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Answers to Even-Numbered Cumulative Review Exercises

28. The ionization energy of an atom represents the energy required to remove an electron from the atom in the gas phase.

Moving from top to bottom in a vertical group on the periodic table, the ionization energies decrease. The ionization

energies increase when going from left to right within a horizontal row within the periodic table. The relative sizes of

atoms also vary systematically with the location of an element on the periodic table. Within a given vertical group, the

atoms become progressively larger when going from the top

of the group to the bottom. Moving from left to right within

a horizontal row on the periodic table, the atoms become progressively smaller.

30. To form an ionic compound, a metallic element reacts with a

nonmetallic element, with the metallic element losing electrons to form a positive ion and the nonmetallic element

gaining electrons to form a negative ion. The aggregate form

of such a compound consists of a crystal lattice of alternating

positively and negatively charged ions: a given positive ion is

attracted by surrounding negatively charged ions, and a given

negative ion is attracted by surrounding positively charged

ions. Similar electrostatic attractions exist in three dimensions throughout the crystal of the ionic solid, leading to a

very stable system (with very high melting and boiling

points, for example). As evidence for the existence of ionic

bonding, ionic solids do not conduct electricity (the ions are

rigidly held), but melts or solutions of such substances do

conduct electric current. For example, when sodium metal

and chlorine gas react, a typical ionic substance (sodium chloride) results: 2Na(s) ϩ Cl2(g) S 2NaϩClϪ(s).

32. Electronegativity represents the relative ability of an atom in

a molecule to attract shared electrons to itself. The larger the

difference in electronegativity between two atoms joined in a

bond, the more polar is the bond. Examples depend on student choice of elements.

34. It has been observed over many, many experiments that when

an active metal like sodium or magnesium reacts with a nonmetal, the sodium atoms always form Naϩ ions and the magnesium atoms always form Mg2ϩ ions. It has also been observed that when nonmetallic elements like nitrogen, oxygen,

or fluorine form simple ions, the ions are always N3Ϫ, O2Ϫ, and

FϪ, respectively. Observing that these elements always form

the same ions and those ions all contain eight electrons in the

outermost shell, scientists speculated that a species that has an

octet of electrons (like the noble gas neon) must be very fundamentally stable. The repeated observation that so many elements, when reacting, tend to attain an electron configuration

that is isoelectronic with a noble gas led chemists to speculate

that all elements try to attain such a configuration for their

outermost shells. Covalently and polar covalently bonded

molecules also strive to attain pseudo–noble gas electron configurations. For a covalently bonded molecule like F2, each

F atom provides one electron of the pair of electrons that constitutes the covalent bond. Each F atom feels also the influence

of the other F atom’s electron in the shared pair, and each

F atom effectively fills its outermost shell.

36. Bonding between atoms to form a molecule involves only the

outermost electrons of the atoms, so only these valence electrons are shown in the Lewis structures of molecules. The

most important requisite for the formation of a stable compound is that each atom of a molecule attain a noble gas electron configuration. In Lewis structures, arrange the bonding

and nonbonding valence electrons to try to complete the

octet (or duet) for as many atoms as possible.

38. You could choose practically any molecules for your discussion. Let’s illustrate the method for ammonia, NH3. First,

count the total number of valence electrons available in the

molecule (without regard to their source). For NH3, since ni-



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trogen is in Group 5, one nitrogen atom would contribute five

valence electrons. Since hydrogen atoms have only one electron each, the three hydrogen atoms provide an additional

three valence electrons, for a total of eight valence electrons

overall. Next, write down the symbols for the atoms in the

molecule, and use one pair of electrons (represented by a line)

to form a bond between each pair of bound atoms.



H



N



H



H

These three bonds use six of the eight valence electrons. Because each hydrogen already has its duet and the nitrogen

atom has only six electrons around it so far, the final two valence electrons must represent a lone pair on the nitrogen.



H



N



H



H

40. Boron and beryllium compounds sometimes do not fit the

octet rule. For example, in BF3, the boron atom has only six

valence electrons in its outermost shell, whereas in BeF2, the

beryllium atom has only four electrons in its outermost shell.

Other exceptions to the octet rule include any molecule with

an odd number of valence electrons (such as NO or NO2).

42. Number of Valence

Pairs



Bond

Angle



Examples



180Њ



BeF2, BeH2



3



120Њ



BCl3



4



109.5Њ



2



CH4, CCl4, GeF4



44. (a) [Kr]5s2; (b) [Ne]3s23p1; (c) [Ne]3s23p5; (d) [Ar]4s1;

(e) [Ne]3s23p4; (f) [Ar]4s23d 104p3

46.



