Define chemistry and consider the occupations that use chemistryl Examine the scientific methodl Review scientific notationl Convert numbers to and from scientific notationl Understand the rules of significant digitsl Solve problems that involve densityl Take time to consider what is chemistry?l Lab Activity: Massing Small Objects

Safety First and Always

Safety is important in all activities, but the chemistry laboratory is an area where safety must come first.Always have an adult present to supervise laboratory activities; never work alone when dealing withhazardous materials. Familiarize yourself in advance with the hazards found in each experiment, so that youwill know the precautions to take with each chemical you will use in an experiment. Know the location andproper use of all safety equipment. Use common sense and ask a teacher or your mentor for assistance if youare in doubt about the safety of any procedure. Always notify an adult of any problems you encounter.

Some specific rules will help you remember how to proceed:

l Choose a safe place to work.l Clear your work area of unnecessary materials such as books and papers before you start. Never taste

laboratory materials: no food, drinks, or gum are allowed when working with chemicals.l Wear safety goggles at all times.l Follow the instructions for each experiment; do not perform unauthorized experiments.l Tie back long hair. Avoid wearing draping sleeves or other clothing or jewelry that could catch on

equipment.l Familiarize yourself with equipment in advance. Always know how equipment works before beginning an

experiment.l Avoid awkward transfers. Pour from large containers into wide-mouthed beakers or containers, not test

tubes or other small-necked flasks.l Carry chemicals defensively: be sure there are no obstacles in your path and that you have a suitable

surface on which to place your container.l Glass and metal stay hot longer than they look hot. Let all hot containers cool before handling. Use hot

pads to protect counters and surfaces where needed.l Label your chemicals with the substance name, the date, and your name. The substance name must be

complete and legible.l Wash your hands thoroughly after each experiment.l Lock chemicals away out of reach of children.l Clean up after each experiment to minimize hazards.l Treat all chemicals with respect and caution to minimize hazards and maximize your fun in lab.

Safety Equipment

l Fire extinguishers should be used according to instructions. You will need to read the procedures on yourfire extinguisher in advance so that if an emergency does arise, you will be prepared. Fire extinguishersare found in almost all chemistry laboratories. Your local fire department or fire marshals office canprovide more information.

l Fire blankets are designed to be mounted vertically on a wall in a laboratory. A fire blanket can help tosmother flames that may be on your back or other areas that are difficult to reach. To use a fire blanket,grab the handle and pull away from the wall as you roll your body towards the wall.

l Laboratory aprons and coats are a useful layer of protection over your clothing. They serve to minimizeexposure to hazardous chemicals that may spill on you. They should be easy to remove. Protectiveclothing should be made of flame retardant materials that are tightly woven to prevent seepage of liquidsonto your clothing.

l Safety showers are designed to provide a large volume of water to rinse your entire body quickly in the

event of a large spill or splash. They are particularly useful if spills or splashes cause your head and bodyto be covered in hazardous substances. A safety shower should provide a large volume of water andshould rinse you from the top down, just as your shower at home does.

l Eye Wash stations provide gently running water that you get close to, so that your eyes can be rinsedwith water for 15 minutes. Eye wash stations often resemble your kitchen or bathroom sink.

l Cardboard boxes are kept nearby for safe disposal of sharp items, such as broken glass. Trash bags andother soft containers are easily punctured by broken glass or sharp metal points or edges. Be sure thatyou have a disposable container of cardboard, thick plastic, or other material that is safe from punctureby broken glass.

l First aid kits may be actual boxed kits or simply your collection of first aid supplies that are assembledand ready to use in an easily accessible location. You should include band-aids, adhesive tape, headacheremedies, sterile eye wash, burn cream for heat burns, and other useful items such as tweezers, sterilegauze, small scissors.

What is Chemistry?

Chemistry is often called the central science because it is the study of all substances and the changes theyundergo. Chemistry is going on all around us and in our bodies all the time. The processes that help usbreathe, think, laugh, and eat all involve chemistry. The processes that allow us to have automobiles andtelephones, space shuttles and televisions, all involve chemical processes. Without chemistry we would nothave foods packaged and preserved safely for storage on convenient supermarket shelves or for shipping tolocations far from where the food was produced. The wide variety we enjoy in our diets, the safety andconvenience of home refrigerators, and many other items are available to us because of our understandingand application of chemistry. We could not exist without chemistry.

So, what does this mean to you when you meet someone who is a chemist? The chemist typically has abachelors degree in chemistry or a closely-related field, or perhaps a Masters degree or Ph.D. in chemistry orchemical engineering. Sometimes the job title chemist does not necessarily mean specialized degrees, butmay refer to job functions that include monitoring chemical mixing in an industrial setting.

Jobs for chemists most often include several activities: research, analysis, writing scientific papers and othercommunications, monitoring a process, or changing a process to make it better. Research can take manyforms, but it usually refers to the investigation of new ideas or of applying known ideas in new ways.Analysis often means using some type of instrument to measure the physical and chemical properties ofsubstances. Writing and communicating ideas to other people is an important function of chemists and otherscientists that people sometimes overlook. It can be fun and challenging to try to explain a complicated newidea in a way that gets everyone excited about it!

Many occupations require some knowledge of chemistry even though the job title might not be chemist."Biologists must know the chemical processes at work in cells to understand the functioning of livingorganisms. A wetlands ecologist is a type of biologist dealing with all types of creatures found in or nearbodies of water. The substances in the water and their possible toxic effects on wildlife are of particularconcern to the wetlands ecologist. Archeologists must understand how materials are preserved, how todetermine the age or materials, and how different materials change over time. Firefighters must know theproper use of fire retardant materials as well as the type of fire they are fighting so that the appropriateextinguishing substance, water or chemical foam, may be used. Engineers must understand the properties ofthe materials they use to fabricate bridges, roads, buildings, and intricate electronic devices. A hairstylistmust understand the properties of hair care products designed to clean, condition, color, and style hair, andall of these substances must be handled appropriately to insure safe and happy customers. A permanentthat only last a few days isnt very permanent!

Those are just a few of the many occupations that require a special knowledge of chemistry. For example, it isnot obvious to others that chemistry is important for candy manufacturers who must know how to effectivelyhandle sugar solutions of varying concentrations and densities to produce the best treats. A chef must knowthe importance of proper cooking times and temperatures to insure that raw food is cooked safely. Chemicalchanges that occur during cooking are important both for proper taste and safe food. Photographers oftendevelop their own photographs, which requires a knowledge of silver bromide and other developing materials.A mortician must understand the properties and usage of embalming fluids. A swimming pool maintenanceworker needs knowledge of acids and bases and pH balance. A pilot needs to understand changes inatmospheric pressure. A jeweler must understand properties of crystals and metals to achieve results that arepleasing to the eye, as well as suitable for different applications. An architect must understand the propertiesof different woods and metals so that they can be used in building applications. Laboratory technicians inmany fields (including medical labs) also need knowledge of chemistry, especially acids, bases, and

concentrations. So get ready, as we explore these topics and more.

Units of Measure

Units of measure allow us to make sense out of the numerical measurements we encounter. Standardmeasurements for length, mass, and volume are familiar to us. Both the traditional (English) system and theSI (systeme international) system of measurement have standard units for measuring these quantities. Wewill deal with the SI system here, which is most commonly used for scientific applications. The advantages ofthe SI system include the easy division of all quantities by ten or some power of ten (which also makesestimation easier and more practical) and world-wide acceptance of this system as a standard. Even GreatBritain, home of the English system, has recently begun converting their national measurement system to theSI system. Only the US still maintains the English system as a national standard.

Name of Standard Units

Units of MeasurePhysical Quantity Symbol

Mass kg kilogram

Length (distance) m meter

Volume L liter

Time s (or sec) second

Electric current A ampere

Temperature K Kelvin

Amount of Substance mol mole

Luminous intensity cd candela

Common PrefixesPrefix Symbol Explanation

kilo- 103 kilometer means 103 meters, whichequals 1,000 meters (1 km= 1,000 m)

centi- 10-2 centimeter means 10-2 meters, whichequals 0.01 meter (1 cm = .01 meterand 100 cm = 1 m)

milli- 10-3 millimeter means 10-3 meters (1mm =0.001 m and 1,000 mm = 1 m)

The prefixes are used to indicate a larger or smaller unit than the standard unit. For example, a kilometer is

103 times one meter, meaning that 1 km = 1000 m. A km is a LARGER unit than a meter. Lets look at anexample that indicates a smaller unit. A milliliter, mL, is a very common unit of measure; in fact, you mayhave noticed that most soda and juice bottles give the volumes in milliliters as well as ounces. A milliliter is

10-3 times one liter, meaning that 1000 milliliters = 1 liter. The liter is the larger unit in this case.

