Atoms, Molecules,and Ions41A Look Ahead We begin with a historical perspective of the search for the fundamentalunits of matter. The modern version of atomic theory was laid by JohnDalton in the nineteenth century, who postulated that elements are composedof extremely small particles, called atoms. All atoms of a given element areidentical, but they are different from atoms of all other elements. (2.1) We note that, through experimentation, scientists have learned that an atomis composed of three elementary particles: proton, electron, and neutron. Theproton has a positive charge, the electron has a negative charge, and theneutron has no charge. Protons and neutrons are located in a small regionat the center of the atom, called the nucleus, while electrons are spread outabout the nucleus at some distance from it. (2.2) We will learn the following ways to identify atoms. Atomic number is thenumber of protons in a nucleus; atoms of different elements have differentatomic numbers. Isotopes are atoms of the same element having a differentnumber of neutrons. Mass number is the sum of the number of protons andneutrons in an atom. Because an atom is electrically neutral, the number ofprotons is equal to the number of electrons in it. (2.3) Next we will see how elements can be grouped together according to theirchemical and physical properties in a chart called the periodic table. Theperiodic table enables us to classify elements (as metals, metalloids, andnonmetals) and correlate their properties in a systematic way. (2.4) We will see that atoms of most elements interact to form compounds, whichare classifi ed as molecules or ionic compounds made of positive (cations)and negative (anions) ions. (2.5) We learn to use chemical formulas (molecular and empirical) to representmolecules and ionic compounds and models to represent molecules. (2.6) We learn a set of rules that help us name the inorganic compounds. (2.7) Finally, we will briefl y explore the organic world to which we will returnin a later chapter. (2.8)Since ancient times humans have pondered the nature of matter. Our modernideas of the structure of matter began to take shape in the early nineteenthcentury with Daltons atomic theory. We now know that all matter is made ofatoms, molecules, and ions. All of chemistry is concerned in one way or anotherwith these species.2.1 The Atomic Theory2.2 The Structure of the Atom2.3 Atomic Number, MassNumber, and Isotopes2.4 The Periodic Table2.5 Molecules and Ions2.6 Chemical Formulas2.7 Naming Compounds2.8 Introduction to OrganicCompoundsChapter OutlineStudent InteractiveActivityAnimationsCathode Ray Tube (2.2)Millikan Oil Drop (2.2)Alpha, Beta, and GammaRays (2.2)a-Particle Scattering (2.2)Media PlayerRutherfords Experiment (2.2)Formation of an IonicCompound (2.7)Chapter SummaryARISExample Practice ProblemsEnd of Chapter ProblemsQuantum TutorsEnd of Chapter Problems42 Atoms, Molecules, and Ions2.1 The Atomic TheoryIn the fi fth century b.c. the Greek philosopher Democritus expressed the belief thatall matter consists of very small, indivisible particles, which he named atomos (meaninguncuttable or indivisible). Although Democritus idea was not accepted by manyof his contemporaries (notably Plato and Aristotle), somehow it endured. Experimentalevidence from early scientifi c investigations provided support for the notion ofatomism and gradually gave rise to the modern defi nitions of elements and compounds.In 1808 an English scientist and school teacher, John Dalton, formulated aprecise defi nition of the indivisible building blocks of matter that we call atoms.Daltons work marked the beginning of the modern era of chemistry. The hypothesesabout the nature of matter on which Daltons atomic theory is based can besummarized as follows:1. Elements are composed of extremely small particles called atoms.2. All atoms of a given element are identical, having the same size, mass, andchemical properties. The atoms of one element are different from the atoms ofall other elements.3. Compounds are composed of atoms of more than one element. In any compound,the ratio of the numbers of atoms of any two of the elements present is either aninteger or a simple fraction.4. A chemical reaction involves only the separation, combination, or rearrangementof atoms; it does not result in their creation or destruction.Figure 2.1 is a schematic representation of the last three hypotheses.Daltons concept of an atom was far more detailed and specifi c than Democritus.The second hypothesis states that atoms of one element are different from atoms ofall other elements. Dalton made no attempt to describe the structure or compositionof atomshe had no idea what an atom is really like. But he did realize that thedifferent properties shown by elements such as hydrogen and oxygen can be explainedby assuming that hydrogen atoms are not the same as oxygen atoms.The third hypothesis suggests that, to form a certain compound, we need not onlyatoms of the right kinds of elements, but specifi c numbers of these atoms as well.John Dalton (17661844). English chemist, mathematician, and philosopher. In addition to the atomictheory, he also formulated several gas laws and gave the fi rst detailed description of color blindness, fromwhich he suffered. Dalton was described as an indifferent experimenter, and singularly wanting in thelanguage and power of illustration. His only recreation was lawn bowling on Thursday afternoons. Perhapsit was the sight of those wooden balls that provided him with the idea of the atomic theory.(b)Atoms of element X Atoms of element Y Compounds of elements X and Y(a)Figure 2.1 (a) According toDaltons atomic theory, atoms ofthe same element are identical,but atoms of one element aredifferent from atoms of otherelements. (b) Compound formedfrom atoms of elements X and Y.In this case, the ratio of theatoms of element X to the atomsof element Y is 2:1. Note that achemical reaction results only inthe rearrangement of atoms, notin their destruction or creation.This idea is an extension of a law published in 1799 by Joseph Proust, a Frenchchemist. Prousts law of defi nite proportions states that different samples of the samecompound always contain its constituent elements in the same proportion by mass.Thus, if we were to analyze samples of carbon dioxide gas obtained from differentsources, we would fi nd in each sample the same ratio by mass of carbon to oxygen.It stands to reason, then, that if the ratio of the masses of different elements in a givencompound is fi xed, the ratio of the atoms of these elements in the compound alsomust be constant.Daltons third hypothesis supports another important law, the law of multipleproportions . According to the law, if two elements can combine to form more thanone compound, the masses of one element that combine with a fi xed mass of the otherelement are in ratios of small whole numbers. Daltons theory explains the law ofmultiple proportions quite simply: Different compounds made up of the same elementsdiffer in the number of atoms of each kind that combine. For example, carbon formstwo stable compounds with oxygen, namely, carbon monoxide and carbon dioxide.Modern measurement techniques indicate that one atom of carbon combines with oneatom of oxygen in carbon monoxide and with two atoms of oxygen in carbon dioxide.Thus, the ratio of oxygen in carbon monoxide to oxygen in carbon dioxide is 1:2.This result is consistent with the law of multiple proportions ( Figure 2.2 ).Daltons fourth hypothesis is another way of stating the law of conservation ofmass, which is that matter can be neither created nor destroyed. Because matter ismade of atoms that are unchanged in a chemical reaction, it follows that mass mustbe conserved as well. Daltons brilliant insight into the nature of matter was the mainstimulus for the rapid progress of chemistry during the nineteenth century.Review of ConceptsThe atoms of elements A (blue) and B (orange) form two compounds shownhere. Do these compounds obey the law of multiple proportions?Joseph Louis Proust (17541826). French chemist. Proust was the fi rst person to isolate sugar from grapes.According to Albert Einstein, mass and energy are alternate aspects of a single entity called mass-energy.Chemical reactions usually involve a gain or loss of heat and other forms of energy. Thus, when energyis lost in a reaction, for example, mass is also lost. Except for nuclear reactions (see Chapter 23), however,changes of mass in chemical reactions are too small to detect. Therefore, for all practical purposes massis conserved.Carbon monoxideCarbon dioxideRatio of oxygen incarbon monoxide tooxygen in carbon dioxide: 1:2OC21OC11_ __ _Figure 2.2 An illustration of thelaw of multiple proportions.2.2 The Structure of the AtomOn the basis of Daltons atomic theory, we can defi ne an atom as the basic unit of anelement that can enter into chemical combination. Dalton imagined an atom that wasboth extremely small and indivisible. However, a series of investigations that began inthe 1850s and extended into the twentieth century clearly demonstrated that atomsactually possess internal structure; that is, they are made up of even smaller particles,which are called subatomic particles. This research led to the discovery of three suchparticleselectrons, protons, and neutrons.2.2 The Structure of the Atom 4344 Atoms, Molecules, and IonsThe ElectronIn the 1890s, many scientists became caught up in the study of radiation , the emissionand transmission of energy through space in the form of waves. Information gainedfrom this research contributed greatly to our understanding of atomic structure. Onedevice used to investigate this phenomenon was a cathode ray tube, the forerunner ofthe television tube ( Figure 2.3 ). It is a glass tube from which most of the air has beenevacuated. When the two metal plates are connected to a high-voltage source, thenegatively charged plate, called the cathode, emits an invisible ray. The cathode rayis drawn to the positively charged plate, called the anode, where it passes through ahole and continues traveling to the other end of the tube. When the ray strikes thespecially coated surface, it produces a strong fl uorescence, or bright light.In some experiments, two electrically charged plates and a magnet were added tothe outside of the cathode ray tube (see Figure 2.3 ). When the magnetic fi eld is on andthe electric fi eld is off, the cathode ray strikes point A. When only the electric fi eld ison, the ray strikes point C. When both the magnetic and the electric fi elds are off orwhen they are both on but balanced so that they cancel each others infl uence, the raystrikes point B. According to electromagnetic theory, a moving charged body behaveslike a magnet and can interact with electric and magnetic fi elds through which it passes.Because the cathode ray is attracted by the plate bearing positive charges and repelledby the plate bearing negative charges, it must consist of negatively charged particles.We know these negatively charged particles as electrons . Figure 2.4 shows the effectof a bar magnet on the cathode ray.An English physicist, J. J. Thomson, used a cathode ray tube and his