Appendix C
A-9
Answers to Practice Problems
D
A P P E N D I X
Answers to Selected
Questions and Problems
Chapter 1 – Matter and Energy
1.1
1.3
1.5
1.7
1.9
1.11
1.13
1.15
1.17
1.19
1.21
1.23
1.25
1.27
1.29
1.31
1.33
1.35
1.37
1.39
1.41
1.43
1.45
(a) mass; (b) chemical property; (c) mixture; (d) element;
(e) energy; (f) physical property; (g) liquid; (h) density;
(i) homogeneous mixture; (j) solid
(a) 2.95 × 104; (b) 8.2 × 10−5; (c) 6.5 × 108;
(d) 1.00 × 10−2
(a) 0.0000186; (b) 10,000,000; (c) 453,000; (d) 0.0061
(a) 6.2 × 103; (b) 3.5 × 107; (c) 2.9 × 10−3;
(d) 2.5 × 10−7; (e) 8.20 × 105; (f) 1.6 × 10−6
(a) 3; (b) 2; (c) 4; (d) 2; (e) 3
(a) 1.5; (b) 1.5; (c) 14; (d) 1.20
(a) 2.8; (b) 0.28; (c) 2.8; (d) 0.049
(a) 1.21; (b) 0.204; (c) 1.84; (d) 42.2; (e) 0.00710
(a) 0.036 m; (b) 3.57 × 105 g; (c) 0.07650 L; (d) 8.4670 cm;
(e) 5.97 × 10−7 m; (f) 92.7 cm; (g) 7.62 × 104 g;
(h) 865 L; (i) 17.5 in; (j) 0.5214 lb; (k) 2.1 qt
(a) 0.589 mi; (b) 14.9 lb; (c) 6.617 × 10−2 gal;
(d) 2.30 × 102 mL; (e) 4.50 × 103 nm; (f) 0.952 m;
(g) 3.12 × 102 kg; (h) 5 × 102 mL; (i) 4.10 ft;
(j) 1.19 × 10−3 lb; (k) 6.6 × 10−9 gal
(a) 7.38 × 104 ft/min; (b) 1.50 in3; (c) 0.697 lb/in3
(a) homogeneous mixture if the dye is evenly mixed into
the water; (b) element; (c) homogeneous mixture;
(d) heterogeneous mixture
Gasoline, automobile exhaust, oxygen gas, and the iron
pipe are matter. Sunlight is energy.
Elements are composed of only one type of atom. Compounds are made up of two or more different elements.
Metals are lustrous (shiny) and conduct heat and electricity. In addition you can form wires with metals (ductile)
and you can make foil out of them by hitting them with a
hammer (malleable).
(a) titanium; (b) tantalum; (c) thorium; (d) technetium;
(e) thallium
(a) boron; (b) barium; (c) beryllium; (d) bromine;
(e) bismuth
(a) nitrogen; (b) iron; (c) manganese; (d) magnesium;
(e) aluminum; (f) chlorine
(a) Fe; (b) Pb; (c) Ag; (d) Au; (e) Sb
Ir is the symbol for the element iridium. Iron’s symbol is
Fe.
The correct symbol is No.
The only pure substance is the salt in the salt shaker (if it
is not iodized salt). Hamburger: heterogeneous mixture.
Salt: pure substance. Soft drink: heterogeneous mixture
(until it goes flat). Ketchup: heterogeneous mixture.
1.47
1.49
1.51
The chemical formula is N2O4.
Image D represents a mixture of an element and a compound.
The elements are O2, P4, and He. Fe2O3, NaCl, and H2O
are compounds.
1.53
Liquid state
1.55
1.57
1.59
1.61
1.63
1.65
1.67
1.69
1.71
1.73
1.75
1.77
1.79
1.81
1.83
gas
(a) gas; (b) liquid; (c) solid
solid
O2(aq)
physical properties
(a) 0.045 g; (b) 1.6 × 10−3 oz; (c) 9.9 × 10−5 lb
(a) 0.10 mg; (b) 1.0 × 102 μg; (c) 1.0 × 10−7 kg
(a) 1.2 × 103 mL; (b) 1.2 × 103 cm3; (c) 1.2 × 10−3 m3
1.6 × 102 mL, 0.16 L
1.3 g/cm3, 1.3 g/mL
38.5 mL
The molecules in the liquid state are closer together than
molecules in the gas state. More matter in the same volume means that the density is higher.
oil (least) < plastic < water (greatest)
329 K
Scale
1.89
Freezing
Boiling
Difference
Celsius
0°C
100°C
100°C
Kelvin
213.15 K
373.15 K
100 K
32°F
212°F
180°F
Fahrenheit
1.85
1.87
Solid state
no
Physical properties are (a) mass, (b) density, and (e) melting point. Chemical properties are (c) flammability,
(d) resistance to corrosion, and (f) reactivity with water.