Now lets consider temperature. We know temperature to be the measure of how cold or hot something is.Temperature scales are typically based on some readily observable property or event. The boiling point(temperature at which a liquid vaporizes to a gas) and melting point (temperature at which a solid becomes aliquid) of water have been used as these reference points. A scale of temperature measurement usedprimarily in the United States is the Fahrenheit scale The Fahrenheit scale's reference points range from 32 to212 degrees. In brief, 32F is freezing and 212F is boiling.

Most measurements worldwide are in Celsius, the centigrade temperature scale, indicated by C. The Celsiussystem was devised by a an eighteenth century Swedish astronomer named Andres Celsius. The boiling pointof water on this scale is 100C, and the freezing point of water is 0C.

The third temperature scale we will consider is the Kelvin scale. This is called the absolute temperature scalesince it is not dependent on the observation of known properties but is instead based on the energyassociated with matter on an atomic level. The idea of absolute zero comes from the Kelvin temperature scale(0 K = absolute zero). Notice that there is no degree sign before the K symbol.

Precision and accuracy are terms that are often used to mean the same thing in everyday conversation, but,in fact, they are not really the same. Precision is the ability of a measuring instrument to give the samemeasurement over and over again. Precision is often called repeatability. Accuracy, on the other hand,refers to the closeness of a measurement to the accepted value. Accuracy refers to correctness of ameasurement

Keep in mind that all measurement involves a person reading a scale of some sort. The object beingmeasured may fall exactly on a specific measurement or may fall between specific readings on the scale. Theperson making the measurement determines exactly what reading is correct and often estimates themeasurement. The estimates that are made may be very good, but they always involve some degree ofuncertainty. The other factor that contributes to uncertainty in measurement is the physical limitations of themeasuring device. No device is absolutely perfect, though some devices may be extremely accurate andprecise. Thus, a measurement could be repeatable, but not correct (precise but not accurate). Or ameasurement could be correct once, but not repeatable (accurate but not precise); it might be neither ofthese or both.

It is useful to have a way to indicate how precise and how accurate measured values are for any quantity wemeasure. We use significant figures to help us do this. Significant figures in a measured number include allthe digits known for certain, plus one digit that is uncertain (the estimated digit would be the uncertain digitin a measurement). Even electronic instruments make an estimate in the last digit. The precision of anyinstrument is ultimately determined by how well the instrument is constructed, as well as how well theinstrument is maintained. To determine which digits are significant in a correctly written measurement, usethese rules and examples to guide you. Remember that all non-zero digits are significant; that is, theyprovide information for us to some degree of certainty

For instance, the measurement 5.2 cm has 2 significant digits, while 5.23 cm has 3 significant digits. Thedigits for which we need rules are the zeros. Zeros are sometimes significant, and sometimes are not. Zerosbetween two nonzero numbers are significant. In the measurement 3.005, there are 4 significant figures.Zeros at the end of a number are significant if they are to the right of a decimal point. In 7.0, there are 2significant figures (the zero is significant; it would not be written if it were not). Zeros to the left are notsignificant, as in 0.045 (there are 2 significant figures here). There is one situation in which it is difficult todecide if a zero is significant; if a zero is to the right of nonzero numbers, but there is no decimal. In 3500,there is nothing to indicate if one or both of the zeros is significant or not; there may be 2, 3 or 4 significantfigures. If both are meant to be significant, the number could be written with a decimal at the end (3500.).The decimal would then indicate that both zeros are significant. This is a bit awkward, so we typically use

scientific notation to correctly express significant figures for these types of cases: 4.200 x 103 would correctly

show 4 significant figures, 4.20 x 103 would show 3, and 4.2 x 103 would clearly indicate only 2 significantfigures.

Rules for Calculations Involving Significant Figures:

For addition and subtraction, the answer should have the same number of digits to the right of the decimalplace as are found in the number having the fewest digits to the right of the decimal place. For example,when adding 56.01234 g and 4.29 g, your answer will be correctly expressed as 60.30 g.

DISCUSSION

Chemistry affects our lives in so many ways. You have to understand chemistry to really understand howlasers work, why cells produce CO

2, or why detergents clean your clothes. What in the world isn't chemistry?

Lab Activity: Massing Small Objects

Mass is measured in the laboratory with a balance. To complete this lab, you will need a balance and massset. Consult your teacher for this equipment. You may be worried that the balance included in your kit is notas accurate as more modern-looking electronic balances. Actually, it should have the same precision as theelectronic balances which are used in most classrooms. Your balance will be precise enough to find the

difference in the mass of two quarters.

In order to take advantage of the accuracy of your balance, special procedures must be used. You will use thebalance often to determine the mass of chemicals involved in chemical reactions. Learn how to use thebalance properly and your laboratory results will be much better (which will bear upon your grade in theclass).

Massing an Object

1. Always use some type of container to hold the chemicals you will be massing. Examples: milk bottlecaps, short Styrofoam cups, etc. Remember that the container will have some mass of its own.

2. Place the object to be massed on the left pan of the balance.3. Add and/or remove masses, usually starting with the larger and moving to the smaller, to and from the

pan on the right until the pans balance.4. For greatest accuracy, scientists never touch any of the weights with their fingers. This would leave

deposits from fingerprints on the weights. Handle the weights with the forceps included in your weightset. Be careful not to allow any chemicals to contact the weights or the pans of the balance.

5. The units of the masses which are made of brass are in grams. The units of the small aluminum massesare in mg.

Examples:

a 10 g mass and a 100 mg mass = 10.100 g

a 10 g mass and two 100 mg masses = 10.200 g

a 10 g mass and a 10 mg mass = 10.010 g

6. Find the mass of a small empty container. You can not assume that two containers have the same massif they are the same. Two Styrofoam cups from the same package will usually differ in mass by morethan 0.05 g.

7. Add the chemical to be massed out.8. Find the mass of the container and its contents.9. Subtract the mass of the container to find the mass of its contents.

10. When writing the mass of the object, always include the units as part of the answer.

Chapter 2

Lesson 2: Calculations and Graphing in Chemistry

OBJECTIVES

l Understand how to solve problems dealing with densityl Examine the algebra needed to solve chemistry problemsl Understand that units must be treated as factors in chemistry problems and solve problems using these unit factorsl Define and recognize unit multipliers and compile a list of unit multipliers that can be used to make conversionsl Work problems using unit multipliers to make conversions.l Learn how to measure the mass and volume of waterl Graph the results of a lab and use the graph to determine the density of a solutionl Determine the percent error in a problem

DISCUSSION Density

Density is the ratio of mass over volume. Density is an important characteristic for identifying substances.When you see a rock, you visually appraise its size and estimate its mass. You have estimated its density. Ifyou pick the rock up and find it is five times heavier than you anticipated you say, "This must be a kind of rockthat is different from usual." If you pick up a piece of metal and notice that it is gray colored and has a highmass to volume ratio, you might guess that the metal is lead because lead is the most common metal with ahigh density.

There are many kinds of ratios: distance/time (speed), price/ pound (unit price), individuals/acre (populationdensity), etc. All of these things can be expressed in a simple equation. If you learn how to use theseequations, you will have a powerful tool to solve problems.

If you know the density and volume of an object, you can calculate its mass:

When you solve the problem for volume, the equation becomes:

Lab Activity Hints

Density

Density is a characteristic property of a pure substance. Density is the ratio of mass over volume, D = m/V.The concept of density is one we encounter everyday. When we get a glass of water to drink, if we add icecubes (solid water), the ice cubes float. The ice cubes are less dense than the liquid water in which they float.If you go to a gas station after a rainfall, you may see puddles with swirling patterns on top the gasoline thatwashes into the puddles is less dense than the water, so the gasoline floats. If you are preparing a cake, therecipe may call for oil as well as water. When the oil and the water are placed in the same measuring cup, weobserve the oil floating on the water; we can clearly observe 2 layers. In this case the water is more densethan the oil. On a molecular level, density refers to the arrangement of the atoms and molecules of asubstance, the "packing." Substances whose atoms or molecules are packed more closely together are denser.In general, solids are more dense than liquids, which are more dense than gases. Water solid (ice) is animportant exception to this rule, and we will discuss this in more detail when we discuss intermolecular forcesin a later chapter.