knowledgeof electromagnetic theory to determine the ratio of electric charge to the mass of anindividual electron. The number he came up with was 21.76 3 108 C/g, where Cstands for coulomb, which is the unit of electric charge. Thereafter, in a series ofexperiments carried out between 1908 and 1917, R. A. Millikan succeeded in measuringthe charge of the electron with great precision. His work proved that the chargeon each electron was exactly the same. In his experiment, Millikan examined themotion of single tiny drops of oil that picked up static charge from ions in the air.He suspended the charged drops in air by applying an electric fi eld and followed theirAnimationCathode Ray TubeHigh voltage+Anode CathodeABCSNFluorescent screenFigure 2.3 A cathode ray tubewith an electric fi eld perpendicularto the direction of the cathoderays and an external magneticfi eld. The symbols N and Sdenote the north and south polesof the magnet. The cathode rayswill strike the end of the tube atA in the presence of a magneticfi eld, at C in the presence of anelectric fi eld, and at B when thereare no external fi elds present orwhen the effects of the electricfi eld and magnetic fi eld canceleach other.Electrons are normally associated withatoms. However, they can also be studiedindividually.Joseph John Thomson (18561940). British physicist who received the Nobel Prize in Physics in 1906for discovering the electron.Robert Andrews Millikan (18681953). American physicist who was awarded the Nobel Prize in Physicsin 1923 for determining the charge of the electron.AnimationMillikan Oil Dropmotions through a microscope ( Figure 2.5 ). Using his knowledge of electrostatics,Millikan found the charge of an electron to be 21.6022 3 10219 C. From these datahe calculated the mass of an electron:mass of an electron 5chargecharge/mass521.6022 3 10219 C21.76 3 108 C/g5 9.10 3 10228 gThis is an exceedingly small mass.RadioactivityIn 1895, the German physicist Wilhelm Rntgen noticed that cathode rays causedglass and metals to emit very unusual rays. This highly energetic radiation penetratedmatter, darkened covered photographic plates, and caused a variety of substances tofl uoresce. Because these rays could not be defl ected by a magnet, they could notcontain charged particles as cathode rays do. Rntgen called them X rays becausetheir nature was not known.(a) (b) (c)Figure 2.4 (a) A cathode ray produced in a discharge tube. The ray itself is invisible, but the fl uorescence of a zinc sulfi de coatingon the glass causes it to appear green. (b) The cathode ray is bent downward when a bar magnet is brought toward it. (c) When thepolarity of the magnet is reversed, the ray bends in the opposite direction.AtomizerViewingmicroscopeCharged plateX rayto producecharge onoil droplet(_)Smallhole(_)Charged plate Oil droplets Figure 2.5 Schematic diagramof Millikans oil drop experiment.2.2 The Structure of the Atom 45Wilhelm Konrad Rntgen (18451923). German physicist who received the Nobel Prize in Physics in 1901for the discovery of X rays.46 Atoms, Molecules, and IonsNot long after Rntgens discovery, Antoine Becquerel, a professor of physicsin Paris, began to study the fl uorescent properties of substances. Purely by accident,he found that exposing thickly wrapped photographic plates to a certain uraniumcompound caused them to darken, even without the stimulation of cathode rays. LikeX rays, the rays from the uranium compound were highly energetic and could not bedefl ected by a magnet, but they differed from X rays because they arose spontaneously.One of Becquerels students, Marie Curie, suggested the name radioactivity todescribe this spontaneous emission of particles and/or radiation. Since then, any elementthat spontaneously emits radiation is said to be radioactive.Three types of rays are produced by the decay, or breakdown, of radioactivesubstances such as uranium. Two of the three are defl ected by oppositely chargedmetal plates ( Figure 2.6 ). Alpha (a) rays consist of positively charged particles, calleda particles, and therefore are defl ected by the positively charged plate. Beta (b) rays,or b particles, are electrons and are defl ected by the negatively charged plate. Thethird type of radioactive radiation consists of high-energy rays called gamma (g)rays . Like X rays, g rays have no charge and are not affected by an external fi eld.The Proton and the NucleusBy the early 1900s, two features of atoms had become clear: they contain electrons,and they are electrically neutral. To maintain electric neutrality, an atom must containan equal number of positive and negative charges. Therefore, Thomson proposed thatan atom could be thought of as a uniform, positive sphere of matter in which electronsare embedded like raisins in a cake ( Figure 2.7 ). This so-called plum-pudding modelwas the accepted theory for a number of years.Antoine Henri Becquerel (18521908). French physicist who was awarded the Nobel Prize in Physics in1903 for discovering radioactivity in uranium.Marie (Marya Sklodowska) Curie (18671934). Polish-born chemist and physicist. In 1903 she and herFrench husband, Pierre Curie, were awarded the Nobel Prize in Physics for their work on radioactivity. In1911, she again received the Nobel prize, this time in chemistry, for her work on the radioactive elementsradium and polonium. She is one of only three people to have received two Nobel prizes in science. Despiteher great contribution to science, her nomination to the French Academy of Sciences in 1911 was rejectedby one vote because she was a woman! Her daughter Irene, and son-in-law Frederic Joliot-Curie, sharedthe Nobel Prize in Chemistry in 1935.+Radioactive substanceLead blockFigure 2.6 Three types of raysemitted by radioactive elements.b rays consist of negativelycharged particles (electrons) andare therefore attracted by thepositively charged plate. Theopposite holds true for a raysthey are positively charged andare drawn to the negativelycharged plate. Because g rayshave no charges, their path isunaffected by an externalelectric fi eld.AnimationAlpha, Beta, and Gamma RaysPositive charge spreadover the entire sphereFigure 2.7 Thomsons model ofthe atom, sometimes describedas the plum-pudding model,after a traditional English dessertcontaining raisins. The electronsare embedded in a uniform,positively charged sphere.In 1910 the New Zealand physicist Ernest Rutherford, who had studied withThomson at Cambridge University, decided to use a particles to probe the structure ofatoms. Together with his associate Hans Geiger and an undergraduate named ErnestMarsden, Rutherford carried out a series of experiments using very thin foils of goldand other metals as targets for a particles from a radioactive source ( Figure 2.8 ). Theyobserved that the majority of particles penetrated the foil either undefl ected or with onlya slight defl ection. But every now and then an a particle was scattered (or defl ected) ata large angle. In some instances, an a particle actually bounced back in the directionfrom which it had come! This was a most surprising fi nding, for in Thomsons modelthe positive charge of the atom was so diffuse that the positive a particles should havepassed through the foil with very little defl ection. To quote Rutherfords initial reactionwhen told of this discovery: It was as incredible as if you had fi red a 15-inch shell ata piece of tissue paper and it came back and hit you.Rutherford was later able to explain the results of the a-scattering experiment interms of a new model for the atom. According to Rutherford, most of the atom mustbe empty space. This explains why the majority of a particles passed through the goldfoil with little or no defl ection. The atoms positive charges, Rutherford proposed, areall concentrated in the nucleus, which is a dense central core within the atom. Wheneveran a particle came close to a nucleus in the scattering experiment, it experienced a largerepulsive force and therefore a large defl ection. Moreover, an a particle traveling directlytoward a nucleus would be completely repelled and its direction would be reversed.The positively charged particles in the nucleus are called protons . In separateexperiments, it was found that each proton carries the same quantity of charge as anelectron and has a mass of 1.67262 3 10224 gabout 1840 times the mass of theoppositely charged electron.At this stage of investigation, scientists perceived the atom as follows: The massof a nucleus constitutes most of the mass of the entire atom, but the nucleus occupiesonly about 1/1013 of the volume of the atom. We express atomic (and molecular)dimensions in terms of the SI unit called the picometer ( pm ) , where1 pm 5 1 3 10212 mAnimationa-Particle ScatteringDetecting screen SlitGold foil(a) (b)ParticleemitterFigure 2.8 (a) Rutherfordsexperimental design for measuringthe scattering of a particles by apiece of gold foil. Most of the aparticles passed through the goldfoil with little or no defl ection. Afew were defl ected at wide angles.Occasionally an a particle wasturned back. (b) Magnifi ed view ofa particles passing through andbeing defl ected by nuclei.2.2 The Structure of the Atom 47Ernest Rutherford (18711937). New Zealand physicist. Rutherford did most of his work in England(Manchester and Cambridge Universities). He received the Nobel Prize in Chemistry in 1908 for hisinvestigations into the structure of the atomic nucleus. His often-quoted comment to his students was thatall science is either physics or stamp-collecting.Johannes Hans Wilhelm Geiger (18821945). German physicist. Geigers work focused on the structureof the atomic nucleus and on radioactivity. He invented a device for measuring radiation that is now commonlycalled the Geiger counter.Ernest Marsden (18891970). English physicist. It is gratifying to know that at times an undergraduatecan assist in winning a Nobel Prize. Marsden went on to contribute signifi cantly to the development ofscience in New Zealand.A common non-SI unit for atomic length isthe angstrom (; 1 = 100 pm).Media PlayerRutherfords Experiment48 Atoms, Molecules, and IonsA typical atomic radius is about 100 pm, whereas the radius of an atomic nucleus isonly about 5 3 1023 pm. You can appreciate the relative sizes of an atom and itsnucleus by imagining that if an atom were the size of a sports stadium, the volumeof its nucleus would be comparable to that of a small marble. Although the protonsare confi ned to the nucleus of the atom, the electrons are conceived of as being spreadout about the nucleus at some distance from it.The concept of atomic radius is useful experimentally, but we should not inferthat atoms have well-defi ned boundaries or surfaces. We will learn later that the outerregions of atoms are relatively fuzzy.The NeutronRutherfords model of atomic structure left one major problem unsolved. It was knownthat hydrogen, the simplest atom, contains only one proton and that the helium atomcontains two protons. Therefore, the ratio of the mass of a helium atom to that of ahydrogen atom should be 2:1. (Because electrons are much lighter than protons, theircontribution to atomic mass can be ignored.) In reality, however, the ratio is 4:1.Rutherford