H



O



H



4 electron pairs tetrahedrally oriented on

O; nonlinear (bent, V-shaped) geometry;

HOOOH bond angle slightly less than

109.5° because of lone pairs



H



P



H



4 electron pairs tetrahedrally oriented on

P; trigonal pyramidal geometry; HOPOH

bond angles slightly less than 109.5° because of lone pair



H



Br

Br



C



4 electron pairs tetrahedrally oriented on

C; overall tetrahedral geometry;

BrOCOBr bond angles 109.5°



Br



Br

Ϫ



O

O



Cl



O



4 electron pairs tetrahedrally oriented

on Cl; overall tetrahedral geometry;

OOClOO bond angles 109.5°



O

F

B



F



F

F



Be



F



3 electron pairs trigonally oriented on B

(exception to octet rule); overall trigonal

geometry; FOBOF bond angles 120°

2 electron pairs linearly oriented on Be

(exception to octet rule); overall linear

geometry; FOBeOF bond angle 180°



Chapters 13–15

2. The pressure of the atmosphere represents the mass of the

gases in the atmosphere pressing down on the surface of the

earth. The device most commonly used to measure the pres-



A54 Answers to Even-Numbered Cumulative Review Exercises



4.



6.



8.



10.



12.



sure of the atmosphere is the mercury barometer, shown in

Figure 13.2. A simple experiment to demonstrate the pressure

of the atmosphere is shown in Figure 13.1.

Boyle’s law says that the volume of a gas sample will decrease

if you squeeze it harder (at constant temperature, for a fixed

amount of gas). Two mathematical statements of Boyle’s

law are

P ϫ V ϭ constant

P1 ϫ V1 ϭ P2 ϫ V2

These two mathematical formulas say the same thing: if the

pressure on a sample of gas is increased, the volume of the

sample will decrease. A graph of Boyle’s law data is given as

Figure 13.5: this type of graph (xy ϭ k) is known to mathematicians as a hyperbola.

Charles’s law says that if you heat a sample of gas, the volume

of the sample will increase (assuming the pressure and

amount of gas remain the same). When the temperature is

given in kelvins, Charles’s law expresses a direct proportionality (if you increase T, then V increases), whereas Boyle’s law

expresses an inverse proportionality (if you increase P, then V

decreases). Two mathematical statements of Charles’s law are

V ϭ bT and (V1/T1) ϭ (V2/T2). With this second formulation,

we can determine volume–temperature information for a

given gas sample under two sets of conditions. Charles’s law

holds true only if the amount of gas remains the same (the

volume of a gas sample would increase if more gas were present) and also if the pressure remains the same (a change in

pressure also changes the volume of a gas sample). A graph

of volume versus temperature (at constant pressure) for an

ideal gas is a straight line with an intercept at –273 ЊC (see

Figure 13.7).

Avogadro’s law says that the volume of a sample of gas is directly proportional to the number of moles (or molecules)

of gas present (at constant temperature and pressure).

Avogadro’s law holds true only for gas samples compared

under the same conditions of temperature and pressure.

Avogadro’s law expresses a direct proportionality: the more

gas in a sample, the larger the sample’s volume.

The “partial” pressure of an individual gas in a mixture of

gases represents the pressure the gas would exert in the same

container at the same temperature if it were the only gas present. The total pressure in a mixture of gases is the sum of the

individual partial pressures of the gases present in the mixture. The fact that the partial pressures of the gases in a mixture are additive suggests that the total pressure in a container

is a function of the number of molecules present, and not of

the identity of the molecules or of any other property (such as

the molecules’ inherent atomic size).

The main postulates of the kinetic molecular theory for gases

are: (a) gases consist of tiny particles (atoms or molecules),

and the size of these particles themselves is negligible compared with the bulk volume of a gas sample; (b) the particles

in a gas are in constant random motion, colliding with each

other and with the walls of the container; (c) the particles in

a gas sample do not assert any attractive or repulsive forces on

one another; and (d) the average kinetic energy of the gas

particles is directly related to the absolute temperature of

the gas sample. The pressure exerted by a gas results from the

molecules colliding with (and pushing on) the walls of the

container; the pressure increases with temperature because, at

a higher temperature, the molecules move faster and hit the

walls of the container with greater force. A gas fills the volume

available to it because the molecules in a gas are in constant

random motion: the randomness of the molecules’ motion

means that they eventually will move out into the available

volume until the distribution of molecules is uniform; at constant pressure, the volume of a gas sample increases as the



14.



16.



18.



20.