Physical changes are (a) boiling acetone, (b) dissolving
oxygen gas in water, and (e) screening rocks from sand.
Chemical changes are (c) combining hydrogen and oxygen
to make water, (d) burning gasoline, and (f) conversion of
ozone to oxygen.
H2
A-9
Appendix D
A-10
1.91
Symbolic: Cl2(g)
gram is shown.
Answers to Selected Questions and Problems
Cl2(l). A molecular-level dia-
1.121
Condensation
Gas
1.93
1.95
1.119
Liquid
chemical change
physical change
1.123
1.125
1.127
1.129
1.131
You could do some library research and find the densities
of many different woods. By placing samples of the different woods in water, you can determine if the theory is
correct.
At high altitude, the air pressure is lower. As a result, when
a balloon rises, it expands. Since the mass of air in the
balloon has not changed but the volume has increased, the
density of the balloon is lower.
zinc
(a) He; (b) Ne; (c) Ar; (d) Kr; (e) Xe; (f) Rn
physical
8 mg
1.6 × 10−4 lb/oz; 0.02 to 0.03 lb
Chapter 2 – Atoms, Ions, and the Periodic Table
2.1
2.3
2.5
2.7
1.97
1.99
1.101
1.103
1.105
1.107
1.109
1.111
1.113
1.115
1.117
physical change
Anna and Bill would have observed kinetic energy from
the movement of the welder and the motion of the sparks.
The sparks would have glowed, indicating heat, light, and
chemical energy.
The molecules in image A have greater kinetic energy
because they are moving faster.
Any object that would move if allowed has potential
energy (e.g., a picture hanging on a wall).
The people walking, the wheel chair rolling, and the suitcase being pushed all have kinetic energy. The people, the
wall art, and objects on the tables all have potential energy.
Many objects have potential and kinetic energy.
Consider as an example a car going down the road. A car
going up a hill is converting chemical energy in the fuel
to mechanical energy to reach the top. At the top of the
hill, the car has potential energy. If it rolls down the hill,
it gains kinetic energy but loses potential energy. Other
energy conversions can be found while driving a car.
When the batter swings the bat, potential energy (metabolic energy) is converted into kinetic energy (moving
bat). The bat strikes the ball, and the kinetic energy of the
bat is transferred to the kinetic energy of the ball. As the
ball leaves the bat, it rises against the Earth’s gravity and
kinetic energy is converted to potential energy. When the
ball starts dropping again, potential energy is converted to
kinetic energy.
BMI = 21.7 kg/m2; healthy
As the water leaves the top of the fountain, it possesses
kinetic energy (going up). That energy is converted to
potential energy. As the water falls, the potential energy
converts back into kinetic energy.
It is used to explain and predict scientific results.
(a) hypothesis; (b) observation; (c) theory; (d) observation
2.9
2.11
2.13
2.15
(a) neutron; (b) law of conservation of mass; (c) proton;
(d) main-group element; (e) relative atomic mass;
(f) mass number; (g) isotope; (h) cation; (i) subatomic particle; (j) alkali metal; (k) periodic table
Dalton used the laws of conservation of mass (Lavoisier)
and definite proportions (Proust).
They differ in their atomic masses and chemical properties.
Compounds contain discrete numbers of atoms of each
element that form them. Because all the atoms of an element have the same relative atomic mass, the mass ratio
of the elements in a compound is always the same (law of
definite proportions).
No. Hydrogen atoms are not conserved. A molecule of
hydrogen, H2, should be added to the left panel.
Thomson’s cathode-ray experiment
electrons
The nucleus of helium has two protons and two neutrons.
Two electrons can be found outside the nucleus.
proton
neutron
2 e–
2.17
2.19
2.21
2.23
2.25
2.27
neutron
Carbon has six protons. The relative atomic mass of a
carbon atom is 12.01 amu indicating the presence of six
neutrons. Protons and neutrons have approximately equal
masses, so the nuclear mass is approximately two times
the mass of the protons.
(a) 1; (b) 8; (c) 47
protons
number of protons
(b) atomic number
A-11
Appendix D Answers to Selected Questions and Problems
2.29
2.31
The following table displays the atomic, neutron, and mass
numbers for the isotopes of hydrogen.