and others postulated that there must be another type of subatomic particlein the atomic nucleus; the proof was provided by another English physicist,James Chadwick, in 1932. When Chadwick bombarded a thin sheet of berylliumwith a particles, a very high-energy radiation similar to g rays was emitted by themetal. Later experiments showed that the rays actually consisted of a third type ofsubatomic particles, which Chadwick named neutrons, because they proved to beelectrically neutral particles having a mass slightly greater than that of protons. Themystery of the mass ratio could now be explained. In the helium nucleus there aretwo protons and two neutrons, but in the hydrogen nucleus there is only one protonand no neutrons; therefore, the ratio is 4:1.Figure 2.9 shows the location of the elementary particles (protons, neutrons, andelectrons) in an atom. There are other subatomic particles, but the electron, the proton,If the size of an atom were expanded tothat of this sports stadium, the size of thenucleus would be that of a marble.James Chadwick (18911972). British physicist. In 1935 he received the Nobel Prize in Physics forproving the existence of neutrons.ProtonNeutronFigure 2.9 The protons andneutrons of an atom are packedin an extremely small nucleus.Electrons are shown as cloudsaround the nucleus.and the neutron are the three fundamental components of the atom that are importantin chemistry. Table 2.1 shows the masses and charges of these three elementaryparticles.2.3 Atomic Number, Mass Number, and IsotopesAll atoms can be identifi ed by the number of protons and neutrons they contain. Theatomic number (Z) is the number of protons in the nucleus of each atom of an element.In a neutral atom the number of protons is equal to the number of electrons,so the atomic number also indicates the number of electrons present in the atom. Thechemical identity of an atom can be determined solely from its atomic number. Forexample, the atomic number of fl uorine is 9. This means that each fl uorine atom has9 protons and 9 electrons. Or, viewed another way, every atom in the universe thatcontains 9 protons is correctly named fl uorine.The mass number (A) is the total number of neutrons and protons present in thenucleus of an atom of an element. Except for the most common form of hydrogen,which has one proton and no neutrons, all atomic nuclei contain both protons andneutrons. In general, the mass number is given bymass number 5 number of protons 1 number of neutrons5 atomic number 1 number of neutrons(2.1)The number of neutrons in an atom is equal to the difference between the mass numberand the atomic number, or (A 2 Z). For example, if the mass number of a particularboron atom is 12 and the atomic number is 5 (indicating 5 protons in thenucleus), then the number of neutrons is 12 2 5 5 7. Note that all three quantities(atomic number, number of neutrons, and mass number) must be positive integers, orwhole numbers.Atoms of a given element do not all have the same mass. Most elements havetwo or more isotopes, atoms that have the same atomic number but different massnumbers. For example, there are three isotopes of hydrogen. One, simply known ashydrogen, has one proton and no neutrons. The deuterium isotope contains one protonand one neutron, and tritium has one proton and two neutrons. The accepted wayto denote the atomic number and mass number of an atom of an element (X) is asfollows:mass numberatomic number8n8nZAXChargeParticle Mass (g) Coulomb Charge UnitElectron* 9.10938 3 10228 21.6022 3 10219 21Proton 1.67262 3 10224 11.6022 3 10219 11Neutron 1.67493 3 10224 0 0*More refi ned measurements have given us a more accurate value of an electrons mass than Millikans.TABLE 2.1 Mass and Charge of Subatomic Particles2.3 Atomic Number, Mass Number, and Isotopes 49Protons and neutrons are collectivelycalled nucleons.50 Atoms, Molecules, and IonsThus, for the isotopes of hydrogen, we write11H 21H 31Hhydrogen deuterium tritiumAs another example, consider two common isotopes of uranium with mass numbersof 235 and 238, respectively:23592U 23892UThe fi rst isotope is used in nuclear reactors and atomic bombs, whereas the secondisotope lacks the properties necessary for these applications. With the exception ofhydrogen, which has different names for each of its isotopes, isotopes of elementsare identifi ed by their mass numbers. Thus, the preceding two isotopes are calleduranium-235 (pronounced uranium two thirty-fi ve) and uranium-238 (pronounceduranium two thirty-eight).The chemical properties of an element are determined primarily by the protonsand electrons in its atoms; neutrons do not take part in chemical changes under normalconditions. Therefore, isotopes of the same element have similar chemistries,forming the same types of compounds and displaying similar reactivities.Example 2.1 shows how to calculate the number of protons, neutrons, and electronsusing atomic numbers and mass numbers.1 1H 12 H 13 HReview of Concepts(a) Name the only element having an isotope that contains no neutrons.(b) Explain why a helium nucleus containing no neutrons is likely to be unstable.EXAMPLE 2.1Give the number of protons, neutrons, and electrons in each of the following species:(a) 2011Na, (b) 2211Na, (c) 17O, and (d) carbon-14.Strategy Recall that the superscript denotes the mass number ( A ) and the subscriptdenotes the atomic number ( Z ). Mass number is always greater than atomic number.(The only exception is 11H, where the mass number is equal to the atomic number.) Ina case where no subscript is shown, as in parts (c) and (d), the atomic number canbe deduced from the element symbol or name. To determine the number of electrons,remember that because atoms are electrically neutral, the number of electrons is equalto the number of protons.Solution (a) The atomic number is 11, so there are 11 protons. The mass number is20, so the number of neutrons is 20 2 11 5 9. The number of electrons is thesame as the number of protons; that is, 11.(b) The atomic number is the same as that in (a), or 11. The mass number is 22, so thenumber of neutrons is 22 2 11 5 11. The number of electrons is 11. Note that thespecies in (a) and (b) are chemically similar isotopes of sodium.(c) The atomic number of O (oxygen) is 8, so there are 8 protons. The mass number is17, so there are 17 2 8 5 9 neutrons. There are 8 electrons.(d) Carbon-14 can also be represented as 14C. The atomic number of carbon is 6, sothere are 14 2 6 5 8 neutrons. The number of electrons is 6.Practice Exercise How many protons, neutrons, and electrons are in the followingisotope of copper: 63Cu?Similar problems: 2.15, 2.16.2.4 The Periodic TableMore than half of the elements known today were discovered between 1800 and1900. During this period, chemists noted that many elements show strong similaritiesto one another. Recognition of periodic regularities in physical and chemicalbehavior and the need to organize the large volume of available informationabout the structure and properties of elemental substances led to the developmentof the periodic table, a chart in which elements having similar chemical and physicalproperties are grouped together. Figure 2.10 shows the modern periodic tablein which the elements are arranged by atomic number (shown above the elementsymbol) in horizontal rows called periods and in vertical columns known as groupsor families, according to similarities in their chemical properties. Note that elements112116 and 118 have recently been synthesized, although they have not yetbeen named.The elements can be divided into three categoriesmetals, nonmetals, and metalloids.A metal is a good conductor of heat and electricity while a nonmetal isusually a poor conductor of heat and electricity. A metalloid has properties that areintermediate between those of metals and nonmetals. Figure 2.10 shows that the2.4 The Periodic Table 51MetalsMetalloidsNonmetals1H3Li11Na19K37Rb55Cs87Fr20Ca38Sr56Ba88Ra21Sc39Y57La89Ac22Ti40Zr72Hf104Rf23V41Nb73Ta105Db24Cr42Mo74W106Sg25Mn43Tc75Re107Bh26Fe44Ru76Os108Hs27Co45Rh77Ir109Mt28Ni46Pd78Pt29Cu47Ag79Au30Zn48Cd80Hg31Ga49In81Tl32Ge50Sn82Pb33As51Sb83Bi34Se52Te84Po35Br53I85At36Kr54Xe86Rn13Al14Si15P16S17Cl18Ar5B6C7N8O9F10Ne2He4Be12Mg58Ce90Th59Pr91Pa60Nd92U61Pm93Np62Sm94Pu63Eu95Am64Gd96Cm65Tb97Bk66Dy98Cf67Ho99Es68Er100Fm69Tm101Md70Yb102No71Lu103Lr110Ds112 113 114 115 116 (117) 11811A22A33B44B55B66B77B98B111B8 10 122B133A144A155A166A177A188A111RgFigure 2.10 The modern periodic table. The elements are arranged according to the atomic numbers above their symbols. With theexception of hydrogen (H), nonmetals appear at the far right of the table. The two rows of metals beneath the main body of the tableare conventionally set apart to keep the table from being too wide. Actually, cerium (Ce) should follow lanthanum (La), and thorium (Th)should come right after actinium (Ac). The 118 group designation has been recommended by the International Union of Pure andApplied Chemistry (IUPAC) but is not yet in wide use. In this text, we use the standard U.S. notation for group numbers (1A8A and1B8B). No names have yet been assigned to elements 112116, and 118. Element 117 has not yet been synthesized.52C H E M I S T R Yin ActionThe majority of elements are naturally occurring. How arethese elements distributed on Earth, and which are essentialto living systems?Earths crust extends from the surface to a depth of about40 km (about 25 mi). Because of technical diffi culties, scientistshave not been able to study the inner portions of Earth as easilyas the crust. Nevertheless, it is believed that there is a solid coreconsisting mostly of iron at the center of Earth. Surrounding thecore is a layer called the mantle, which consists of hot fl uidcontaining iron, carbon, silicon, and sulfur.Of the 83 elements that are found in nature, 12 make up99.7 percent of Earths crust by mass. They are, in decreasingorder of natural abundance, oxygen (O), silicon (Si), aluminum(Al), iron (Fe), calcium (Ca), magnesium (Mg), sodium (Na),potassium (K), titanium (Ti), hydrogen (H), phosphorus (P),and manganese (Mn). In discussing the natural abundance of theDistribution of Elements on Earth and in Living Systemselements, we should keep in mind that (1) the elements are notevenly distributed throughout Earths crust, and (2) most elementsoccur in combined forms. These facts provide the basisfor most methods of obtaining pure elements from their compounds,as we will see in later chapters.The accompanying table lists the essential elements in thehuman body. Of special interest are the trace elements, such asiron (Fe), copper (Cu), zinc (Zn), iodine (I), and cobalt (Co),which together make up about 0.1 percent of the bodys mass.These elements are necessary for biological functions such asgrowth, transport of oxygen for metabolism, and defenseagainst disease. There is a delicate balance in the amounts ofthese elements in our bodies. Too much or too little over anextended period of time can lead to serious illness, retardation,or even death.2900 km 3480 kmCrustCoreMantleStructure of Earths interior.Essential Elements in the Human BodyElement Percent by Mass* Element Percent by Mass*Oxygen 65 Sodium 0.1Carbon 18 Magnesium 0.05Hydrogen 10 Iron ,0.05Nitrogen 3 Cobalt ,0.05Calcium 1.6 Copper ,0.05Phosphorus 1.2 Zinc ,0.05Potassium 0.2 Iodine ,0.05Sulfur 0.2 Selenium ,0.01Chlorine 0.2 Fluorine ,0.01*Percent by mass gives the mass of the element in grams present in a 100-g sample.