22.



temperature is increased because with each collision having

greater force, the container must expand so that the molecules are farther apart if the pressure is to remain constant.

The molecules are much closer together in solids and liquids

than in gaseous substances and interact with each other to a

much greater extent. Solids and liquids have much greater

densities than do gases, and are much less compressible, because so little room exists between the molecules in the solid

and liquid states (the volume of a solid or liquid is not affected very much by temperature or pressure). We know that

the solid and liquid states of a substance are similar to each

other in structure, since it typically takes only a few kilojoules

of energy to melt 1 mole of a solid, whereas it may take

10 times more energy to convert a liquid to the vapor state.

The normal boiling point of water—that is, water’s boiling

point at a pressure of exactly 760 mm Hg—is 100 ЊC. Water

remains at 100 ЊC while boiling, because the additional energy added to the sample is used to overcome attractive forces

among the water molecules as they go from the condensed,

liquid state to the gaseous state. The normal (760 mm Hg)

freezing point of water is exactly 0 ЊC. A cooling curve for water is given in Figure 14.2.

Dipole–dipole forces arise when molecules with permanent

dipole moments try to orient themselves so that the positive

end of one polar molecule can attract the negative end of another polar molecule. Dipole–dipole forces are not nearly as

strong as ionic or covalent bonding forces (only about 1% as

strong as covalent bonding forces) since electrostatic attraction is related to the magnitude of the charges of the attracting species and drops off rapidly with distance. Hydrogen

bonding is an especially strong dipole–dipole attractive force

that can exist when hydrogen atoms are directly bonded to

the most electronegative atoms (N, O, and F). Because the hydrogen atom is so small, dipoles involving NOH, OOH, and

FOH bonds can approach each other much more closely

than can other dipoles; because the magnitude of dipole–

dipole forces is related to distance, unusually strong attractive forces can exist. The much higher boiling point of water

than that of the other covalent hydrogen compounds of the

Group 6 elements is evidence for the special strength of hydrogen bonding.

The vaporization of a liquid requires an input of energy to

overcome the intermolecular forces that exist between the

molecules in the liquid state. The large heat of vaporization of

water is essential to life since much of the excess energy striking the earth from the sun is dissipated in vaporizing water.

Condensation refers to the process by which molecules in the

vapor state form a liquid. In a closed container containing a

liquid with some empty space above the liquid, an equilibrium occurs between vaporization and condensation. When

the liquid is first placed in the container, the liquid phase begins to evaporate into the empty space. As the number of molecules in the vapor phase increases, however, some of these

molecules begin to reenter the liquid phase. Eventually, each

time a molecule of liquid somewhere in the container enters

the vapor phase, another molecule of vapor reenters the liquid

phase. No further net change occurs in the amount of liquid

phase. The pressure of the vapor in such an equilibrium situation is characteristic for the liquid at each temperature. A simple experiment to determine the vapor pressure of a liquid is

shown in Figure 14.10. Typically, liquids with strong intermolecular forces have smaller vapor pressures (they have more

difficulty in evaporating) than do liquids with very weak intermolecular forces.

The electron sea model explains many properties of metallic elements. This model pictures a regular array of metal atoms set

in a “sea” of mobile valence electrons. The electrons can



Answers to Even-Numbered Cumulative Review Exercises



24.



26.



28.

30.

32.

34.

36.

38.



move easily throughout the metal to conduct heat or electricity, and the lattice of atoms and cations can be deformed

with little effort, allowing the metal to be hammered into a

sheet or stretched into wire. An alloy is a material that contains a mixture of elements that overall has metallic properties. Substitutional alloys consist of a host metal in which

some of the atoms in the metal’s crystalline structure are replaced by atoms of other metallic elements. For example, sterling silver is an alloy in which some silver atoms have been replaced by copper atoms. An interstitial alloy is formed when

other, smaller atoms enter the interstices (holes) between

atoms in the host metal’s crystal structure. Steel is an interstitial alloy in which carbon atoms enter the interstices of a crystal of iron atoms.

A saturated solution contains as much solute as can dissolve at

a particular temperature. Saying that a solution is saturated

does not necessarily mean that the solute is present at a high

concentration—for example, magnesium hydroxide dissolves

only to a very small extent before the solution is saturated. A

saturated solution is in equilibrium with undissolved solute: as

molecules of solute dissolve from the solid in one place in the

solution, dissolved molecules rejoin the solid phase in another

part of the solution. Once the rates of dissolving and solid formation become equal, no further net change occurs in the concentration of the solution and the solution is saturated.