DISCUSSION What a Concept! 1=1!

You know that one times a number is still that same number.

Unit multipliers may not look like one times a number, but they are. It is often helpful to reduce the problemto the simplest examples. Let's look at some simple examples.

Because these ratios are all equal to one, they are called unit multipliers. You can multiply a number by aunit multipliers and still have the same number.

The real power of unit multipliers is converting from one unit to another. Combined with what we have learnedabout working with units in the previous section, this concept becomes a very useful process.

Unit Conversions You Should Know

Below is a partial list of unit multipliers most commonly used during the course. You may wish to add to thislist or write the list where you can refer to it easily.

This not an exhaustive list. Many of the unit multipliers are just a matter of learning the metric prefixes.

Unit Multipliers

Units as Factors in Chemistry Problems

Units have an important function: the unit specifies what is measured. For instance, if a sign on the interstateindicated that the next exit is 5.5 away, how far would it be? Maybe 5.5 miles, or 5.5 kilometers, both commondistance units, or maybe something else! That may not seem too important, until you consider the fact that a5 mile run is about 1.6 times longer than a 5K run. So units do matter!

We use units in our calculations, so we must know how they work. Units add, subtract, multiply, and divide,according to the examples provided. You are probably familiar with this concept from math class; if not, theseexamples will help.

Use these sample problems to help you answer the questions.

ex 1. Units in division:a. 55.2 g/125 mL = 0.442 g/mLb. 10 gallons x 35 miles/gallon = 350 miles

ex 2. Units in multiplication:

c. 12.0 cm x 10.5 cm x 6.7 cm = 844 cm3 (cubic centimeters)

d. 4.5 ft x 5.3 ft = 24 ft2 ("square feet")

ex 3. Example word problems:e. A sample of metal has dimensions length =7.2 cm, width = 1.5 cm, and height = 2.4 cm. What volume of

space does the metal occupy? (A: 26 cm3)f. The same sample of metal weighs 2.95 g, and density = mass/volume. What is the density of the metal?(Hint: use the volume you calculated in the previous example.)

Answer: D = 2.95 g/ 26 cm3 = 0.11 g/cm3

Scientists use a wide variety of abbreviations and symbols to stand for units. No matter what the unit is, thesame rules of unit manipulation apply.

Here are more examples of this process:

ex 4.

ex 5.

The same approach to units that we have already used also works when we are converting between differenttypes of units, even in different systems of measurement. Following this approach helps us keep track of theunits we have in our starting quantity and the units we want in our answer.

ex 6. How many gallons are in 46.0 L? Look for a conversion factor that relates gallons and liters in the tableprovided.

To solve our conversion problem, we need the factor 1.0 gal = 3.785 L. Begin with the information we have,our measured volume, 46 L, and the information we want (its equivalent in gallons):

and

We arranged our conversion factor in a specific manner: the units we wanted in our answer remained on top(in the numerator), while the units in our original number were in the denominator. We multiplied 46.0 by1.00, then divided by 3.785.

DISCUSSION Why Make Graphs?

Long tables of data are tedious and often hard to follow. Not only that, it is difficult see trends in the tables'data that may be important to research. There are other ways to display data, however. For example, a groupof research assistants studying nuclear energy might compile all their data and graph the information using acomputer spreadsheet. The data available on one graph would have taken 20 pages of tables. The graph wouldalso show the data in a clear and more useful manner, making the researcher's analysis much easier.(Learning how to make these graphs using a computer spreadsheet would be very beneficial. Consult yourteacher to determine if this option is available to you.)

Problem Solution

(10)(x)(x)(y) = x

(x)(y)(y)(10) y

Problem Solution

(25)(x)(z)(z) x=

xz

(x)(y)(z) 25 y

Conversion Factors

Mass Length Volume

2.205 lb. = 1.000 kg 1.094 yd. = 1.000 m 1.0 quart = 0.946 L

1.000 lb. = 453.6 g 0.3937 in. = 1.000 cm 1000 ft3 = 28.3 L

1.000 oz. = 28.35 g 0.6214 mile = 1.000 km 1.00 gal = 3.785 L

Havemultiply by the correct conversion factor

Want

46.0 L X gallons

46.0 L x1.00 gal

= 12.2 gallons3.785 L

Graphing

Graphing is a technique we use to arrange information in a visual manner. We can observe trends and drawconclusions from graphical information quickly and easily. Graphs can also help us to predict future valuesonce we have established a pattern. Many graphs are depicted in line form, because linear graphs are moreeasily used to extrapolate (continue a set of data along a trend), interpolate (find a value between two datapoints), and detect unreliable data. The slope of a line graph is readily obtained and often provides usefulinformation.

Here is an example showing how to find the slope of line.

Finding the slope of a line

1. Locate two points on the graph as shown.

2. Draw in the line as indicated.

3. Determine the x and y values for each.

4. The slope is often referred to as rise over run, or the change in the y-axis between those two points dividedby the change in the x-axis between those two points.

5. In this case, notice that the ratio of rise over run has units of distance over time. The slope of this line isthe speed the object was going in km per minute

To graph a set of data, the information must first be collected by the experimenter. If we wish to measure howlong it takes boiling water to cool to room temperature, we record the temperature of the water at differenttime intervals after first boiling the water. Since the experimenter controlled time by choosing when toobserve the temperature, time is the independent variable. The independent variable should be shown onthe horizontal axis of the graph, the x-axis. Temperature is the dependent variable, since thetemperature observed by the experimenter depends on the time at which it was observed. The dependentvariable should be shown on the vertical axis of the graph, the y-axis.

Next, we must decide how many units we need on each axis of our graph. We want our graph to be easy todraw and read. The best graphs also fill more than half the available graph paper in each direction. For theindependent variable on the x-axis, if we had 9 times at which we made measurements, we could separatethem each along one ninth of the horizontal space of the graph paper. Alternatively, we might make graphinga bit easier and mark 10 divisions on our x-axis, even though we need only 9. This makes our graphingsimpler, which can help us reduce mistakes. The same principle would apply to any set of measurements;using a total number of divisions that can be divided by 5 or 10 is often easier to graph. Having more divisionsthan absolutely necessary also allows space for extrapolation. For our dependent variable, we use the same

process. We would also have 9 temperature readings that we observed, one at each time interval. Here is anexample graph of our time and temperature data:

This graph can be generated using a pencil and graph paper, a spreadsheet such as MS Excel, or a graphingcalculator such as a TI-83.

Percent Error

Percent error is a number that is easy to calculate yet has an important meaning. Percent error refers to thepercent difference between the correct answer and an answer obtained by measurement or theoreticalcalculation. Most often percent error is used to measure the accuracy of experimental measurements. Recallthat accuracy refers to closeness to true, correct value. So, a small percent error refers to a very accuratemeasurement, and a high percent error refers to a measurement that is not very accurate. To calculatepercent error, you need to know the correct value for a measurement (in some situations this is simply theaccepted value); you must also know your measured value in the same units as the correct value. To expresspercent error, use this equation:

ex 7. Measured value of a length: 9.88 m Correct value: 9.55 m

Percent error (measured compared to correct):

Remember that if a percent error calculation results in a negative number, it does not mean that you haveless than zero error; a negative percent error indicates that your measured value is simply lower than thecorrect value by the calculated percent.

Table of Collected Data

Time (min) Temp (C)

0 99.5

5 92.5

10 85

15 78

20 71

25 65

30 57

35 50

40 40

Measured value correct valuex 100 = % error

Correct value

9.88 m - 9.55 mx 100 = 3.45% error

9.55 m

Chapter 3

Lesson 3: Energy and Matter

OBJECTIVES

l Name and describe the three basic forms of energy: potential, kinetic, and electromagnetic.l State and apply the law of conservation of energyl Differentiate between the terms temperature and heatl Compare the Fahrenheit, Celsius, and Kelvin temperature scalesl Convert temperatures from Celsius to Kelvin temperature scalesl Explain what is meant by absolute zerol Name and describe the four states of matterl Compare physical and chemical properties of matterl Define and recognize examples of elements.l Learn the correct way to write the symbol of an elementl Define and recognize examples of compounds

Energy

Energy is the ability to do work to accomplish change. Energy is found in many forms: light, heat, sound,mechanical, electrical. Energy can be classified into three basic forms. Potential energy refers to storedenergy. Potential energy can be stored in the form of chemical bonds (chemical potential energy) or asenergy of position (e.g., a rock at the top of a hill).