(a) Natural abundance of the elementsin percent by mass. For example, oxygensabundance is 45.5 percent. Thismeans that in a 100-g sample of Earthscrust there are, on the average, 45.5 gof the element oxygen. (b) Abundanceof elements in the human body in percentby mass.Magnesium 2.8%Oxygen45.5% Oxygen65%Silicon27.2% Hydrogen 10%Carbon18%Calcium 4.7%All others 5.3%All others 1.2%Phosphorus 1.2%Calcium 1.6%Nitrogen 3%Iron 6.2%Aluminum 8.3%(a) (b)majority of known elements are metals; only 17 elements are nonmetals, and 8 elementsare metalloids. From left to right across any period, the physical and chemicalproperties of the elements change gradually from metallic to nonmetallic.Elements are often referred to collectively by their periodic table group number(Group 1A, Group 2A, and so on). However, for convenience, some element groupshave been given special names. The Group 1A elements (Li, Na, K, Rb, Cs, and Fr)are called alkali metals, and the Group 2A elements (Be, Mg, Ca, Sr, Ba, and Ra) arecalled alkaline earth metals . Elements in Group 7A (F, Cl, Br, I, and At) are knownas halogens, and elements in Group 8A (He, Ne, Ar, Kr, Xe, and Rn) are called noblegases, or rare gases .The periodic table is a handy tool that correlates the properties of the elementsin a systematic way and helps us to make predictions about chemical behavior. Wewill take a closer look at this keystone of chemistry in Chapter 8.The Chemistry in Action essay on p. 52 describes the distribution of the elementson Earth and in the human body.Review of ConceptsIn viewing the periodic table, do chemical properties change more markedlyacross a period or down a group?2.5 Molecules and IonsOf all the elements, only the six noble gases in Group 8A of the periodic table (He,Ne, Ar, Kr, Xe, and Rn) exist in nature as single atoms. For this reason, they arecalled monatomic (meaning a single atom) gases. Most matter is composed ofmolecules or ions formed by atoms.MoleculesA molecule is an aggregate of at least two atoms in a defi nite arrangement heldtogether by chemical forces (also called chemical bonds ). A molecule may containatoms of the same element or atoms of two or more elements joined in a fi xed ratio,in accordance with the law of defi nite proportions stated in Section 2.1. Thus, a moleculeis not necessarily a compound, which, by defi nition, is made up of two or moreelements (see Section 1.4). Hydrogen gas, for example, is a pure element, but it consistsof molecules made up of two H atoms each. Water, on the other hand, is amolecular compound that contains hydrogen and oxygen in a ratio of two H atomsand one O atom. Like atoms, molecules are electrically neutral.The hydrogen molecule, symbolized as H2, is called a diatomic molecule becauseit contains only two atoms. Other elements that normally exist as diatomic moleculesare nitrogen (N2) and oxygen (O2), as well as the Group 7A elementsfl uorine (F2),chlorine (Cl2), bromine (Br2), and iodine (I2). Of course, a diatomic molecule cancontain atoms of different elements. Examples are hydrogen chloride (HCl) and carbonmonoxide (CO).The vast majority of molecules contain more than two atoms. They can be atomsof the same element, as in ozone (O3), which is made up of three atoms of oxygen,or they can be combinations of two or more different elements. Molecules containingmore than two atoms are called polyatomic molecules . Like ozone, water (H2O) andammonia (NH3) are polyatomic molecules.2.5 Molecules and Ions 53We will discuss the nature of chemicalbonds in Chapters 9 and 10.1A2A 3A 4A 5A 6A 7A8AN O FClBrIHElements that exist as diatomic molecules.54 Atoms, Molecules, and IonsIonsAn ion is an atom or a group of atoms that has a net positive or negative charge.The number of positively charged protons in the nucleus of an atom remains the sameduring ordinary chemical changes (called chemical reactions), but negatively chargedelectrons may be lost or gained. The loss of one or more electrons from a neutralatom results in a cation, an ion with a net positive charge. For example, a sodiumatom (Na) can readily lose an electron to become a sodium cation, which is representedby Na1:Na Atom Na1 Ion11 protons 11 protons11 electrons 10 electronsOn the other hand, an anion is an ion whose net charge is negative due to an increasein the number of electrons. A chlorine atom (Cl), for instance, can gain an electronto become the chloride ion Cl2:Cl Atom Cl2 Ion17 protons 17 protons17 electrons 18 electronsSodium chloride (NaCl), ordinary table salt, is called an ionic compound because itis formed from cations and anions.An atom can lose or gain more than one electron. Examples of ions formed bythe loss or gain of more than one electron are Mg21, Fe31, S22, and N32. These ions,as well as Na1 and Cl2, are called monatomic ions because they contain only oneatom. Figure 2.11 shows the charges of a number of monatomic ions. With very fewexceptions, metals tend to form cations and nonmetals form anions.In addition, two or more atoms can combine to form an ion that has a net positiveor net negative charge. Polyatomic ions such as OH2 (hydroxide ion), CN2(cyanide ion), and NH14(ammonium ion) are ions containing more than one atom.11A22A33B44B55B66B77B98B111B8 10 122B133A144A155A166A177A188ALi+Na+K+Rb+Cs+Ca2+Sr2+Ba2+Ag+Ni Zn2+ Se2 Br 2+Ni3+Mn2+Mn3+Cr2+Cr3+Cd2+ Te2 IAl3+ S2 ClC4 N3 O2 FMg2+Fe2+Fe3+Co2+Co3+Cu+Cu2+P3Sn2+Sn4+Pb2+Pb4+Hg2+Hg2+Au+Au3+2Figure 2.11 Common monatomic ions arranged according to their positions in the periodic table. Note that the Hg221 ion containstwo atoms.In Chapter 8, we will see why atoms ofdifferent elements gain (or lose) a specifi cnumber of electrons.2.6 Chemical FormulasChemists use chemical formulas to express the composition of molecules and ioniccompounds in terms of chemical symbols. By composition we mean not only the elementspresent but also the ratios in which the atoms are combined. Here we areconcerned with two types of formulas: molecular formulas and empirical formulas.Molecular FormulasA molecular formula shows the exact number of atoms of each element in the smallestunit of a substance. In our discussion of molecules, each example was given withits molecular formula in parentheses. Thus, H2 is the molecular formula for hydrogen,O2 is oxygen, O3 is ozone, and H2O is water. The subscript numeral indicates thenumber of atoms of an element present. There is no subscript for O in H2O becausethere is only one atom of oxygen in a molecule of water, and so the number oneis omitted from the formula. Note that oxygen (O2) and ozone (O3) are allotropes ofoxygen. An allotrope is one of two or more distinct forms of an element. Two allotropicforms of the element carbondiamond and graphiteare dramatically differentnot only in properties but also in their relative cost.Molecular ModelsMolecules are too small for us to observe directly. An effective means of visualizingthem is by the use of molecular models. Two standard types of molecular models arecurrently in use: ball-and-stick models and space-fi lling models ( Figure 2.12 ). In balland-stick model kits, the atoms are wooden or plastic balls with holes in them. Sticksor springs are used to represent chemical bonds. The angles they form between atomsapproximate the bond angles in actual molecules. With the exception of the H atom,the balls are all the same size and each type of atom is represented by a specifi c color.In space-fi lling models, atoms are represented by truncated balls held together by snap2.6 Chemical Formulas 55MolecularformulaStructuralformulaBall-and-stickmodelSpace-fillingmodelHydrogenH2HHWaterH2OHOHAmmoniaNH3HNHWHMethaneCH4HWHCHWHFigure 2.12 Molecular and structural formulas and molecular models of four common molecules.See back endpaper for color codes foratoms.56 Atoms, Molecules, and Ionsfasteners, so that the bonds are not visible. The balls are proportional in size to atoms.The fi rst step toward building a molecular model is writing the structural formula,which shows how atoms are bonded to one another in a molecule. For example, it isknown that each of the two H atoms is bonded to an O atom in the water molecule.Therefore, the structural formula of water is H}O}H. A line connecting the twoatomic symbols represents a chemical bond.Ball-and-stick models show the three-dimensional arrangement of atoms clearly,and they are fairly easy to construct. However, the balls are not proportional to thesize of atoms. Furthermore, the sticks greatly exaggerate the space between atoms ina molecule. Space-fi lling models are more accurate because they show the variationin atomic size. Their drawbacks are that they are time-consuming to put together andthey do not show the three-dimensional positions of atoms very well. We will useboth models extensively in this text.Empirical FormulasThe molecular formula of hydrogen peroxide, a substance used as an antiseptic and asa bleaching agent for textiles and hair, is H2O2. This formula indicates that each hydrogenperoxide molecule consists of two hydrogen atoms and two oxygen atoms. The ratio ofhydrogen to oxygen atoms in this molecule is 2:2 or 1:1. The empirical formula ofhydrogen peroxide is HO. Thus, the empirical formula tells us which elements are presentand the simplest whole-number ratio of their atoms, but not necessarily the actualnumber of atoms in a given molecule. As another example, consider the compoundhydrazine (N2H4), which is used as a rocket fuel. The empirical formula of hydrazine isNH2. Although the ratio of nitrogen to hydrogen is 1:2 in both the molecular formula(N2H4) and the empirical formula (NH2), only the molecular formula tells us the actualnumber of N atoms (two) and H atoms (four) present in a hydrazine molecule.Empirical formulas are the simplest chemical formulas; they are written by reducingthe subscripts in the molecular formulas to the smallest possible whole numbers.Molecular formulas are the true formulas of molecules. If we know the molecularformula, we also know the empirical formula, but the reverse is not true. Why, then,do chemists bother with empirical formulas? As we will see in Chapter 3, when chemistsanalyze an unknown compound, the fi rst step is usually the determination of thecompounds empirical formula. With additional information, it is possible to deducethe molecular formula.For many molecules, the molecular formula and the empirical formula are oneand the same. Some examples are water (H2O), ammonia (NH3), carbon dioxide(CO2), and methane (CH4).Examples 2.2 and 2.3 deal with writing molecular formulas from molecular modelsand writing empirical formulas from molecular formulas.H2O2The word empirical means derived fromexperiment. As we will see in Chapter 3,empirical formulas are determinedexperimentally.Similar problems: 2.47, 2.48.OHCMethanolEXAMPLE 2.2Write the molecular formula of methanol, an organic solvent and antifreeze, from itsball-and-stick model, shown in the margin.Solution Refer