Adding more solvent to a solution to dilute the solution does

not change the number of moles of solute present, but changes

only the volume in which the solute is dispersed. If molarity is

used to describe the solution’s concentration, then the number

of liters is changed when solvent is added and the number of

moles per liter (the molarity) changes, but the actual number

of moles of solute does not change. For example, 125 mL of

0.551 M NaCl contains 0.0689 mole of NaCl. The solution will

still contain 0.0689 mole of NaCl after 250 mL of water is added

to it. The volume and the concentration will change, but the

number of moles of solute in the solution will not change. The

0.0689 mole of NaCl, divided by the total volume of the diluted

solution in liters, gives the new molarity (0.184 M).

(a) 105 mL; (b) 1.05 ϫ 103 mm Hg

(a) 6.96 L; (b) Phydrogen ϭ 5.05 atm; Phelium ϭ 1.15 atm;

(c) 2.63 atm

0.550 g CO2; 0.280 L CO2 at STP

(a) 9.65% NaCl; (b) 2.75 g CaCl2; (c) 11.4 g NaCl

(a) 0.505 M; (b) 0.0840 M; (c) 0.130 M

(a) 226 g; (b) 18.4 M; (c) 0.764 M; (d) 1.53 N;

(e) 15.8 mL



Chapters 16–17

2. A conjugate acid–base pair consists of two species related to

one another by the donating or accepting of a single proton,

Hϩ. An acid has one more Hϩ than its conjugate base; a base

has one less Hϩ than its conjugate acid.

Brønsted–Lowry acids:

HCl(aq) ϩ H2O(l) S ClϪ(aq) ϩ H3Oϩ(aq)

H2SO4(aq) ϩ H2O(l) S HSO4Ϫ(aq) ϩ H3Oϩ(aq)

H3PO4(aq) ϩ H2O(l) S H2PO4Ϫ(aq) ϩ H3Oϩ(aq)

NH4ϩ(aq) ϩ H2O(l) S NH3(aq) ϩ H3Oϩ(aq)

Brønsted–Lowry bases:

NH3(aq) ϩ H2O(l) S NH4ϩ(aq) ϩ OHϪ(aq)

HCO3Ϫ(aq) ϩ H2O(l) S H2CO3(aq) ϩ OHϪ(aq)

NH2Ϫ(aq) ϩ H2O(l) S NH3(aq) ϩ OHϪ(aq)

H2PO4Ϫ(aq) ϩ H2O(l) S H3PO4(aq) ϩ OHϪ(aq)

4. The strength of an acid is a direct result of the position of the

acid’s dissociation (ionization) equilibrium. Acids whose dissociation equilibrium positions lie far to the right are called



A55



strong acids. Acids whose equilibrium positions lie only

slightly to the right are called weak acids. For example, HCl,

HNO3, and HClO4 are strong acids, which means that they

are completely dissociated in aqueous solution (the position

of equilibrium is very far to the right):

HCl(aq) ϩ H2O(l) S ClϪ(aq) ϩ H3Oϩ(aq)

HNO3(aq) ϩ H2O(l) S NO3Ϫ(aq) ϩ H3Oϩ(aq)

HClO4(aq) ϩ H2O(l) S ClO4Ϫ(aq) ϩ H3Oϩ(aq)

Since these are very strong acids, their anions (ClϪ, NO3Ϫ,

ClO4Ϫ) must be very weak bases, and solutions of their sodium

salts will not be basic.

6. The pH of a solution is defined as pH ϭ Ϫlog[Hϩ(aq)] for a

solution. In pure water, the amount of Hϩ(aq) ion present is

equal to the amount of OHϪ(aq) ion—that is, pure water is

neutral. Since [Hϩ] ϭ 1.0 ϫ 10Ϫ7 M in pure water, the pH of

pure water is Ϫlog[1.0 ϫ 10Ϫ7 M] ϭ 7.00. Solutions in which

[Hϩ] Ͼ 1.0 ϫ 10Ϫ7 M (pH Ͻ 7.00) are acidic; solutions in which

[Hϩ] Ͻ 1.0 ϫ 10Ϫ7 M (pH Ͼ 7.00) are basic. The pH scale is

logarithmic: a pH change of one unit corresponds to a change

in the hydrogen ion concentration by a factor of ten. An analogous logarithmic expression is defined for the hydroxide ion

concentration in a solution: pOH ϭ Ϫlog[OHϪ(aq)]. The concentrations of hydrogen ion and hydroxide ion in water (and

in aqueous solutions) are not independent of one another, but

rather are related by the dissociation equilibrium constant for

water, Kw ϭ [Hϩ][OHϪ] ϭ 1.0 ϫ 10Ϫ14 at 25 ЊC. From this expression it follows that pH ϩ pOH ϭ 14.00 for water (or an aqueous solution) at 25 ЊC.

