10: Electron Configurations and the Periodic Table
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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