to the labels (also see back endpapers). There are four H atoms, oneC atom, and one O atom. Therefore, the molecular formula is CH4O. However, thestandard way of writing the molecular formula for methanol is CH3OH because it showshow the atoms are joined in the molecule.Practice Exercise Write the molecular formula of chloroform, which is used as asolvent and a cleansing agent. The ball-and-stick model of chloroform is shown in themargin on p. 57.Formula of Ionic CompoundsThe formulas of ionic compounds are usually the same as their empirical formulasbecause ionic compounds do not consist of discrete molecular units. For example, asolid sample of sodium chloride (NaCl) consists of equal numbers of Na1 and Cl2ions arranged in a three-dimensional network ( Figure 2.13 ). In such a compound thereis a 1:1 ratio of cations to anions so that the compound is electrically neutral. As youcan see in Figure 2.13 , no Na1 ion in NaCl is associated with just one particular Cl2ion. In fact, each Na1 ion is equally held by six surrounding Cl2 ions and vice versa.Thus, NaCl is the empirical formula for sodium chloride. In other ionic compounds,the actual structure may be different, but the arrangement of cations and anions issuch that the compounds are all electrically neutral. Note that the charges on thecation and anion are not shown in the formula for an ionic compound.ClHCChloroformEXAMPLE 2.3Write the empirical formulas for the following molecules: (a) acetylene (C2H2), which isused in welding torches; (b) glucose (C6H12O6), a substance known as blood sugar; and(c) nitrous oxide (N2O), a gas that is used as an anesthetic gas (laughing gas) and asan aerosol propellant for whipped creams.Strategy Recall that to write the empirical formula, the subscripts in the molecularformula must be converted to the smallest possible whole numbers.Solution(a) There are two carbon atoms and two hydrogen atoms in acetylene. Dividing thesubscripts by 2, we obtain the empirical formula CH.(b) In glucose there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms.Dividing the subscripts by 6, we obtain the empirical formula CH2O. Note that ifwe had divided the subscripts by 3, we would have obtained the formula C2H4O2.Although the ratio of carbon to hydrogen to oxygen atoms in C2H4O2 is the same asthat in C6H12O6 (1:2:1), C2H4O2 is not the simplest formula because its subscriptsare not in the smallest whole-number ratio.(c) Because the subscripts in N2O are already the smallest possible whole numbers, theempirical formula for nitrous oxide is the same as its molecular formula.Practice Exercise Write the empirical formula for caffeine (C8H10N4O2), a stimulantfound in tea and coffee.Similar problems: 2.45, 2.46.Sodium metal reacting with chlorine gas toform sodium chloride.2.6 Chemical Formulas 57(a) (b) (c)Figure 2.13 (a) Structure of solid NaCl. (b) In reality, the cations are in contact with the anions. In both (a) and (b), the smaller spheresrepresent Na1 ions and the larger spheres, Cl2 ions. (c) Crystals of NaCl.58 Atoms, Molecules, and IonsFor ionic compounds to be electrically neutral, the sum of the charges on the cationand anion in each formula unit must be zero. If the charges on the cation and anion arenumerically different, we apply the following rule to make the formula electrically neutral:The subscript of the cation is numerically equal to the charge on the anion, and thesubscript of the anion is numerically equal to the charge on the cation. If the chargesare numerically equal, then no subscripts are necessary. This rule follows from the factthat because the formulas of ionic compounds are usually empirical formulas, the subscriptsmust always be reduced to the smallest ratios. Let us consider some examples. Potassium Bromide. The potassium cation K1 and the bromine anion Br2 combineto form the ionic compound potassium bromide. The sum of the charges is11 1 (21) 5 0, so no subscripts are necessary. The formula is KBr. Zinc Iodide. The zinc cation Zn21 and the iodine anion I2 combine to form zinciodide. The sum of the charges of one Zn21 ion and one I2 ion is 12 1 (21) 511. To make the charges add up to zero we multiply the 21 charge of the anionby 2 and add the subscript 2 to the symbol for iodine. Therefore the formulafor zinc iodide is ZnI2. Aluminum Oxide. The cation is Al31 and the oxygen anion is O22. The followingdiagram helps us determine the subscripts for the compound formed by thecation and the anion:Al2 O3Al 3 _ O 2 _The sum of the charges is 2(13) 1 3(22) 5 0. Thus, the formula for aluminumoxide is Al2O3.Note that in each of the above three examples,the subscripts are in the smallest ratios.When magnesium burns in air, it forms bothmagnesium oxide and magnesium nitride.EXAMPLE 2.4Write the formula of magnesium nitride, containing the Mg21 and N32 ions.Strategy Our guide for writing formulas for ionic compounds is electrical neutrality;that is, the total charge on the cation(s) must be equal to the total charge on theanion(s). Because the charges on the Mg21 and N32 ions are not equal, we know theformula cannot be MgN. Instead, we write the formula as MgxNy, where x and y aresubscripts to be determined.Solution To satisfy electrical neutrality, the following relationship must hold:(12)x 1 (23)y 5 0Solving, we obtain x/y 5 3/2. Setting x 5 3 and y 5 2, we writeMg3 N2Mg N 2 _ 3 _Check The subscripts are reduced to the smallest whole number ratio of the atomsbecause the chemical formula of an ionic compound is usually its empirical formula.Practice Exercise Write the formulas of the following ionic compounds: (a) chromiumsulfate (containing the Cr31 and SO422 ions) and (b) titanium oxide (containing the Ti41and O22 ions).Refer to Figure 2.11 for charges of cationsand anions.Similar problems: 2.43, 2.44.2.7 Naming CompoundsWhen chemistry was a young science and the number of known compounds was small,it was possible to memorize their names. Many of the names were derived from theirphysical appearance, properties, origin, or applicationfor example, milk of magnesia,laughing gas, limestone, caustic soda, lye, washing soda, and baking soda.Today the number of known compounds is well over 20 million. Fortunately, itis not necessary to memorize their names. Over the years chemists have devised aclear system for naming chemical substances. The rules are accepted worldwide,facilitating communication among chemists and providing a useful way of labelingan overwhelming variety of substances. Mastering these rules now will prove benefi -cial almost immediately as we proceed with our study of chemistry.To begin our discussion of chemical nomenclature, the naming of chemical compounds,we must fi rst distinguish between inorganic and organic compounds. Organiccompounds contain carbon, usually in combination with elements such as hydrogen,oxygen, nitrogen, and sulfur. All other compounds are classifi ed as inorganic compounds.For convenience, some carbon-containing compounds, such as carbon monoxide(CO), carbon dioxide (CO2), carbon disulfi de (CS2), compounds containing thecyanide group (CN2), and carbonate (CO322) and bicarbonate (HCO32) groups areconsidered to be inorganic compounds. Section 2.8 gives a brief introduction toorganic compounds.To organize and simplify our venture into naming compounds, we can divideinorganic compounds into four categories: ionic compounds, molecular compounds,acids and bases, and hydrates.Ionic CompoundsIn Section 2.5 we learned that ionic compounds are made up of cations (positive ions)and anions (negative ions). With the important exception of the ammonium ion, NH41,all cations of interest to us are derived from metal atoms. Metal cations take theirnames from the elements. For example,For names and symbols of the elements,see front end papers.1A2A 3A 4A 5A 6A 7A8AN OAl SFClBrILiNaKRbCsMgCaSrBaThe most reactive metals (green) and themost reactive nonmetals (blue) combine toform ionic compounds.Review of ConceptsMatch each of the diagrams shown here with the following ionic compounds:Al2O3, LiH, Na2S, Mg(NO3)2. (Green spheres represent cations and red spheresrepresent anions.)(a) (b) (c) (d)2.7 Naming Compounds 59Element Name of CationNa sodium Na1 sodium ion (or sodium cation)K potassium K1 potassium ion (or potassium cation)Mg magnesium Mg21 magnesium ion (or magnesium cation)Al aluminum Al31 aluminum ion (or aluminum cation)Many ionic compounds are binary compounds, or compounds formed from justtwo elements. For binary compounds, the fi rst element named is the metal cation,followed by the nonmetallic anion. Thus, NaCl is sodium chloride. The anion is namedMedia PlayerFormation of an Ionic Compound60 Atoms, Molecules, and Ionsby taking the fi rst part of the element name (chlorine) and adding -ide. Potassiumbromide (KBr), zinc iodide (ZnI2), and aluminum oxide (Al2O3) are also binary compounds.Table 2.2 shows the -ide nomenclature of some common monatomic anionsaccording to their positions in the periodic table.The -ide ending is also used for certain anion groups containing different elements,such as hydroxide (OH2) and cyanide (CN2). Thus, the compounds LiOH andKCN are named lithium hydroxide and potassium cyanide, respectively. These and anumber of other such ionic substances are called ternary compounds, meaning compoundsconsisting of three elements. Table 2.3 lists alphabetically the names of anumber of common cations and anions.Certain metals, especially the transition metals, can form more than one type of cation.Take iron as an example. Iron can form two cations: Fe21 and Fe31. An older nomenclaturesystem that is still in limited use assigns the ending -ous to the cation with fewerpositive charges and the ending -ic to the cation with more positive charges:Fe21 ferrous ionFe31 ferric ionThe names of the compounds that these iron ions form with chlorine would thus beFeCl2 ferrous chlorideFeCl3 ferric chlorideThis method of naming ions has some distinct limitations. First, the -ous and -icsuffi xes do not provide information regarding the actual charges of the two cationsinvolved. Thus, the ferric ion is Fe31, but the cation of copper named cupric hasthe formula Cu21. In addition, the -ous and -ic designations provide names foronly two different elemental cations. Some metallic elements can assume three ormore different positive charges in compounds. Therefore, it has become increasinglycommon to designate different cations with Roman numerals. This is called theStock system. In this system, the Roman numeral I indicates one positive charge,II means two positive charges, and so on. For example, manganese (Mn) atoms canassume several different positive charges:Mn21: MnO manganese(II) oxideMn31: Mn2O3 manganese(III) oxideMn41: MnO2 manganese(IV) oxideThese names are pronounced manganese-two oxide, manganese-three oxide, andmanganese-four oxide. Using the Stock system, we denote the ferrous ion and theGroup 4A Group 5A Group 6A Group 7AC carbide (C42)* N nitride (N32) O oxide (O22) F fl uoride (F2)Si silicide (Si42) P phosphide (P32) S sulfi de (S22) Cl chloride (Cl2)Se selenide (Se22) Br bromide (Br2)Te telluride (Te22) I iodide (I2)*The word carbide is also used for the anion C2 22.TABLE 2.2 The -ide Nomenclature