8. Chemists envision that a reaction can take place between

molecules only if the molecules physically collide with each

other. Furthermore, when molecules collide, the molecules

must collide with enough force for the reaction to be

successful (there must be enough energy to break bonds in

the reactants), and the colliding molecules must be positioned with the correct relative orientation for the products

(or intermediates) to form. Reactions tend to be faster if

higher concentrations are used for the reaction because if

there are more molecules present per unit volume there will

be more collisions between molecules in a given time period.

Reactions are faster at higher temperatures because at higher

temperatures the reactant molecules have a higher average kinetic energy, and the number of molecules that will collide

with sufficient force to break bonds increases.

10. Chemists define equilibrium as the exact balancing of two exactly opposing processes. When a chemical reaction is begun

by combining pure reactants, the only process possible initially is

reactants S products

However, for many reactions, as the concentration of product

molecules increases, it becomes more and more likely that

product molecules will collide and react with each other,

products S reactants

giving back molecules of the original reactants. At some point

in the process the rates of the forward and reverse reactions

become equal, and the system attains chemical equilibrium.

To an outside observer the system appears to have stopped reacting. On a microscopic basis, though, both the forward and

reverse processes are still going on. Every time additional molecules of the product form, however, somewhere else in the

system molecules of product react to give back molecules of

reactant.

Once the point is reached that product molecules are

reacting at the same speed at which they are forming, there is

no further net change in concentration. At the start of the re-



A56 Answers to Even-Numbered Cumulative Review Exercises

action, the rate of the forward reaction is at its maximum,

while the rate of the reverse reaction is zero. As the reaction

proceeds, the rate of the forward reaction gradually decreases

as the concentration of reactants decreases, whereas the rate

of the reverse reaction increases as the concentration of products increases. Once the two rates have become equal, the reaction has reached a state of equilibrium.

12. The equilibrium constant for a reaction is a ratio of the concentration of products present at the point of equilibrium to

the concentration of reactants still present. A ratio means that

we have one number divided by another number (for example, the density of a substance is the ratio of a substance’s

mass to its volume). Since the equilibrium constant is a ratio,

there are an infinite number of sets of data that can give the

same ratio: for example, the ratios 8/4, 6/3, 100/50 all have

the same value, 2. The actual concentrations of products and

reactants will differ from one experiment to another involving a particular chemical reaction, but the ratio of the amount

of product to reactant at equilibrium should be the same for

each experiment.

14. Your paraphrase of Le Châtelier’s principle should go something like this: “When you make any change to a system in

equilibrium, this throws the system temporarily out of

equilibrium, and the system responds by reacting in whatever

direction it will be able to reach a new position of

equilibrium.” There are various changes that can be made to

a system in equilibrium. Here are examples of some of them.

a. The concentration of one of the reactants is increased.

z 2SO3(g)

2SO2(g) ϩ O2(g) y

If additional SO2 or O2 is added to the system at equilibrium,

then more SO3 will result than if no change was made.

b. The concentration of one of the products is decreased by

selectively removing it from the system.

z H2O ϩ CH3COOCH3

CH3COOH ϩ CH3OH y



16.



18.



20.



22.

24.



If H2O were to be removed from the system by, for

example, use of a drying agent, then more CH3COOCH3

would result than if no change was made.

c. The reaction system is compressed to a smaller volume.

z 2NH3(g)

3H2(g) ϩ N2(g) y

If this system is compressed to smaller volume, then more

NH3 would be produced than if no change was made.

d. The temperature is increased for an endothermic reaction.

z Na2CO3 ϩ H2O ϩ CO2

2NaHCO3 ϩ heat y

If heat is added to this system, then more product would

be produced than if no change was made.

e. The temperature is decreased for an exothermic process.

z PCl5 ϩ heat

PCl3 ϩ Cl2 y

If heat is removed from this system (by cooling), then more

PCl5 would be produced than if no change was made.