Kinetic energy is the second basic form of energy. Kinetic energy is the energy possessed by an object as aresult of its motion (e.g., a baseball traveling through the air towards home plate).

The third major form of energy is electromagnetic energy, also known as radiant energy. This refers toenergy traveling in waves, in a straight line, at the speed of light. This type of energy includes the visible lightof all colors that we can observe with our own eyes. Electromagnetic energy includes the ultraviolet energy(UV rays) that can give us sunburn. Electromagnetic energy also includes the infrared energy that radiatesfrom a conventional oven, when we open the door to check our food. We do not see infrared energy, but wecan definitely feel it on our hands and face.

There are specific patterns of behavior of different forms of energy that have been observed and organized into overarching principles calledlaws. The Second Law of Thermodynamics states that energy is neither created nor destroyed in ordinary chemical reactions. This lawmeans that the total amount of energy remains the same, but that energy can be transferred from one place to another, or transformed fromone form to another. An example of transferring energy is placing a pan of water on the stove top for several minutes. The infrared energyfrom the burners heats first the pan and then the water the heat energy is transferred from an area of high amounts of heat energy (theburner) to an area of relatively low amount of heat energy (the water). An example of transforming energy from one form to another isburning fuel in an automobile. The combustion reaction allows the energy stored in the chemical bonds of the fuel, to be released as heatenergy during the course of the reaction. Operating a car involves a series of transformations and transfers. For example, burning gas todrive a car involves converting chemical energy to thermal energy to mechanical energy.

Combustion processes are the most common way we have to release chemical energy in a useful form, heat.Heat is defined as the amount of energy in a system. Temperature is a measure of the average kineticenergy of a system (or the average energy of the molecular motions of a system). Basically, each moleculehas its own kinetic energy. The average kinetic energy is a measure of the average kinetic energy of all themolecules in a substance. Electrical power plants use a series of transformations to convert coal, oil, or waterpower (hydropower) into electricity. The energy is stored as potential chemical energy in the coal and oil and,for hydropower, as energy of position in the water at the top of a fall.

Photosynthesis and Energy Storage

Energy can also be converted from its electromagnetic form into storage. Photosynthesis is an example ofthis type of storage process. Ultraviolet light from the sun provides the energy plants need to convert carbondioxide (CO2) and water (H2O) into glucose (C6H12O6), which the plant can use for energy, or as a building

material for plant structures. The second product of this reaction is breathable oxygen gas, a real bonus for us!

Energy + CO2

+ H2O C

6H

12O

6+ O

2

Sunlight energy is really not just the light we see, but also the infrared heat energy we feel beating down on

our skin, the ultraviolet rays that can tan or burn and other electromagnetic (radiant) energy. Fossil fuels suchas coal, petroleum (oil), and natural gas all originate in photosynthesis. These fuels begin as plant and animalmaterials composed of the elements carbon (C), hydrogen (H), and oxygen (O). After millions of years ofgeological changes in the earths crust, these materials would be buried by tons of rock and dirt. Under therock and dirt, sunlight and oxygen from the atmosphere were not available. Over time, the extreme conditionsof the rock pressing down on the materials, combined with the extreme heat from the earths core, caused theplant and animal materials to react, releasing water and oxygen. Thus, the fuels that we know today, coal,petroleum, natural gas, all contain carbon and hydrogen. These are known as fossil fuels because they aregenerated from fossils. Some molecules found in crude petroleum also contain a small proportion of oxygen,but it is very small compared to the amount of carbon and hydrogen. The high relative percentage of carbonand hydrogen in substances we typically regard as fuels, is one of their defining characteristics.

The processes by which fossil fuels are formed present a problem. These processes occur so slowly that theamount of fuel that any person requires during their lifetime cannot be regenerated during that lifetime. Infact, many highly industrialized nations such as the United States, Germany, Great Britain, are currently usingfuels many times more quickly than they can be regenerated. You may even be aware of the newspaperarticles and reports that reveal the serious concern many people have over the amount of fossil fuels that areavailable to us. Since fossil fuels cannot be regenerated by natural processes on a useful human time scale, wecan not count on more of these fuels than we already have. Therefore, these resources are considerednonrenewable. A resource that cannot be replenished will eventually run out.

">That is the concern many people have about our rapid rate of use of fossil fuels. Some projections indicate we may run out by the year2050. Some estimates indicate that if we conserve resources we may be able to extend out supply for another 100 years or more. In anyevent, there is still a strong possibility that we may have to turn to other sources of energy in the future. Many countries and many privategroups are already exploring the development of other energy sources (called alternate energy sources). We will discuss some these in alater chapter.

Another concern associated with fossil fuels involves the products of the combustion reactions that release thestored energy in fuels. For example, the combustion of methane (CH

4, the hydrocarbon fuel we know as

natural gas) in air produces carbon dioxide and water according to this balanced equation:

CH4 + 2 O2 CO2 + 2 H2O + Energy

Carbon dioxide is one of the major greenhouse gases. A greenhouse gas is a substance containing two ormore elements that absorbs infrared energy inside the earths atmosphere and radiates that energy back toearth. The gases in our atmosphere behave like the walls of a greenhouse, trapping in heat to keep thetemperature suitable for supporting plant and animal life. It is very beneficial that our atmosphere functions asa greenhouse. Without an atmosphere, the earth would be socold that we could not survive.

The problem with greenhouse gases such as CO2, and combustion reactions that we use to release energy, is

that we are producing CO2

more rapidly than the plants on earth can use it in photo-synthesis. There are also

other processes besides photosynthesis that use CO2. Even combining all the CO2 used up by natural

processes on earth, human activities are still contributing to a rapid increase in the amount of CO2

in our

atmosphere. More CO2

could mean more energy trapped inside the atmosphere. More energy could mean that

the earth would warm up enough to begin melting the polar ice caps and causing drastic changes in localclimates all over the world. There is much discussion over the controversy of what exactly will happen in thefuture, and to what extent human activities can affect the composition of our atmosphere.

You may want to examine more information to come to your own conclusions. Learning to examine theavailable information to draw your own conclusions is an important skill, but it takes practice!

Converting Temperatures: Absolute Zero and the Kelvin Scale

The absolute, or Kelvin, temperature scale is often used in chemistry, particularly when working with gaslaws. The Kelvin degree is the same size as the Celsius degree. Unlike the Fahrenheit and Celsius temperaturescales, which are based on the boiling and freezing points of water, the Kelvin scale is based on absolute zero(0K). Absolute zero is equivalent to -273.15C and is the lowest possible temperature. It is the point where allmolecular motion ceases, and where an ideal gas has a volume of zero. To convert from a temperature fromCelsius to Kelvin, just add 273.15 to the Celsius temperature.

What Is Matter?

Matter can be defined as anything that has mass and takes up space. That means everything! Keep in mindthat atoms are the smallest bits of matter that retain a complete set of physical and chemical properties. Eachatom has a specific arrangement of subatomic particles, which we will discuss in Chapter 4. We can identifywhich atoms are alike and which are different, by observing their properties, and by observing their structure.Atoms are then organized into categories, by structure. The categories are called elements. Keep in mind thatall atoms of an element have a complete set of physical and chemical properties that are characteristic of theelement. You may have heard about interesting newly discovered particles that are the result of breaking apartatoms or parts of atoms in high-energy collisions. While this is interesting to know, keep in mind that almostall of the changes we observe everyday, on an atomic (micro) level, or on a larger, bulk (macro) level, can beaccounted for by the properties of different atoms, which we can identify as specific elements.

Now, on to physical and chemical properties! Physical properties can be observed without changing thecomposition of a substance. Examples of physical properties are shape, color, and state of matter. We willexamine states of matter very soon. Chemical properties can only be observed when you are changing theidentity of a substance, through a chemical reaction. Reactivity is a chemical property, referring to howquickly and easily something reacts. For instance, not all materials are combustible (able to be burned as fuel).Combustion or burning reactions are common, and are familiar to us. Our cars use internal combustionengines and burn gasoline as fuel. When we cook out on a grill, we use charcoal as fuel.