of Some Common Monatomic AnionsAccording to Their Positions in the Periodic TableAlfred E. Stock (18761946). German chemist. Stock did most of his research in the synthesis and characterizationof boron, beryllium, and silicon compounds. He was the fi rst scientist to explore the dangersof mercury poisoning.3B 4B 5B 6B 7B 8B 1B 2BThe transition metals are the elements inGroups 1B and 3B8B (see Figure 2.10).FeCl2 (left) and FeCl3 (right).Keep in mind that the Roman numeralsrefer to the charges on the metal cations.Cation Anionaluminum (Al31) bromide (Br2)ammonium (NH14) carbonate (CO3 22)barium (Ba21) chlorate (ClO32)cadmium (Cd21) chloride (Cl2)calcium (Ca21) chromate (CrO4 22)cesium (Cs1) cyanide (CN2)chromium(III) or chromic (Cr31) dichromate (Cr2O7 22)cobalt(II) or cobaltous (Co21) dihydrogen phosphate (H2PO42)copper(I) or cuprous (Cu1) fl uoride (F2)copper(II) or cupric (Cu21) hydride (H2)hydrogen (H1) hydrogen carbonate or bicarbonate (HCO32)iron(II) or ferrous (Fe21) hydrogen phosphate (HPO4 22)iron(III) or ferric (Fe31) hydrogen sulfate or bisulfate (HSO42)lead(II) or plumbous (Pb21) hydroxide (OH2)lithium (Li1) iodide (I2)magnesium (Mg21) nitrate (NO32)manganese(II) or manganous (Mn21) nitride (N32)mercury(I) or mercurous (Hg2 21)* nitrite (NO22)mercury(II) or mercuric (Hg21) oxide (O22)potassium (K1) permanganate (MnO42)rubidium (Rb1) peroxide (O2 22)silver (Ag1) phosphate (PO4 32)sodium (Na1) sulfate (SO4 22)strontium (Sr21) sulfi de (S22)tin(II) or stannous (Sn21) sulfi te (SO3 22)zinc (Zn21) thiocyanate (SCN2)*Mercury(I) exists as a pair as shown.TABLE 2.3 Names and Formulas of Some Common Inorganic Cationsand Anions2.7 Naming Compounds 61ferric ion as iron(II) and iron(III), respectively; ferrous chloride becomes iron(II)chloride; and ferric chloride is called iron(III) chloride. In keeping with modern practice,we will favor the Stock system of naming compounds in this textbook.Examples 2.5 and 2.6 illustrate how to name ionic compounds and write formulasfor ionic compounds based on the information given in Figure 2.11 and Tables 2.2and 2.3 .EXAMPLE 2.5Name the following compounds: (a) Cu(NO3)2, (b) KH2PO4, and (c) NH4ClO3.Strategy Note that the compounds in (a) and (b) contain both metal and nonmetalatoms, so we expect them to be ionic compounds. There are no metal atoms in (c) butthere is an ammonium group, which bears a positive charge. So NH4ClO3 is also an(Continued)62 Atoms, Molecules, and IonsSimilar problems: 2.57(b), (e), (f).EXAMPLE 2.6Write chemical formulas for the following compounds: (a) mercury(I) nitrite, (b) cesiumsulfi de, and (c) calcium phosphate.Strategy We refer to Table 2.3 for the formulas of cations and anions. Recall that theRoman numerals in the Stock system provide useful information about the charges ofthe cation.Solution(a) The Roman numeral shows that the mercury ion bears a 11 charge. According toTable 2.3 , however, the mercury(I) ion is diatomic (that is, Hg2 21) and the nitrite ionis NO22. Therefore, the formula is Hg2(NO2)2.(b) Each sulfi de ion bears two negative charges, and each cesium ion bears one positivecharge (cesium is in Group 1A, as is sodium). Therefore, the formula is Cs2S.(c) Each calcium ion (Ca21) bears two positive charges, and each phosphate ion (PO4 32)bears three negative charges. To make the sum of the charges equal zero, we mustadjust the numbers of cations and anions:3(12) 1 2(23) 5 0Thus, the formula is Ca3(PO4)2.Practice Exercise Write formulas for the following ionic compounds: (a) rubidiumsulfate and (b) barium hydride.Note that the subscripts of this ioniccompound are not reduced to the smallestratio because the Hg(I) ion exists as a pairor dimer.Similar problems: 2.59(a), (b), (d), (h), (i).ionic compound. Our reference for the names of cations and anions is Table 2.3 . Keepin mind that if a metal atom can form cations of different charges (see Figure 2.11 ), weneed to use the Stock system.Solution(a) The nitrate ion (NO32) bears one negative charge, so the copper ion must have twopositive charges. Because copper forms both Cu1 and Cu21 ions, we need to usethe Stock system and call the compound copper(II) nitrate.(b) The cation is K1 and the anion is H2PO42 (dihydrogen phosphate). Becausepotassium only forms one type of ion (K1), there is no need to use potassium(I) inthe name. The compound is potassium dihydrogen phosphate.(c) The cation is NH41 (ammonium ion) and the anion is ClO32. The compound isammonium chlorate.Practice Exercise Name the following compounds: (a) PbO and (b) Li2SO3.Molecular CompoundsUnlike ionic compounds, molecular compounds contain discrete molecular units. Theyare usually composed of nonmetallic elements (see Figure 2.10 ). Many molecularcompounds are binary compounds. Naming binary molecular compounds is similar tonaming binary ionic compounds. We place the name of the fi rst element in the formulafi rst, and the second element is named by adding -ide to the root of the element name.Some examples areHCl hydrogen chlorideHBr hydrogen bromideSiC silicon carbideIt is quite common for one pair of elements to form several different compounds.In these cases, confusion in naming the compounds is avoided by the use of Greekprefi xes to denote the number of atoms of each element present ( Table 2.4 ). Considerthe following examples:CO carbon monoxideCO2 carbon dioxideSO2 sulfur dioxideSO3 sulfur trioxideNO2 nitrogen dioxideN2O4 dinitrogen tetroxideThe following guidelines are helpful in naming compounds with prefi xes: The prefi x mono- may be omitted for the fi rst element. For example, PCl3 isnamed phosphorus trichloride, not monophosphorus trichloride. Thus, the absenceof a prefi x for the fi rst element usually means there is only one atom of that elementpresent in the molecule. For oxides, the ending a in the prefi x is sometimes omitted. For example, N2O4may be called dinitrogen tetroxide rather than dinitrogen tetraoxide.Exceptions to the use of Greek prefi xes are molecular compounds containinghydrogen. Traditionally, many of these compounds are called either by their common,nonsystematic names or by names that do not specifi cally indicate the number of Hatoms present:B2H6 diboraneCH4 methaneSiH4 silaneNH3 ammoniaPH3 phosphineH2O waterH2S hydrogen sulfi deNote that even the order of writing the elements in the formulas for hydrogen compoundsis irregular. In water and hydrogen sulfi de, H is written fi rst, whereas it appearslast in the other compounds.Writing formulas for molecular compounds is usually straightforward. Thus, thename arsenic trifl uoride means that there are three F atoms and one As atom in eachmolecule, and the molecular formula is AsF3. Note that the order of elements in theformula is the same as in its name.Binary compounds containing carbon andhydrogen are organic compounds; they donot follow the same naming conventions.We will discuss the naming of organiccompounds in Chapter 24.2.7 Naming Compounds 63EXAMPLE 2.7Name the following molecular compounds: (a) SiCl4 and (b) P4O10.Strategy We refer to Table 2.4 for prefi xes. In (a) there is only one Si atom so we donot use the prefi x mono.Solution (a) Because there are four chlorine atoms present, the compound is silicontetrachloride.(b) There are four phosphorus atoms and ten oxygen atoms present, so the compound istetraphosphorus decoxide. Note that the a is omitted in deca.Practice Exercise Name the following molecular compounds: (a) NF3 and (b) Cl2O7.Similar problems: 2.57(c), (i), (j).Greek Prefi xes Used inNaming MolecularCompoundsTABLE 2.4Prefi x Meaningmono- 1di- 2tri- 3tetra- 4penta- 5hexa- 6hepta- 7octa- 8nona- 9deca- 1064 Atoms, Molecules, and IonsFigure 2.14 summarizes the steps for naming ionic and binary molecularcompounds.EXAMPLE 2.8Write chemical formulas for the following molecular compounds: (a) carbon disulfi deand (b) disilicon hexabromide.Strategy Here we need to convert prefi xes to numbers of atoms (see Table 2.4 ). Becausethere is no prefi x for carbon in (a), it means that there is only one carbon atom present.Solution (a) Because there are two sulfur atoms and one carbon atom present, theformula is CS2.(b) There are two silicon atoms and six bromine atoms present, so the formula is Si2Br6.Practice Exercise Write chemical formulas for the following molecular compounds:(a) sulfur tetrafl uoride and (b) dinitrogen pentoxide.Cation hasonly one chargeIonicCation: metal or NH4+Anion: monatomic orpolyatomic Binary compoundsof nonmetalsCompoundMolecularCation has morethan one chargeNamingNamingNaming Alkali metal cations Alkaline earth metal cations Ag+, Al3+, Cd2+, Zn2+ Other metal cations Name metal first If monatomic anion,add ide to theroot of the elementname If polyatomic anion,use name of anion(see Table 2.3) Name metal first Specify charge ofmetal cation withRoman numeralin parentheses If monatomic anion,add ide to theroot of the elementname If polyatomic anion,use name of anion(see Table 2.3) Use prefixes forboth elements present(Prefix monousually omitted forthe first element) Add ide to theroot of the secondelementFigure 2.14 Steps for naming ionic and binary molecular compounds.Similar problems: 2.59(g), (j).2.7 Naming Compounds 65Anion Corresponding AcidF2 (fl uoride) HF (hydrofl uoric acid)Cl2 (chloride) HCl (hydrochloric acid)Br2 (bromide) HBr (hydrobromic acid)I2 (iodide) HI (hydroiodic acid)CN2 (cyanide) HCN (hydrocyanic acid)S22 (sulfi de) H2S (hydrosulfuric acid)TABLE 2.5 Some Simple AcidsNote that these acids all exist as molecularcompounds in the gas phase.Acids and BasesNaming AcidsAn acid can be described as a substance that yields hydrogen ions (H1) when dissolvedin water. (H1 is equivalent to one proton, and is often referred to that way.)Formulas for acids contain one or more hydrogen atoms as well as an anionic group.Anions whose names end in -ide form acids with a hydro- prefi x and an -icending, as shown in Table 2.5 . In some cases two different names seem to be assignedto the same chemical formula.HCl hydrogen chlorideHCl hydrochloric acidThe name assigned to the compound depends on its physical state. In the gaseous orpure liquid state, HCl is a molecular compound called hydrogen chloride. When it isdissolved in water, the molecules break up into H1 and Cl2 ions; in this state, thesubstance is called hydrochloric acid.Oxoacids are acids that contain hydrogen, oxygen, and another element (the centralelement). The formulas of oxoacids are usually written with the H fi rst, followedby the central element and then O. We use the following fi ve common acids as ourreferences in naming oxoacids:H2CO3 carbonic acidHClO3 chloric acidHNO3 nitric acidH3PO4 phosphoric acidH2SO4 sulfuric acidOften two or more oxoacids have the same central atom but a different number of Oatoms. Starting with our reference oxoacids whose names all end with -ic, we usethe following rules to name these compounds.1. Addition of one O atom to the -ic acid: The acid is called per . . . -ic acid.Thus, adding an O atom to HClO3 changes chloric acid to perchloric acid,HClO4.2. Removal of one O atom from the -ic acid: The acid is called -ous acid. Thus,nitric acid, HNO3, becomes nitrous acid, HNO2.3. Removal of two O atoms from the -ic acid: The acid