Specific answer depends on student choices. In general, for a

weak acid HA and a weak base B,

z AϪ ϩ H3Oϩ

HA ϩ H2O y

z HBϩ ϩ OHϪ

B ϩ H2O y

z NH4ϩ(aq)(acid) ϩ

(a) NH3(aq)(base) ϩ H2O(l)(acid) y

OHϪ (aq)(base);

z HSO4Ϫ(aq)(base) ϩ

(b) H2SO4 (aq)(acid) ϩ H2O(l)(base) y

H3Oϩ (aq)(acid);

z OHϪ(aq)(acid) ϩ

(c) O2Ϫ(s)(base) ϩ H2O(l)(acid) y

OHϪ (aq)(base);

z NH3(aq)(acid) ϩ

(d) NH2Ϫ(aq)(base) ϩ H2O(l)(acid) y

OHϪ (aq)(base);

z

(e) H2PO4Ϫ(aq)(acid) ϩ OHϪ(aq)(base) y

HPO42Ϫ(aq)(base) ϩ H2O(l)(acid)

(a) pH ϭ 2.851; pOH ϭ 11.149; (b) pOH ϭ 2.672; pH ϭ

11.328; (c) pH ϭ 2.288; pOH ϭ 11.712; (d) pOH ϭ 3.947;

pH ϭ 10.053

7.8 ϫ 105

0.220 g/L



I N D E X A N D G L O S S A RY



Page numbers followed by n refer to margin notes. Page numbers followed by f refer to figures. Page numbers followed by t refer to tables.