States of matter refer to the physical arrangements of substances. There are three common states ofmatter: solid, liquid, gas. A fourth state of matter, plasma, is also known. Each state has specificcharacteristics associated with it. We will describe the characteristics of solids, liquids and gases, byconsidering a sample of a common substance, water, as our example in the following table:

CHARACTERISTICS OF STATES OF MATTERSolid Liquid Gas Plasma

Definite Shape No Definite Shape No Definite Shape No Definite Shape

DefiniteVolume

Definite Volume No DefiniteVolume

No Definite Volume

High EnergyParticles

Example Example Example Example

ice (solidwater)

water in the drinkable form with which we are mostfamiliar

steam the sun

Plasma is different from the other three states of matter. Plasma is a high-energy state that is not necessarily composed of one specificsubstance. Plasma is not commonly found on the surface of the earth. We can, however, generate small amounts of plasma in scientificinstruments under controlled conditions.

Lets examine our three common states of matter in more detail. Solids have a definite shape and volume. Anice cube holds it shape, and only takes up a certain amount of space. As the ice melts, we see that the liquidthat forms, water, no longer has a specific shape. The liquid water takes the shape of the container it is in, orspreads out in an indefinite shape it is on a surface such as a counter top. The water in an ice cube will onlyspread so far; it does have a definite volume. You cannot wet the entire kitchen floor with one ice cube!

When we boil liquid water in a pan on the stovetop, we observe a cloud above the pan. The liquid water has evaporated into gaseous water,steam. We can cool the gas cloud, and it will condense, or return to the liquid state. You have seen water condense, if you have ever boiledwater and noticed that the countertop nearby (which is cooler) has water drops on it. If you have a cool beverage in a glass, and waterappears on the outside, your cup is not leaking, you have simply condensed water from the air on the cool surface of the glass.

To melt the ice, we do nothing more than remove it from the freezer. The warmth of the air provides sufficientenergy to melt the ice to liquid. But to boil water, we need a source of energy, and a lot of it! Our stovetopburners or microwave oven can provide the energy.

To sum up what we know about the physical states of matter: the identity of the substance, water, does not change just because the ice meltsor the water boils. The water is still water.

We can relate what we know about states of matter (a physical property) to the larger topic of physical andchemical properties. Physical and chemical properties can be remembered best, when we use examples ofphysical and chemical changes to remind us.

Ex 1. Physical changes

Physical changes are often reversible, and do not change the identity of the substance. Chopped wood is still wood.

Ex 2. Chemical Changes

Boiling an egg is not easily reversible. In fact, boiling (cooking) an egg causes a chemical change in thesubstances in the egg, making cooked eggs more digestible for humans. Rusted iron is not good for buildinganymore; burned wood cannot be used as fuel again.

Elements and Compounds

Elements are organized on the Periodic Table by atomic structure, which we will discuss in more detail later.The Periodic Table also reveals the periodic law, which states that the physical and chemical properties of theelements are consequences of their atomic numbers. Elements are organized into groups (vertical columns)with similar physical and chemical properties. All atoms of the same element have the same complete set ofphysical and chemical properties belonging to that element. There are about 110 known elements. ManyPeriodic Tables do not show 110 elements, if they were printed or published before all the elements werediscovered. The first 40 elements are encountered most often. Many of the newly discovered elements are onlyknown to exist under laboratory conditions. Elements are identified by a name and a short version of thename, or symbol. The symbol may be one, two or three letters. Each element has a unique name and symbol.Carbon has symbol C, oxygen, O, calcium, Ca. The first letter of a symbol is always capitalized. The elementsymbol can represent one atom of the element or a group of atoms, depending on the context. Elements areclassified as pure substances, because each element has a unique set of physical and chemical properties.

A molecule is a group of two or more atoms that are chemically combined. The atoms can be the same, as inO

2, a molecule of oxygen gas. The atoms can also be of different elements; a substance that contains two or

more elements in fixed proportions is called a compound. Water (H2O), carbon dioxide (CO

2), and table salt

(sodium chloride, NaCl) are all compounds. The law of definite proportions states that elements in a puresubstance will always exist in the same proportions. Water is always two hydrogen atoms and one oxygenatom per molecule of water, H2O. Measured by mass, water is 11.1% H and 88.9% O. Hydrogen peroxide has

Chopping wood Boiling water Alcohol evaporating

Butter melting glass breaking Wire bending

Boiling an egg Iron rusting

Wood rotting Milk spoiling

the formula H2O

2, which is 5.9% H and 94.1% O. Water, H

2O, tastes good to drink and we must have it to

live. Hydrogen peroxide, H2O2, is a disinfectant, used to kill germs in cuts and scrapes. Do not drink hydrogen

peroxide!

DISCUSSION: A New Language

Here you start to look at the letters and words that chemists use to talk about chemistry. If you think it lookslike a different language, you are correct. There are specific, correct ways of expressing yourself in chemistry.

We start here with the basic letters of the language: the symbols for the elements. The chemist puts thesesymbols together into words and the words into sentences. A big part of learning chemistry is learning thislanguage.

The symbols for elements are written with the first letter capitalized and the remaining letters of the symbol inlowercase. Why is this important?

l Co is the symbol for cobalt, a metallic compound like iron with a slightly pink color in it.l CO is carbon monoxide, a colorless gas which attaches to your hemoglobin and prevents the passage of oxygen to your cells.

Notice that if you write the symbol for cobalt incorrectly with both letters capitalized, you would be sayingsomething very different.

LAB HINTS

Some elements are attracted to magnets; check around the house, particularly on the refrigerator! An elementattracted to a magnet could be separated from the other substances in the group. Some substances aresoluble in hot water, but not in cold. Try cold water to dissolve one substance, and set the water aside. Try aseparate sample of warm water (handle with care!) to try to dissolve something else from the mixture. Allowthe water to evaporate from the containers, to observe any solid that remains. Some solids do not dissolve inwater. A filter or screen might be used to remove them. Coffee filters and strainers allow some substances togo through, while trapping others.

Chapter 4

Lesson 4: Atomic Structure

OBJECTIVES

l Define the word atoml Describe the four points of Dalton's atomic theory of matterl Identify and describe the two kinds of electrical chargesl Describe how particles with the same charge affect each otherl Describe how particles of different charges affect each otherl Discuss how atoms are related to electricityl Explain how cathode rays and radioactivity are related to atomic structurel Explain how Rutherford's experiment showed the existence of the nucleusl Describe alpha, beta, and gamma radiationl Name and describe the three subatomic particlesl Determine the number of protons, neutrons, and electrons in an atom or ionl Explain how an ion differs from an atoml Determine number of subatomic particles from the symbol of a nuclidel Identify isotopesl Calculate average atomic mass from relative abundancel Describe the changes that accompany nuclear reactionsl Define radioactivityl Perform calculations with half-lives

Atomic Structure, Atomic Number, Mass Number

The atom is the basic unit of matter. It is the smallest particle of an element that still has the characteristicsof that element. Every atom has a positively charged central nucleus, composed of positively chargedprotons (1+) and uncharged neutrons. The nucleus is surrounded by a number of negatively chargedelectrons (1-). Like magnets, particles of opposite charge are attracted to one another, and particles with thesame charge are repelled from one another. The structure of the atom that we are familiar with today has along history to which many people contributed. The discovery of these particles and their arrangement, relativemasses, and charges required many experiments.

The word atom comes from the Greek word atomos, meaning "indivisible," and indeed the idea that all matteris composed of tiny, indivisible particles goes back to ancient Greece. The first scientist in modern times torevive this idea was a British scientist named John Dalton. In 1803, after observing that elements in a givencompound always exist in the same proportions, Dalton formulated his atomic theory of matter. The theoryhad four main points:

l Each element is composed of extremely small particles called atomsl All atoms of a given element are identical; the atoms of different elements are different and have different properties (including

different masses)l Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are neither created nor destroyed in

chemical reactionsl Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and

kind of atoms.

Michael Faraday was an English chemist who, in 1832-33, carried out a series of experiments attempting touse electricity to isolate elements from known compounds. His work led him to discover that the amount ofelectricity applied to a sample compound is related to the amount of an element that is isolated. He concludedthat the structure of the atom must play a role in this.

Building on Faraday's experiments, English physicist J. J. Thompson discovered that electricity consisted of tinysubatomic particles called electrons. In 1897, Thompson was the first person to measure the mass-to-chargeratio of a single electron. He found that the particles he observed were common to all atoms. Thomson alsoknew that the negative charge he measured must be balanced by a particle of positive charge, since there wasno overall charge on an atom. Thomson devised his own model of the atom, which is often described as theblueberry muffin model: he pictured an atom as a uniform ball of positive charge, with the electrons scattereduniformly throughout, like blueberries in a muffin.