is called hypo . . . -ousacid. Thus, when HBrO3 is converted to HBrO, the acid is called hypobromousacid.ClHClH3O+When dissolved in water, the HCl moleculeis converted to the H1 and Cl2 ions.The H1 ion is associated with one ormore water molecules, and is usuallyrepresented as H3O1.OHNHNO3OHCH2CO366 Atoms, Molecules, and IonsThe rules for naming oxoanions, anions of oxoacids, are as follows:1. When all the H ions are removed from the -ic acid, the anions name ends with-ate. For example, the anion CO322 derived from H2CO3 is called carbonate.2. When all the H ions are removed from the -ous acid, the anions name endswith -ite. Thus, the anion ClO22 derived from HClO2 is called chlorite.3. The names of anions in which one or more but not all the hydrogen ions havebeen removed must indicate the number of H ions present. For example, considerthe anions derived from phosphoric acid:H3PO4 phosphoric acidH2PO42 dihydrogen phosphateHPO422 hydrogen phosphatePO432 phosphateNote that we usually omit the prefi x mono- when there is only one H in theanion. Figure 2.15 summarizes the nomenclature for the oxoacids and oxoanions, andTable 2.6 gives the names of the oxoacids and oxoanions that contain chlorine.OPHH3PO4hypo ous acidous acidReference ic acidper ic aciditeateper ateRemoval ofall H+ ionshypo iteOxoacid Oxoanion+[O][O][O]Figure 2.15 Naming oxoacidsand oxoanions.Acid AnionHClO4 (perchloric acid) ClO42 (perchlorate)HClO3 (chloric acid) ClO32 (chlorate)HClO2 (chlorous acid) ClO22 (chlorite)HClO (hypochlorous acid) ClO2 (hypochlorite)TABLE 2.6 Names of Oxoacids and Oxoanions That Contain ChlorineExample 2.9 deals with the nomenclature for an oxoacid and an oxoanion.2.7 Naming Compounds 67EXAMPLE 2.9Name the following oxoacid and oxoanion: (a) H3PO3 and (b) IO42.Strategy To name the acid in (a), we fi rst identify the reference acid, whose nameends with ic, as shown in Figure 2.15 . In (b), we need to convert the anion to itsparent acid shown in Table 2.6 .Solution (a) We start with our reference acid, phosphoric acid (H3PO4). BecauseH3PO3 has one fewer O atom, it is called phosphorous acid.(b) The parent acid is HIO4. Because the acid has one more O atom than our referenceiodic acid (HIO3), it is called periodic acid. Therefore, the anion derived from HIO4is called periodate.Practice Exercise Name the following oxoacid and oxoanion: (a) HBrO and(b) HSO42.Similar problem: 2.58(f).Naming BasesA base can be described as a substance that yields hydroxide ions (OH2) when dissolvedin water. Some examples areNaOH sodium hydroxideKOH potassium hydroxideBa(OH)2 barium hydroxideAmmonia (NH3), a molecular compound in the gaseous or pure liquid state, isalso classifi ed as a common base. At fi rst glance this may seem to be an exceptionto the defi nition of a base. But note that as long as a substance yields hydroxideions when dissolved in water, it need not contain hydroxide ions in its structureto be considered a base. In fact, when ammonia dissolves in water, NH3 reactspartially with water to yield NH41 and OH2 ions. Thus, it is properly classifi ed asa base.HydratesHydrates are compounds that have a specifi c number of water molecules attached tothem. For example, in its normal state, each unit of copper(II) sulfate has fi ve watermolecules associated with it. The systematic name for this compound is copper(II)sulfate pentahydrate, and its formula is written as CuSO4 ? 5H2O. The water moleculescan be driven off by heating. When this occurs, the resulting compound is CuSO4,which is sometimes called anhydrous copper(II) sulfate; anhydrous means that thecompound no longer has water molecules associated with it ( Figure 2.16 ). Some otherhydrates areBaCl2 ? 2H2O barium chloride dihydrateLiCl ? H2O lithium chloride monohydrateMgSO4 ? 7H2O magnesium sulfate heptahydrateSr(NO3)2 ? 4H2O strontium nitrate tetrahydrate68 Atoms, Molecules, and IonsFamiliar Inorganic CompoundsSome compounds are better known by their common names than by their systematicchemical names. Familiar examples are listed in Table 2.7 .2.8 Introduction to Organic CompoundsThe simplest type of organic compounds is the hydrocarbons, which contain onlycarbon and hydrogen atoms. The hydrocarbons are used as fuels for domestic andindustrial heating, for generating electricity and powering internal combustion engines,and as starting materials for the chemical industry. One class of hydrocarbons is calledthe alkanes. Table 2.8 shows the names, formulas, and molecular models of the fi rstten straight-chain alkanes, in which the carbon chains have no branches. Note that allthe names end with - ane . Starting with C5H12, we use the Greek prefi xes in Table 2.4to indicate the number of carbon atoms present.The chemistry of organic compounds is largely determined by the functionalgroups, which consist of one or a few atoms bonded in a specifi c way. For example,when an H atom in methane is replaced by a hydroxyl group (}OH), an aminoFigure 2.16 CuSO4 ? 5H2O (left)is blue; CuSO4 (right) is white.Formula Common Name Systematic NameH2O Water Dihydrogen monoxideNH3 Ammonia Trihydrogen nitrideCO2 Dry ice Solid carbon dioxideNaCl Table salt Sodium chlorideN2O Laughing gas Dinitrogen monoxideCaCO3 Marble, chalk, limestone Calcium carbonateCaO Quicklime Calcium oxideCa(OH)2 Slaked lime Calcium hydroxideNaHCO3 Baking soda Sodium hydrogen carbonateNa2CO3 ? 10H2O Washing soda Sodium carbonate decahydrateMgSO4 ? 7H2O Epsom salt Magnesium sulfate heptahydrateMg(OH)2 Milk of magnesia Magnesium hydroxideCaSO4 ? 2H2O Gypsum Calcium sulfate dihydrateTABLE 2.7 Common and Systematic Names of Some CompoundsCH3OHCH3NH2CH3COOHName Formula Molecular ModelMethane CH4Ethane C2H6Propane C3H8Butane C4H10Pentane C5H12Hexane C6H14Heptane C7H16Octane C8H18Nonane C9H20Decane C10H22TABLE 2.8 The First Ten Straight-Chain Alkanesgroup (}NH2), and a carboxyl group (}COOH), the following molecules aregenerated:HH C OHHMethanolHH NH2 CHMethylamineHH C C OHH OAcetic acid2.8 Introduction to Organic Compounds 6970 Atoms, Molecules, and IonsThe chemical properties of these molecules can be predicted based on the reactivityof the functional groups. Although the nomenclature of the major classes of organiccompounds and their properties in terms of the functional groups will not be discusseduntil Chapter 24, we will frequently use organic compounds as examples to illustratechemical bonding, acid-base reactions, and other properties throughout the book.Acid, p. 65Alkali metals, p. 53Alkaline earth metals, p. 53Allotrope, p. 55Alpha (a) particles, p. 46Alpha (a) rays, p. 46Anion, p. 54Atom, p. 43Atomic number ( Z ), p. 49Base, p. 67Beta (b) particles, p. 46Beta (b) rays, p. 46Binary compound, p. 59Cation, p. 54Chemical formula, p. 55Diatomic molecule, p. 53Electron, p. 44Empirical formula, p. 56Families, p. 51Gamma (g) rays, p. 46Groups, p. 51Halogens, p. 53Hydrate, p. 67Inorganiccompounds, p. 59Ion, p. 54Ionic compound, p. 54Isotope, p. 49Key WordsKey Equationmass number 5 number of protons 1 number of neutrons5 atomic number 1 number of neutrons (2. 1)1. Modern chemistry began with Daltons atomic theory,which states that all matter is composed of tiny, indivisibleparticles called atoms; that all atoms of the sameelement are identical; that compounds contain atoms ofdifferent elements combined in whole-number ratios;and that atoms are neither created nor destroyed inchemical reactions (the law of conservation of mass).2. Atoms of constituent elements in a particular compoundare always combined in the same proportions by mass(law of defi nite proportions). When two elements cancombine to form more than one type of compound, themasses of one element that combine with a fi xed massof the other element are in a ratio of small whole numbers(law of multiple proportions).3. An atom consists of a very dense central nucleus containingprotons and neutrons, with electrons movingabout the nucleus at a relatively large distance from it.4. Protons are positively charged, neutrons have no charge,and electrons are negatively charged. Protons and neutronshave roughly the same mass, which is about 1840times greater than the mass of an electron.5. The atomic number of an element is the number of protonsin the nucleus of an atom of the element; it determinesthe identity of an element. The mass number isSummary of Facts and Conceptsthe sum of the number of protons and the number ofneutrons in the nucleus.6. Isotopes are atoms of the same element with the samenumber of protons but different numbers of neutrons.7. Chemical formulas combine the symbols for the constituentelements with whole-number subscripts to showthe type and number of atoms contained in the smallestunit of a compound.8. The molecular formula conveys the specifi c numberand type of atoms combined in each molecule of a compound.The empirical formula shows the simplest ratiosof the atoms combined in a molecule.9. Chemical compounds are either molecular compounds(in which the smallest units are discrete, individual molecules)or ionic compounds, which are made of cationsand anions.10. The names of many inorganic compounds can be deducedfrom a set of simple rules. The formulas can bewritten from the names of the compounds.11. Organic compounds contain carbon and elements likehydrogen, oxygen, and nitrogen. Hydrocarbon is thesimplest type of organic compound.Media PlayerChapter SummaryLaw of conservation ofmass, p. 43Law of defi niteproportions, p. 43Law of multipleproportions, p. 43Mass number ( A ), p. 49Metal, p. 51Metalloid, p. 51Molecular formula, p. 55Molecule, p. 53Monatomic ion, p. 54Neutron, p. 48Noble gases, p. 53Nonmetal, p. 51Nucleus, p. 47Organic compound, p. 59Oxoacid, p. 65Oxoanion, p. 66Periods, p. 51Periodic table, p. 51Polyatomic ion, p. 54Polyatomic molecule, p. 53Proton, p. 47Radiation, p. 44Radioactivity, p. 46Structural formula, p. 56Ternary compound, p. 60Questions and ProblemsStructure of the AtomReview Questions2.1 Defi ne the following terms: (a) a particle, (b) b particle,(c) g ray, (d) X ray.2.2 Name the types of radiation known to be emitted byradioactive elements.2.3 Compare the properties of the following: a particles,cathode rays, protons, neutrons, electrons.2.4 What is meant by the term fundamental particle?2.5 Describe the contributions of the following scientiststo our knowledge of atomic structure: J. J. Thomson,R. A. Millikan, Ernest Rutherford, James Chadwick.2.6 Describe the experimental basis for believing that thenucleus occupies a very small fraction of the volumeof the atom.Problems2.7 The diameter of a helium atom is about 1 3 102 pm.Suppose that we could line up helium atoms side byside in contact with one another. Approximately howmany atoms would it take to make the distance fromend to end 1 cm?2.8 Roughly speaking, the radius of an atom is about10,000 times greater than that of its nucleus. If anatom were magnifi ed so that the radius of its nucleusbecame 2.0 cm, about the size of a marble, what wouldbe the radius of the atom in miles? (1 mi 5 1609 m.)Atomic Number, Mass Number, and IsotopesReview Questions2.9 Use the helium-4 isotope to defi ne atomic number andmass number. Why does a knowledge of atomic numberenable us to deduce the number of electrons