Absolute scale, 35

Absolute zero Ϫ273 °C, 412

Accelerator, particle, 620

Acetic acid

buffered solution and, 534

electric current and, 521f

naming of, 133

solution of, 489, 490f

strength of, 519–520

Acid A substance that produces hydrogen

ions in aqueous solution; proton donor,

132, 514–542

acetic. See Acetic acid

bases and, 514–518

buffered solutions and, 534–535

calculating pH of, 532–533

conjugate, 516

conjugate acid-base pairs, 516–518

diprotic, 521

equivalent of, 497

formation of, 179–182, 180f

hydrochloric. See Hydrochloric acid

hydrofluoric, 157, 260

mineral, 179

naming of, 132–133, 133f, 133t

organic, 521

pH scale and, 525–533

strength of, 518–523, 519n, 520f, 520t,

521n

strong, 180, 519

sulfuric. See Sulfuric acid

water as, 523–525

weak, 519–520, 520f, 520t

conjugate base and, 534

Acid–base indicator, 532

Acid–base pair, conjugate, 516–518

Acid–base reactions, 186–187

in foaming chewing gum, 517

pH scale and, 525–533

strong acids and bases, 180–182, 180f

Acidic solution, 524

pH of, 532–533

Acree, Terry E., 383

Actinide series A group of fourteen

elements following actinium on the

periodic table, in which the 5f orbitals

are being filled, 343

Actinium, 343

Activation energy The threshold energy

that must be overcome to produce a

chemical reaction, 546

Air pollution, measurement of, 22, 22f

Alkali, 179

Alkali metal A Group 1 metal, 92

Alkaline earth metal A Group 2 metal, 92

Allotropes, 97

Alloy A substance that contains a mixture

of elements and has metallic properties,

463–464, 464n

Alpha (␣) particle A helium nucleus

produced in radioactive decay, 616

Alpha-particle production A common

mode of decay for radioactive nuclides in

which the mass number changes, 616

Altitude, 560

Aluminum

calculation of moles, 213–214



cation of, 99

distribution of, 76t

heat capacity of, 297t

ionic compound with oxygen, 367

mass calculations for, 254–256

1-mol sample of, 212t

symbol for, 79t

Aluminum chloride, naming of, 118

Aluminum iodide, mass calculations for,

254–256

Aluminum ion, formation of, 365t

Aluminum oxide

empirical formula for, 230

naming of, 123

Alvarez, Luis W., 1

Ammonia

equilibrium reaction and, 562, 563f

formation of, 268–271

molecular structure of, 385–387, 386f

phosphorus trichloride and, 563–564

reaction with copper, 271–273

reaction with oxygen, 156

synthesis of, 559–561, 560f

Ammonia gas, 156

ideal gas law and, 421–422

Ammonium chlorate, 131

Ammonium ion, polyatomic, 368

Ammonium nitrate, dissolving of,

476–477

Ammonium perchlorate, formula for, 134

Amphoteric substance The fundamental

unit of which elements are composed, 523

Analysis, dimensional, 30–34

Analytical balance, 21t

Anasazi Indians, 89

Anion A negative ion, 99–100

common simple, 117t

ionic bonding and, 368

in naming acids, 132–133, 133f

in naming compounds, 117

oxyanion, 129

Antacids, 261–263

Anthracite coal, 308t

Antimony, symbol for, 79t

Aqueous solution A solution in which

water is the dissolving medium or solvent,

166–202, 167–202, 475

describing reactions in, 177–179

equations for reactions, 178–179

precipitation reaction in, 167–177, 168f

predicting reaction, 167

predicting reaction in, 167

products forming in, 169–171

Argentium (silver), symbol for, 79t

Argon, symbol for, 79t

Argon gas, 424

Arnold, Kathryn E., 325

Arrhenius, Svante, 179–180

Arrhenius concept of acids and bases

A concept postulating that acids produce

hydrogen ions in aqueous solutions,

whereas bases produce hydroxide ions, 516

Arsenic

contamination with, 94

symbol for, 79t

Artificial sweetener, 383

Aspartame, 383



Asphalt, 307t

Atmosphere

carbon dioxide in, 309, 311, 311f

gases of, 403

greenhouse effects on, 309, 311, 311f

radiation and, 326

as unit of measure, 405, 406–407

Atmospheric pressure, 404, 404f

Atom The fundamental unit of which

elements are composed

calculating number of, 214–215

in compounds, 62

conserved in chemical reaction, 151

early models of, 83, 83n

ions of, 98–101, 101f

nuclear, 84

representation of, 59

size of, 350–351, 350f

structure of, 82–85

Atomic mass, 208–209, 209t

Atomic mass unit (amu) A small unit of

mass equal to 1.66 ϫ 1024 grams, 208

calculating mass using, 209

Atomic number (Z) The number of

protons in the nucleus of an atom; each

element has a unique atomic number,

86–88, 615, 615n

Atomic properties, periodic table and,

347–351, 350f

Atomic size, 350–351, 350f

Atomic solid A solid that contains atoms

at the lattice points, 459, 459f, 460f, 461,

463

Atomic structure, 85, 85t

chemical properties and, 85

electrons in, 83

of isotopes, 86–90, 86f

modern concept of, 85, 85f

neutron in, 85

of nuclear atom, 84

plum pudding model of, 83–84

proton in, 85

Atomic theory, 80, 322–357

Bohr model of, 331

Dalton’s, 80

electromagnetic radiation in, 324–327,

324f, 325f, 326f

electron configuration in, 338–346

emission of energy by atoms, 327–328

energy levels of hydrogen, 328–330,

329f, 330f

hydrogen orbitals in, 333–336, 333f,

334f, 335f

Rutherford’s model, 323–324, 324f

wave mechanical model of, 331–332,

336–338

Attractant, light as, 325, 325f

Aurium, symbol for, 79t

Average atomic mass, 208, 209t

Avogadro, Amadeo, 417

Avogadro’s law Equal volumes of gases at

the same temperature and pressure contain

the same number of particles (atoms or

molecules), 417–419, 417f

Avogadro’s number The number of atoms

in exactly 12 grams of pure 12C, equal to

6.022 ϫ 1023, 211



A57



A58 Index and Glossary

Baking soda, 261

Balance, electronic analytical, 21t

Balancing a chemical equation Making

sure that all atoms present in the reactants

are accounted for among the products,

147–157

Barium

distribution of, 76t

symbol for, 79t

Barium chromate, calculating mass of,