The American physicist R. A. Millikan determined the unit of charge using oil drops and an electrical field. Bycalculating the velocity with which a charged drop was deflected from the electrical field, Millikan was able to

calculate the charge on the drop. The mass could also be calculated. By performing many experiments,Millikan was able to calculate the charge on an individual electron and the mass of an individual electron.

The true arrangement of particles in the atom was discovered by Ernest Rutherford, another English physicist.Rutherford performed his famous platinum and gold foil experiments with a college student who was named

Ernest Marsden. They determined that some positively charged helium nuclei, alpha particles (He2+) weredeflected by platinum foil, at angles greater than 90 degrees. Sometimes the particles were deflected almost180 degrees, back the way they came! For an alpha particle to be deflected that much, it must pass very closeto a group of particles that also had a positive charge. The group of particles would have to have a strongpositive charge, concentrated in a relatively small area. That is what the nucleus of an atom is a relativelysmall area with a concentrated positive charge.

The individual protons and neutrons were more difficult to isolate and identify. At this time, it was knownthrough Rutherfords and others experiments that while the proton and electron had the same charge, that themass of the proton was much greater. In any atom, the number of protons and electrons are the same. Theexistence of an uncharged particle was discovered by comparing the number of positive charges (number ofprotons) in an atom with the mass. The number of mass units was almost twice the number of protons. Wherewas the extra mass coming from?

James Chadwicks experiments with beryllium and positively charged alpha particles revealed the presence ofthe neutron. By bombarding beryllium with alpha particles, Chadwick was able to observe that an unchargedparticle was given off, which had the same mass as a proton. Chadwick used a specific type of beryllium,beryllium-9. The number 9 is the mass number, which indicates the total number if protons and neutrons inan atom. The number of protons in an atom of a specific element always identifies the element. The number ofneutrons in atoms of the same element may be different. When neutrons are different in atoms of the sameelement, they are called isotopes. The element neon has two common isotopes, neon-20 and neon-21. Bothisotopes have 10 protons, since neon is element number 10 in the Periodic Table. By subtracting the numberof protons from the mass number, we see that neon-20 has 10 neutrons, and neon-21 has 11 neutrons.Beryllium-9 has 4 protons and 3 neutrons.

Mass number = number of protons + number of neutrons.

Each element is identified by a unique atomic number, which is shown on the Periodic Table. The atomicnumber is equal to the number of protons in each atom of the element. The number of protons and electronsin a neutral atom of an element is always the same.

How does this work? Lets complete the chart for an example, oxygen-16. The mass number is the numberafter the dash, 16.

For an atom of Ca-40, the mass number is the number after the dash, 40. The atomic number can be obtainedby looking for Ca on the Periodic Table. The atomic number is 20, the number of protons.

Mass number number of protons = number of neutrons.

Number of neutrons equals 20. If no charge is specified, the number of electrons is the same as the numberof protons. The number of electrons is therefore 20.

We can follow this process for any isotope for which a mass number is designated. If we do not know the massnumber, we can use the Periodic Table to obtain an average atomic mass, which is usually a decimal, thenround it to the nearest whole number. For instance, the element sulfur, S, has an average atomic mass on thePeriodic Table of 32.07. By rounding this number to the nearest whole number, 32, we now have the massnumber for a common isotope of sulfur. We can obtain the number of protons from the atomic number on thePeriodic Table, and the number of electrons will be the same if no charge is indicated.

Ions and Ion Formation

Symbol Atomic Number Number of Protons Number of Electrons Mass Number Number of Neutrons

O 8 8 8 16 8

Symbol Atomic Number Number of Protons Number of Electrons Mass Number Number of Neutrons

Ca 20 20 20 40 20

The number of electrons can vary if charged particles called ions are formed. How can we determine thenumber of protons and electrons in an ion? Remember that the atomic number always gives the number ofprotons. We can find the atomic number on the Periodic Table, by looking in the same box as the symbol foran element. Each proton has a 1+ charge. An atom of an element has the same number of protons andelectrons, and that means the total charge adds to zero. Atoms of an element are often said to be neutral forthis reason.

Example: (Use the formula: positive charge + negative charge = net charge)

The element sodium, Na, has atomic number 11, meaning 11 protons and 11 electrons: (11+) + (11-) = 0.There is zero net (overall) charge, the atom is neutral; it is not an ion.

An atom of chlorine has atomic number 17, 17 protons and 17 electrons: (17+) + (17-) = 0. Zero net charge,this is a neutral atom.

Lets try this for ions. An ion is an atom that carries a charge either through gaining or losing an electron. Acation is an atom that loses an electron and has a positive charge. An anion is an atom that gains an electronand has a negative charge. We do not observe negative ions by themselves, they are balanced by positiveions. Using an everyday substance as our example, we can determine the number of protons and electrons for

a positive ion, Na1+, and a negative ion, Cl1-. These ions are found in table salt, sodium chloride, NaCl. Theions are formed from the elements sodium, Na, and chlorine, Cl.

The sodium ion, Na1+, has atomic number 11. This means 11 protons, 11 positive charges. The total charge is1+, overall. That means that there is one extra or unbalanced charge. What about the other 10 positivecharges? They are still there in the nucleus, and their charge is balanced by 10 electrons. Use the formulapositive charge + negative charge = net charge.

We already know that 11 protons = 11+, and net charge = 1+ ; so substitute into the equation:

(11+) + number of negative charges = 1+

The math on this one is easy! There are 10 negative charges (10-), which means 10 electrons. The sodium ion,Na1+ , has 10 electrons; a neutral atom of sodium has 11 electrons. The sodium atom loses one electron toform the positive ion. Where does the electron go? Lets take a look at the negative ion in the NaCl compound.

For the chloride ion, Cl1-, the process is the same. The atomic number is 17, indicating 17 protons. The netcharge is 1-. How many of those negative charges are balanced with a positive charge?

Remember: Positive charge + negative charge = net charge

(17+) + number of negative charges = 1-

number of negative charges = 18- , and number of negative charges equals number of electrons, 18!

A neutral chlorine atom has 17 electrons. The chloride ion has 18 electrons. A chlorine atom has to gain anelectron to form the negative ion. Therefore, ions are formed by the loss or gain of electrons. Thisprocess works for any ion, for which you know the symbol and charge. The symbol allows you to determine theatomic number by looking at a Periodic Table.

Calculating the Average Atomic Mass

The average atomic mass of an element is determined by averaging the natural abundance of its isotopes.

Element IsotopeMassNumber

Mass (amu)FractionalAbundance

Average Atomic Mass

Carbon12

C 12 12 (exactly) 98.89%

12.016

13C 13 13.003 1.11

6

Chloride35

Cl 35 34.969 75.53 35.4517

Example: Use the table above to calculate the average atomic mass of silicon.

Multiply the mass by the percent abundance for each of the three isotopes of silicon. (Change the percentageabundance to a decimal by moving the decimal point two places.)

(27.977) (0.9221) = 25.80 amu

(28.976) (0.0470) = 1.362 amu

(29.974) (0.0309) = 0.926 amu

Find the sum of the results. 25.80 + 1.36 + 0.926 = 28.08 amu

Round to three significant figures, 28.1 amu.

The average atomic mass of silicon is 28.1 amu.

Nuclear processes and Radioactivity

The French researcher Antoine Henri Becquerel, made the first observations of natural, spontaneousradioactivity. While studying fluorescent properties of substances, he noticed that photographic plates placednext to certain uranium compounds became darkened, as if they had been exposed to light! This is not whathe had intended to study, but this accidental discovery had an enormous impact. One of the students in his labwas Marie Curie, who went on to win the Nobel Prize in physics in 1903 with her husband Pierre Curie. Shewas awarded another Nobel Prize in 1911 for chemistry for her work on radium and polonium. Marie Curie isone of only three people to win two Nobel Prizes in science. Her daughter and son-in-law shared a Nobel Prizein chemistry in 1935.

Because of these discoveries, Rutherford and other researchers were aware of the existence of elements thatare naturally radioactive. This means that the elements undergo a specific type of change by emitting particlesof different sizes and charges, with differing amounts of energy, from the nuclei of atoms. Nuclear decay(radioactive decay) causes the number of protons in the nucleus to change. Since the identity of an element isdefined by the number of protons in its nuclei, nuclear decay results in the transmutation of one element toanother. This type of change, nuclear change, is very different from ordinary chemical reactions.