presentin an atom?2.10 Why do all atoms of an element have the same atomicnum ber, although they may have different massnumbers?2.11 What do we call atoms of the same elements withdifferent mass numbers?2.12 Explain the meaning of each term in the symbolZAX.Problems2.13 What is the mass number of an iron atom that has28 neutrons?2.14 Calculate the number of neutrons of 239Pu.2.15 For each of the following species, determine the numberof protons and the number of neutrons in the nucleus:2 3He, 24 He, 1224Mg, 2512Mg, 4822Ti, 7935Br, 19578Pt2.16 Indicate the number of protons, neutrons, and electronsin each of the following species:157N, 3316S, 6329Cu, 8438Sr, 13056Ba, 18674W, 20280Hg2.17 Write the appropriate symbol for each of the followingisotopes: (a) Z 5 11, A 5 23; (b) Z 5 28, A 5 64.Questions and Problems 71Electronic Homework ProblemsThe following problems are available at www.aris.mhhe.com if assigned by your instructor as electronic homework.Quantum Tutor problems are also available at the same site.ARIS Problems: 2.13, 2.15, 2.22, 2.32, 2.35, 2.36,2.43, 2.44, 2.46, 2.48, 2.49, 2.50, 2.58, 2.59, 2.60, 2.63,2.65, 2.77, 2.90, 2.91, 2.96, 2.97, 2.100, 2.101, 2.102.Quantum Tutor Problems: 2.43, 2.44, 2.45, 2.46,2.57, 2.58, 2.59, 2.60.72 Atoms, Molecules, and Ions2.18 Write the appropriate symbol for each of the followingisotopes: (a) Z 5 74, A 5 186; (b) Z 5 80; A 5 201.The Periodic TableReview Questions2.19 What is the periodic table, and what is its signifi cancein the study of chemistry?2.20 State two differences between a metal and a nonmetal.2.21 Write the names and symbols for four elements in eachof the following categories: (a) nonmetal, (b) metal,(c) metalloid.2.22 Defi ne, with two examples, the following terms:(a) alkali metals, (b) alkaline earth metals, (c) halogens,(d) noble gases.Problems2.23 Elements whose names end with ium are usually metals;sodium is one example. Identify a nonmetal whosename also ends with ium.2.24 Describe the changes in properties (from metals tononmetals or from nonmetals to metals) as we move(a) down a periodic group and (b) across the periodictable from left to right.2.25 Consult a handbook of chemical and physical data(ask your instructor where you can locate a copy ofthe handbook) to fi nd (a) two metals less dense thanwater, (b) two metals more dense than mercury,(c) the densest known solid metallic element, (d) thedensest known solid nonmetallic element.2.26 Group the following elements in pairs that you wouldexpect to show similar chemical properties: K, F, P,Na, Cl, and N.Molecules and IonsReview Questions2.27 What is the difference between an atom and amolecule?2.28 What are allotropes? Give an example. How areallotropes different from isotopes?2.29 Describe the two commonly used molecular models.2.30 Give an example of each of the following: (a) a monatomiccation, (b) a monatomic anion, (c) a polyatomiccation, (d) a polyatomic anion.Problems2.31 Which of the following diagrams represent diatomicmolecules, polyatomic molecules, molecules that arenot compounds, molecules that are compounds, or anelemental form of the substance?2.32 Which of the following diagrams represent diatomicmolecules, polyatomic molecules, molecules that arenot compounds, molecules that are compounds, or anelemental form of the substance?(a) (b) (c)(a) (b) (c)2.33 Identify the following as elements or compounds:NH3, N2, S8, NO, CO, CO2, H2, SO2.2.34 Give two examples of each of the following: (a) adiatomic molecule containing atoms of the sameelement, (b) a diatomic molecule containing atoms ofdifferent elements, (c) a polyatomic molecule containingatoms of the same element, (d) a polyatomic moleculecontaining atoms of different elements.2.35 Give the number of protons and electrons in each ofthe following common ions: Na1 , Ca21, Al31, Fe21,I2, F2, S22, O22, and N32.2.36 Give the number of protons and electrons in each ofthe following common ions: K1, Mg21, Fe31, Br2,Mn21, C42, Cu21.Chemical FormulasReview Questions2.37 What does a chemical formula represent? What is theratio of the atoms in the following molecular formulas?(a) NO, (b) NCl3, (c) N2O4, (d) P4O62.38 Defi ne molecular formula and empirical formula.What are the similarities and differences between2.49 Which of the following compounds are likely to beionic? Which are likely to be molecular? SiCl4, LiF,BaCl2, B2H6, KCl, C2H42.50 Which of the following compounds are likely to beionic? Which are likely to be molecular? CH4, NaBr,BaF2, CCl4, ICl, CsCl, NF3Naming Inorganic CompoundsReview Questions2.51 What is the difference between inorganic compoundsand organic compounds?2.52 What are the four major categories of inorganiccompounds?2.53 Give an example each for a binary compound and aternary compound.2.54 What is the Stock system? What are its advantagesover the older system of naming cations?2.55 Explain why the formula HCl can represent two differentchemical systems.2.56 Defi ne the following terms: acids, bases, oxoacids,oxoanions, and hydrates.Problems2.57 Name these compounds: (a) Na2CrO4, (b) K2HPO4,(c) HBr (gas), (d) HBr (in water), (e) Li2CO3,(f) K2Cr2O7, (g) NH4NO2, (h) PF3, (i) PF5, (j) P4O6,(k) CdI2, (l) SrSO4, (m) Al(OH)3, (n) Na2CO3 ? 10H2O.2.58 Name these compounds: (a) KClO, (b) Ag2CO3,(c) FeCl2, (d) KMnO4, (e) CsClO3, (f) HIO, (g) FeO,(h) Fe2O3, (i) TiCl4, (j) NaH, (k) Li3N, (l) Na2O,(m) Na2O2, (n) FeCl3 ? 6H2O.2.59 Write the formulas for the following compounds:(a) rubidium nitrite, (b) potassium sulfi de, (c) sodiumhydrogen sulfi de, (d) magnesium phosphate, (e) calciumhydrogen phosphate, (f) potassium dihydrogen phosphate,(g) iodine heptafl uoride, (h) ammonium sulfate,(i) silver perchlorate, (j) boron trichloride.2.60 Write the formulas for the following compounds:(a) copper(I) cyanide, (b) strontium chlorite, (c) perbromicacid, (d) hydroiodic acid, (e) disodium ammoniumphosphate, (f) lead(II) carbonate, (g) tin(II) fl uoride,(h) tetraphosphorus decasulfi de, (i) mercury(II) oxide,(j) mercury(I) iodide, (k) selenium hexafl uoride.Questions and Problems 73HCNOthe empirical formula and molecular formula of acompound?2.39 Give an example of a case in which two moleculeshave different molecular formulas but the same empiricalformula.2.40 What does P4 signify? How does it differ from 4P?2.41 What is an ionic compound? How is electrical neutralitymaintained in an ionic compound?2.42 Explain why the chemical formulas of ionic compoundsare usually the same as their empirical formulas.Problems2.43 Write the formulas for the following ionic compounds:(a) sodium oxide, (b) iron sulfi de (containingthe Fe21 ion), (c) cobalt sulfate (containing theCo31 and SO422 ions), and (d) barium fl uoride. ( Hint:See Figure 2.11 .)2.44 Write the formulas for the following ionic compounds:(a) copper bromide (containing the Cu1 ion), (b) manganeseoxide (containing the Mn31 ion), (c) mercuryiodide (containing the Hg221 ion), and (d) magnesiumphosphate (containing the PO432 ion). ( Hint: See Figure2.11 .)2.45 What are the empirical formulas of the followingcompounds? (a) C2N2, (b) C6H6, (c) C9H20, (d) P4O10,(e) B2H62.46 What are the empirical formulas of the followingcompounds? (a) Al2Br6, (b) Na2S2O4, (c) N2O5,(d) K2Cr2O72.47 Write the molecular formula of glycine, an aminoacid present in proteins. The color codes are: black(carbon), blue (nitrogen), red (oxygen), and gray(hydrogen).HCO2.48 Write the molecular formula of ethanol. The colorcodes are: black (carbon), red (oxygen), and gray(hydrogen).74 Atoms, Molecules, and IonsAdditional Problems2.61 A sample of a uranium compound is found to be losingmass gradually. Explain what is happening to thesample.2.62 In which one of the following pairs do the two speciesresemble each other most closely in chemical properties?Explain. (a) 11H and 11H1, (b) 147N and 147N32,(c) 126C and 136C.2.63 One isotope of a metallic element has mass number 65and 35 neutrons in the nucleus. The cation derivedfrom the isotope has 28 electrons. Write the symbolfor this cation.2.64 One isotope of a nonmetallic element has mass number127 and 74 neutrons in the nucleus. The anionderived from the isotope has 54 electrons. Write thesymbol for this anion.2.65 Determine the molecular and empirical formulas ofthe compounds shown here. (Black spheres are carbonand gray spheres are hydrogen.)2.71 Explain why anions are always larger than the atomsfrom which they are derived, whereas cations arealways smaller than the atoms from which they arederived. ( Hint: Consider the electrostatic attractionbetween protons and electrons.)2.72 (a) Describe Rutherfords experiment and how it led tothe structure of the atom. How was he able to estimatethe number of protons in a nucleus from the scatteringof the a particles? (b) Consider the 23Na atom. Giventhat the radius and mass of the nucleus are3.04 3 10215 m and 3.82 3 10223 g, respectively,calculate the density of the nucleus in g/cm3. The radiusof a 23Na atom is 186 pm. Calculate the density of thespace occupied by the electrons in the sodium atom.Do your results support Rutherfords model of an atom?[The volume of a sphere of radius r is (4/3)pr3.]2.73 What is wrong with the name (in parentheses) for eachof the following compounds: (a) BaCl2 (barium dichloride),(b) Fe2O3 [iron(II) oxide], (c) CsNO2 (cesiumnitrate), (d) Mg(HCO3)2 [magnesium(II) bicarbonate]?2.74 What is wrong with the chemical formula for each ofthe following compounds: (a) (NH3)2CO3 (ammoniumcarbonate), (b) CaOH (calcium hydroxide), (c) CdSO3(cadmium sulfi de), (d) ZnCrO4 (zinc dichromate)?2.75 Fill in the blanks in the following table:Symbol 5426Fe21Protons 5 79 86Neutrons 6 16 117 136Electrons 5 18 79Net charge 23 02.76 (a) Which elements are most likely to form ionic compounds?(b) Which metallic elements are most likelyto form cations with different charges?2.77 Write the formula of the common ion derived fromeach of the following: (a) Li, (b) S, (c) I, (d) N, (e) Al,(f) Cs, (g) Mg2.78 Which of the following symbols provides more informationabout the atom: 23Na or 11Na? Explain.2.79 Write the chemical formulas and names of binary acidsand oxoacids that contain Group 7A elements. Do thesame for elements in Groups 3A, 4A, 5A, and 6A.2.80 Of the 117 elements known, only two are liquids atroom temperature (25C). What are they? ( Hint: Oneelement is a familiar metal and the other element is inGroup 7A.)2.81 For the noble gases (the Group 8A elements), 42He,2010Ne, 4018Ar, 8436Kr, and 13254Xe, (a) determine the numberof protons and neutrons in the nucleus of each atom,and (b) determine the ratio of neutrons to protons inthe nucleus of each atom. Describe any general trend(a) (b) (c) (d)2.66 What is wrong with or ambiguous about the phrasefour molecules of NaCl?2.67 The following phosphorus sulfi des are known: P4S3,P4S7, and P4S10. Do these compounds obey the law ofmultiple proportions?2.68 Which of the following are elements, which are moleculesbut not compounds, which are compounds butnot molecules, and which are both compounds andmolecules? (a) SO2, (b) S8, (c) Cs, (d) N2O5, (e) O,(f) O2, (g) O3, (h) CH4, (i) KBr, (j) S, (k) P4, (l) LiF2.69 The following table gives numbers of electrons, protons,and neutrons in atoms or ions of a numbe