493–494, 494n

Barium nitrate, reaction with potassium

chromate, 168–169

Barium sulfate, suspension of, 568

Barometer A device for measuring

atmospheric pressure, 404–405

Base A substance that produces hydroxide

ions in aqueous solution; a proton acceptor,

180

conjugate, 516, 534

equivalent of, 497

formation of, 179–182, 180f

hydroxide ion produced by, 515

pH scale and, 525–533. see also pH scale

strength of, 520, 520f

water as, 523–525

Battery, 600–603, 601f

in hybrid car, 262

Benerito, Dr. Ruth Rogan, 4, 4f

Beryllium

electron configuration of, 339

as exception to octet rule, 380

Beryllium chloride

double bond of, 388–389

Lewis structure of, 382

Beta (␤) particle An electron produced in

radioactive decay, 616

Beta-particle production A decay process

for radioactive nuclides in which the mass

number remains constant and the atomic

number increases by one. The net effect is

to change a neutron to a proton, 616

Binary compound A two-element

compound, 116–123

classes of, 115

empirical formula for, 232–233

formulas for, 134–135

ionic, 368

ionic (type I), 115–119, 122–123

ionic (type II), 119–123, 126, 128–129

nonmetal (type III), 124–126, 128–129

Binary ionic compound A two-element

compound consisting of a cation and an

anion, 116. See also Ionic compound

Biomass, 327

Bismuth, symbol for, 79t

Bituminous coal, 308t

Bohr, Niels, 331, 331f

Bohr model of atom, 331, 331f

Boiling, heating to, 453

Boiling point, normal, 449

Bombardier beetle, 153

Bond The force that holds two atoms

together in a compound, 358–401. See also

Bonding

double, 376, 387–391, 388t

electronegativity and, 361–363, 362f,

362t, 363f

ionic, 368–369, 368f, 369f

Lewis structures, 370–382. see also Lewis

structure

molecular structure and, 381–382, 381f

polarity and dipole moments, 364, 364f

single, 376



stable electron configurations, 365–367,

365t, 367t

triple, 376

types of, 359–361, 361f

VSEPR model of, 382–387, 385f

Bond angle, 381, 381f

Bond energy The energy required to break a

given chemical bond, 360

Bond polarity, 361

Bonding

hydrogen, 454–456, 454f, 455f, 456f

intermolecular, 450, 450f

in metals, 463–464, 464n

in solids, 460–465, 461f, 461t, 462f,

463f

Bonding pair An electron pair found in the

space between two atoms, 371

Boron

electron configuration of, 339

1-mol sample of, 212t

symbol for, 79t

Boron trifluoride

as exception to octet rule, 380

Lewis structure of, 382–384

Box diagram, 338

Boyle, Robert, 75, 75f, 407

Boyle’s law The volume of a given sample

of gas at constant temperature varies

inversely with the pressure, 407–411, 408f

calculating pressure using, 410–411

calculating volume using, 409–410

Brain, PET scan of, 625

Breeder reactor A nuclear reactor in which

fissionable fuel is produced while the

reactor runs, 629

Broccoli, 377

Bromine

as diatomic molecule, 96, 96t

ions of, 100

Lewis structure of, 372

symbol for, 79t

Brønsted, Johannes, 516

Brønsted Lowry model A model

proposing that an acid is a proton donor

and that a base is a proton acceptor, 516

Buckminsterfullerine, 97, 97f

Buffer

characteristics of, 535

Buffered solution A solution where there is

a presence of a weak acid and its conjugate

base; a solution that resists a change in its

pH when either hydroxide ions or protons

are added, 534

Butane

formula for, 307t

Cadmium, symbol for, 79t

Calcium

distribution of, 76t

electron configuration of, 342–343

in human body, 77t

ionic compound with oxygen, 366–367

symbol for, 79t

Calcium carbonate

decomposition of, 556–557, 564

equilibrium reaction and, 561–562,

562f

Calcium chloride

formula for, 134

naming of, 123

Calcium fluoride, dissolving of, 568

Calculation

density in, 44–45

of energy requirements, 295–297



mass, 254–256

significant figures in, 27–29

specific heat capacity, 298–301

stoichiometric, 259–260

Calorie A unit of measurement for energy;

1 calorie is the quantity of energy required

to heat 1 gram of water by 1 Celsius

degree, 294–295

Calorimeter A device used to determine the

heat associated with a chemical or

physical change, 302

Car, hybrid, 262–263

Carbon

as atomic solid, 463

conversion of graphite to diamond,

304–305

distribution of, 76t

electron configuration of, 339

heat capacity of, 297t

in human body, 77t

isotopes of, 89

Lewis structure of, 371

symbol for, 79t

Carbon-14 dating, 623

Carbon dioxide

carbonation and, 521

climate effects of, 309, 311, 311f

double bonds of, 388–390

empirical formula for, 227–228

formation of, 150

global warming and, 375

green chemistry and, 479

greenhouse effect and, 326

Lewis structure of, 374–377

as molecular solid, 461

as pollutant, 403

reaction with lithium, 259–260

reaction with water, 547

sequestration of, 375

Carbon monoxide

as pollutant, 403

reaction with hydrogen, 273–275

reaction with steam, 551, 551f, 552f

Carbonation, 521

Carbonic anhydrase, 547

Carboxyl group The —COOH group in an

organic acid, 521

Catalyst A substance that speeds up a

reaction without being consumed, 547

Caterpillar, gypsy moth, 522

Cathode In a galvanic cell, the electrode at

which reduction occurs, 599–602

Cathodic protection The connection of an

active metal, such as magnesium, to steel

in order to protect the steel from corrosion,

604

Cation A positive ion, 99

common simple, 117t

common type II, 120t

ionic bonding and, 368

in naming compounds, 117

in solution, 476

Cell, fuel, 262–263

Celsius scale, 35–42

conversion from Fahrenheit, 41–42

conversion from Kelvin, 37–39

conversion to Fahrenheit, 39–41

conversion to Kelvin, 36–37

Chain reaction (nuclear) A selfsustaining fission process caused by the

production of neutrons that proceed to split

other nuclei, 627, 627f

Change of state, energy required for,

450–453, 450f, 450n