Two types of particles are emitted during natural nuclear decay: alpha particles and beta particles. Alphaparticles consist of two neutrons and two protons bound together--the equivalent of a positively charged

helium ion (He2+). Lets ignore the charge for a moment, and look at the symbol in a new way. Alpha particlesare often written in the following form:

The 4 in the upper left is the mass number. The 2 in the lower left is the atomic number. The helium ion stillhas its 2+ charge; it is just not shown in many nuclear reactions. The arrangement of particles in the nucleusand the changes that occur in radioactive decay processes get special attention, and the charge is typicallyignored. Charges are most often important in ordinary chemical reactions, where they receive much moreconsideration.

An example of alpha decay is the naturally occurring decay of uranium-238

37Cl 37 36.966 24.47

17

Silicon28

Si 28 27.977 92.21

28.09

14

29Si 29 28.976 4.70

14

30Si 30 29.974 3.09

14

4

2He

238

234 4

An alpha particle also has a 2+ charge, even though we do not see/write the charge in nuclear reactions. Notethat the element starting material on the left of the arrow, uranium-238, is not the same as the elementproducts helium-4 or thorium-234.

A beta particle is formed when a neutron from the nucleus breaks down into a proton and an electron. Theelectron, called a beta particle, is expelled from the nucleus. It is important to remember that a beta particle(-) is not one of the original set of electrons that are outside the nucleus.

An example of naturally occurring beta decay is thorium-234:

The beta symbol is used to be sure that we realize that this is not an electron from outside the nucleus. Wecan also recognize the correct mass and charge by using another symbol, e:

Note that the element starting material on the left of the arrow, thorium-234, is not the same as the elementproduct protactinium-234.

In addition to alpha and beta particles, nuclear decay also produces gamma rays. Gamma rays are very highenergy electromagnetic radiation with essentially no charge or mass. They have more penetrating power thanX-rays.

Half-Lives

Not all radioactive elements decay at the same rate. Some artificially created isotopes are very unstable andonly exist for a few minutes before breaking down. At the other extreme, the radioactive isotope rubidium-87decays to strontium-87 extraordinarily slowly; it takes 60 billion years (!) for 50% of a particular sample of87Rb to decay. The time it takes for one half of a sample of a radioactive isotope to decay is called the

isotope's half-life. The half-life of 87Rb is therefore 60 billion years. Carbon-14 has a half-life of 5,730 years,and the half-life of the artificial radioisotope iodine-131 is 0.022 years (about eight days).

92

U 90

Th +2

He

234

90Th

234

91Pa + -

234

90Th

234

91Pa +

0

-1e

Chapter 5

Lesson 5: Atomic Structure

OBJECTIVESl Understand the term valence shelll Describe the distinguishing characteristics of an ionic bondl Describe some properties of an ionic compoundl Explain the octet rulel Apply the octet rule to show why atoms form ionsl Draw Lewis dot diagrams to show the valence electrons of an atom of any elementl Recognize anions and cationsl Recognize binary compoundsl Learn the formula and charge for some common ionsl Write correct formulas for binary compounds using common cations and anionsl Describe covalent bondsl Identify single, double, and triple bondsl Write Lewis structures for covalent compounds, including double and triple bonds

Ionic and Covalent Bonds

Now that we have discussed the formation of charged ions from neutral atoms, we need to keep in mind thatpositive ions and negative ions are always found together. They are held together by electric forces;remember the phrase opposites attract as a reminder. Clothes in a clothes dryer that cling together are alsosaid to be held together by static electricity; fabrics may hold charge on their woven surfaces, and if thecharges are different, the two fabrics will be attracted to each other.

Ionic Compounds

Ionic compounds are held together by ionic bonds, and are composed of one or more cations (positivelycharged ions) and one or more anions (negatively charged ions). In an ionic bond, one or more electrons areremoved from one atom and attached to another, resulting in oppositely charged particles that are attracted toeach other. Many ionic compounds are binary compounds, which means that they are composed of only twoatoms (e.g., potassium chloride, KCl). Other ionic compounds have more than two atoms (e.g., CaCl

2). Some

ionic compounds contain polyatomic ions, which are groups of covalently bonded atoms that carry a

collective charge. In sodium nitrate (NaNO3), Na+ is a cation and NO3

- is a polyatomic anion. In formulas for

ionic compounds, the cation always precedes the anion (e.g., O2- + Mg2+ MgO).

Ionic compounds share a set of characteristic properties in common. These defining properties include:

(1) high melting points, often above 700C. Energy is required to melt solids, and ionic solids are heldtogether by ionic bonds, the opposite electrical charges on ions. These bonds are very strong, and must bebroken for an ionic solid to melt.

(2) brittleness; ionic compounds shatter easily, in specific directions, typically along a plane. Think of graniteor rock salt as examples.

(3) solubility in water; almost all ionic compounds are soluble in water. Ionic bonds are broken apart bywater. Separate ions move freely in water.

(4) ionic solids do NOT conduct electricity when they are in the solid state;

(5) ionic solids DO conduct electricity when molten or dissolved in water.

Properties (4) and (5) are related; In the solid state, ions are held rigidly in place by strong bonds; whenmelted to the liquid state, ions can movemore freely. When ionic solids are dissolved in water, they are also mobile and therefore conducting.

Covalent Compounds

Covalent bonds result from the sharing of one or more pairs of electrons between two atoms. Covalent bondsare very strong. The atoms held together by covalent bonds do not usually carry an ionic charge. Covalentcompounds are very different from ionic compounds because they do not have a charge on the individualparticles that make up the compound. The characteristics of covalent compounds (compounds composed ofelements bonded together by covalent bonds) are not always similar to each other. This is different from ioniccompounds, which have quite a few properties in common. Covalent molecules do typically:

(1) have lower boiling points that ionic compounds, often below 300C

(2) do not conduct electricity in the solid, molten, or dissolved state

(3) have a wide variety of solubilities in water and other solvents, depending on the specifics of thestructure of each molecule.

We will revisit bonding later in this chapter.

Counting Atoms and the Mole

Atoms are so small that we cannot see one atom at a time without sophisticated instruments. Even then, wecan only view atoms indirectly. For instance, when we see a sample of an element, we are examining a largebulk sample of atoms of that element. This is called a macroscopic sample. A sample too small to be seenunaided is considered microscopic. Most of our discussions will involve macroscopic groups of atoms or ions ormolecules.

There is a useful term for groups of atoms, ions, or molecules, called the mole. A mole is the name of a group,

with a specific number of items in the group. The number of items in a mole is 6.022 x 1023. This number isalso called Avogadros number, for one of the early researches working on atomic mass. The periodic tableprovides information that is useful, and knowing Avogadros number makes it more useful. The average atomicmass reported on the Periodic Table for each element is the average of the atomic masses for a typical bulksample of the element, given in grams (g). It is also the average mass for a single atom, in atomic mass units(amu). The mole is useful when we consider the masses involved in chemical equations. We may read anequation as the reaction of atoms, ions, or molecules with each other, in specific ratios. We may also readequations as the reaction of bulk samples of atoms, ions, molecules with each other. Either way, the ratios arealways consistent with the actual observations; these ratios must be consistent on microscopic andmacroscopic level. For example, consider the reaction:

2 H2 + O2 2 H2O

This equation means two molecules of hydrogen react with one molecule of oxygen to form 2 molecules ofwater. If there is no digit in front of a formula, a one is understood. If the number of hydrogen molecules

(two) is multiplied by Avogadros number (6.022 x 1023), then we have 1.2044 x 1024 molecules, which equals

two moles of hydrogen. Perform the same step for oxygen (one molecule multiplied by 6.022 x 1023 = 6.022 x

1023 molecules of oxygen = one mole of oxygen) and water, and we can then read: two moles of hydrogen

react with one mole of oxygen to form 2 moles of water. The identity 6.022 x 1023 = 1 mole can now beused to convert numbers of atoms to moles.

Atomic Mass and the Mole

Atomic mass can also be related to number moles, using the values given on the Periodic Table. The averageatomic mass in grams is also the mass of one mole of atoms of the element. For carbon (C), theaverage atomic mass is 12.01 amu per atom and 12.01 grams per one mole. Remember, the coefficients inbalanced equations refer to moles, not grams. For example:

C + O2 CO2

This equation tells us that one mole of carbon reacts with one mole of oxygen to form one mole of carbondioxide (CO2). From the Periodic Table, we find that one mole of C weighs 12.01 grams. In O2, 2 oxygen

atoms are chemically combined, so we must add the masses of both to find that 1 mole of O2

weighs 32

grams. For CO2, adding the masses of one mole of C and 2 moles of O results in a mass of 44.0 grams for one

mole of CO2. 12.01 g/mol and 32.0 g/mol are the molar masses of carbon and oxygen, respectively. The molar

mass of CO2 is therefore 44.0 g/mol. Molar mass can be used to convert between grams and moles for any

substance. All we need to know is the correct formula and the average atomic masses from the Periodic Table.

Problem: 1.5 moles of C = how many grams of C?

You may recognize this strategy from our unit conversions with the metric system.

Lewis Diagrams, Valence Electrons, and the Octet Rule

Valance Electrons

The electrons found at the greatest distance from the nucleus are called valance electrons. Valance electronsaccount for almost all of the chemistry we see in the world around us. These are the electrons involved in ionformation and/or covalent bond formation, depending on the substances involved and the conditions. Ionicbonding involves the complete transfer of electrons from one element to another. Covalent bonds involvesharing one or more pairs of electrons between two atoms. Our discussion of ion formation helps us realize theusefulness of Lewis Diagrams, and helps us understand the concept of valance or outershell electrons.

The Octet Rule

Atoms lose, gain or share electrons for a reason; is not a random process. Electron transfer and sharing occurbecause every spontaneous process tends to lead to a result that is lower in energy than the starting position.Water runs downhill because it has less stored energy at the bottom of the hill than at the top. The sameprinciple applies to the arrangement of electrons. The transfer or sharing of electrons between to atoms canoften lead to a result that is lower in energy. The lowest energy arrangements of electrons involve a filledouter shell. Many common ions have eight electrons in their filled outer shell. The tendency to have eightelectrons in the valance shell after ion formation or bond formation is called the octet rule.

Lewis Structures

Lets examine one of the ways scientists use to represent elements and their outer shell electrons. This isuseful to us, because almost all of the chemistry we observe around us everyday can be explained by transferor sharing of outer shell electrons. G.N. Lewis noticed some patterns in reactivity and devised a method forpredicting the number of bonds many elements would form. The symbol of the element is used to representthe nucleus, and all inner shell electrons; small dots are used to represent outer shell (valance) electrons. Forexample, the element carbon has 4 valance electrons. The Lewis structure for a atom of carbon looks like this:

Structures like this became known as Lewis dot structures. Lewis had several rules for these structures. First,the valence electrons should be placed as far apart as possible around the symbol of the element, as one, two,three or four valance electrons are added. After the fourth electron is added, we begin pairing electrons. Thisleaves us with electrons around a symbol, at the four poles, north, south, east, west.

Examples of Lewis Dot structures for elements having four or fewer valance electrons:

Examples of Lewis Dot structures for elements having more than four valance electrons:

Solution: 1.5 moles C x12.01 g C

= 18.01 grams C1 mol C

C

LithiumBerylliumAluminumCarbon

Li Be Al C

Drawing Lewis Dot structures seems easy now, right? Except for one thing: how do we determine how manyvalance electrons each element has in its outer shell? To answer this question, we use the Periodic Table.

There are two sets or numbers across the top of the Table; they are specified as New Designation, or OriginalDesignation. Numbers in the Original Designation are roman numerals with a letter: IA, IIA, IIIA, IVA, and soon. The Original Designation numbers with an A after them denote the number of valance electrons for theelements in the column below them. For instance, nitrogen (N) is in column VA. Nitrogen therefore has 5valance electrons. If we check the correct Lewis Dot structure given for nitrogen above, we see 5 dots,representing 5 electrons. The New Designation numbers run from 1 to 18, and are Arabic numbers with noletter. For columns 1 and 2, the new designation number is the number of valance electrons; for columns 13-18, we subtract 10 to get the number of valance electrons. For instance, oxygen (O) is in column 16. Subtract10 from 16, and we are left with 6; 6 valance electrons for oxygen. Examine the correct Lewis Dot structurefor oxygen given above, and we see 6 dots for each of the 6 electrons. Nitrogen in column 15 (in this system)would still have 5 valance electrons. For the columns with a letter B: in the Original Designation scheme, and3-12 in the New Designation scheme, the number of valance electrons is not so simply predicted. Theseelements have more complex rules governing their behavior, which we will discuss later. Keep in mind that theelements in columns 3-12 are metals, and therefore have a general tendency to form positively charged ions.

Covalent Bonds in Lewis Structures: Sharing Electrons

Lewis dot structures can also be used to predict the structure of molecules held together by covalent bonds.Lets use the example of water, H2O, as we go through the steps of writing a Lewis structure.

Step 1: Have the correct formula for your compound. We know that H2O is correct for water.

Step 2: Choose a central atom (for more complex molecules such as hydrocarbons, we will give additionalrules later). Water has 2 hydrogen atoms, H and H; and one oxygen atom, O. A general rule of thumb is thatwhen we only have one of an element (in this case O), that is what we choose as the central atom. So, weplace the O in the middle, and the other atoms around it, as far apart as possible from each other on differentsides of the central atom (in this case on opposite sides of the O): H O H

Step 3: Determine the number of valence electrons for each atom by consulting the Periodic Tableto determine in which group an element is. Hydrogen is in group one, so it has one valence electron.Oxygen is in group six, so it has six valence electrons.

2 x H (1 electron each) = 2 electrons1 x O (6 electrons each) = 6 electrons______TOTAL = 8 valance electrons

Step 4: Begin placing valance electrons in the structure, and continue until all valance electrons areplaced.

a. Place a pair of electrons between the central atom and each outer atom, continue until all outeratoms are connected to the center.

H O H

b. Subtract the number now placed, 4, from the total 8, and we see that 4 valance electrons remain.

8 total valance electrons- 4 electrons now placed (after step 4)4 remaining valance electrons to place

Nitrogen Oxygen Fluorine Neon

N

O

F

Ne

c. Examine the structure, and check for compliance with the Octet rule. Does each element shownhave 8 electrons around it? The answer here is no. Before we continue, there is a new piece of informationwe need. Hydrogen is an exception to the Octet rule, and does not need 8 electrons to fill its outershell. Each hydrogen requires only 2 electrons to fill its outer shell. By sharing 2 electrons with the oxygen,each hydrogen has now filled its outer shell. (Remember: 2 electrons between 2 bonds in a Lewis Structuremeans that the electrons are shared between those two atoms in a covalent bond). What about the oxygen?The oxygen is sharing 4 electrons. Oxygen does follow the Octet Rule. How do we satisfy that requirement? Weuse the remaining 4 valance electrons from our original count. To do this correctly, we place the electrons in 2pairs, in the remaining spaces around the oxygen.

d. Now, do a final check of total valance electrons shown in the structure. There are 8 total valanceelectrons, and 8 electrons around oxygen. That is the completed, correct Lewis dot structure for water.

This method works well for a wide variety of simple molecules and can be readily adapted for hydrocarbonsand derivatives.

Here are more examples:

Methane, CH4

Step 1: Name and formula are correctly matched (for these practice examples, this will always be the case.)

Step 2: Central atom is C; place other atoms around it, as far from each other as possible, and close aroundthe central atom.

Step 3: Count total valance electrons4 x H (1 valance electron each) = 4 valance electrons +1 x C (4 valance electrons each) = 4 valance electronsTOTAL = 8 valance electrons

Step 4a: Begin placing valance electrons between the central atom and the outer atoms, continue until allouteratoms are connected to the center.

Step 4b: Subtract the number of placed valance electrons from our total: 8 8 = 0. All valance electrons areplaced; this indicates we may be finished.

Go to Step 4d: check for placement of all electrons, and compliance with the Octet Rule. Since C has 8electrons around it, the Octet Rule is obeyed. Each hydrogen has 2 electrons, so their outer shells are filled.We have completed another correct Lewis dot structure!

It may seem strange that we count the electrons between H and C twice, once for H, and once for C. Keep inmind that the atoms are sharing the electrons, in the same way that you might share a room at home withyour brother or sister. Both of you have a room, even if you are sharing it! Both atoms have a pair ofelectrons, even if they are sharing.

H O H

H

H C H

H

H

H C H

H

Acetylene, C2H

2

Step 1: Correct formulaStep 2: Central atom is C for hydrocarbons; when there are 2 or more carbons, place them side by side in thecenter. Place other atoms around it, as far from each other as possible, and close around the central atom.H C C H

Step 3: Count total valance electrons

2 x H (1 valance electron each) = 2 valance electrons +2 x C (4 valance electrons each) = 8 valance electronsTOTAL = 10 valance electrons

Step 4a: Begin placing valance electrons between the central atom and the outer atoms, continue until al