1
1H
2
1H
3
1H
Atomic number
1
1
1
Neutron number
0
1
2
Mass number
1
2
3
(a) 158 O
(b) 109
47 Ag
35
(c) 17 Cl
Protons
Neutrons
Electrons
8
7
8
47
62
47
17
18
17
Protons
Neutrons
(a) 56
26 Fe
(b) 39
19 K
19
(c) Copper-65 or 65
29 Cu
29
26
2.47
Electrons
(a)
28
30
27
20
(b)
60
+
28 Ni
28
32
27
36
+
(c) 61
28 Ni
28
33
27
(d)
62
+
28 Ni
28
34
27
(e)
64
+
28 Ni
28
36
27
30
Electrons
(a) 23
11Na
11
12
11
(b) 56
25 Mn
25
31
25
8
10
8
9
10
9
They differ in the number of electrons.
(a) an anion with a 1- charge is formed; (b) a cation with
a 2+ charge is formed
(a) Zn2+, cation; (b) P3−, anion
(a) Zn2+
Protons
Electrons
30
28
(b) F-
9
10
+
1
0
(c) H
2.51
2.81
2.83
2.85
2.87
2.89
2.91
2.93
2.95
2.97
2.99
2.101
2.103
2.105
2.107
Protons
Neutrons
Electrons
(a)
37
17 Cl
17
20
18
2.109
(b)
25
2+
12 Mg
12
13
10
2.111
(c)
13 37N
7
6
10
20
20
18
2+
(d) 40
20 Ca
2.59
2.61
Neutrons
Neutrons
2.49
2.53
2.55
2.57
Protons
58
+
28 Ni
Protons
(c) 188O
(d) 199 F
(a) 2 amu; (b) 238 amu
The mass of D2 is two times the mass of H2.
about 40 amu
The numerical values of masses of individual atoms are
very small when measured on the gram scale. The size of
the atomic mass unit allows us to make easier comparisons
and calculations of masses of molecules.
The mass number is the sum of the number of protons and
neutrons and is always an integer value. In contrast, the
mass of an atom is the actual measurement of how much
matter is in the atom and is never exactly an integer value
(except carbon-12).
A mass spectrometer is used to determine the mass of an
atom. The mass number of an atom is the sum of the number of protons and neutrons.
calcium-40
21.50 amu
(a) 58Ni; (b) 64Ni; (c) 58;
(d)
seven protons and six neutrons
2.41
2.43
2.45
2.73
2.75
2.77
2.79
(a) 31 H; (b) 94 Be; (c) 31
15 P
2.37
2.39
2.71
(a) Z = 18, N = 18, A = 36; (b) Z = 18, N = 20, A = 38;
(c) Z = 18, N = 22, A = 40
2.33
2.35
2.63
2.65
2.67
2.69
potassium, K
copper, Cu
7
Li has three protons, three electrons, and four neutrons.
7 +
Li has only two electrons, and 6Li has three neutrons.
Otherwise they are the same as 7Li. Lithium-6 differs the
most in mass.
19 protons and 18 electrons
The atomic mass unit is defined as one-twelfth the mass of
one carbon-12 atom.
2.113
2.115
2.117
2.119
2.121
10,810 amu or 1.081 × 104 amu
2500 amu of boron contains more atoms.
(a) K; (b) Br; (c) Mn; (d) Mg; (e) Ar; (f) Br, K, Mg, Al, Ar
chlorine, Cl
titanium, Ti
(a) metal; (b) nonmetal; (c) metal; (d) metalloid
(a) main group; (b) main group; (c) main group;
(d) actinide; (e) transition metal
Group VIIA, the halogen family, all occur as diatomic
molecules.
neon
Group VIIIA (18), the noble gases
electrons
(a) group IA (1); (b) group IIA (2); (c) group VIIA (17);
(d) group VIA (16)
(a) Na+; (b) O2−; (c) S2−; (d) Cl−; (e) Br−
All alkali metal elements (group 1 or IA excluding
hydrogen)
The mass of oxygen added to form Fe2O3 causes an
increase in the mass.
The mass ratios of Zn/S are the same for the two samples
(within the significant figures given). The mass ratio is
approximately 2.0:1.0.
Electrons have charge and were readily studied in
cathode-ray tubes.
nickel-60
19 protons and 20 neutrons
The most abundant isotopes of cobalt have masses greater
than the masses of the most abundant isotopes of nickel.
As a result, the relative atomic mass of cobalt is greater
than that of nickel.
When there are many isotopes, some of the isotopes can be
present in very low abundance. As a result, their masses
cannot be determined as accurately and their percentage
A-12
2.123
2.125
2.127
2.129
2.131
2.133
Appendix D Answers to Selected Questions and Problems
contributions are also less well known. Both factors result
in a decrease in the precision of the calculated relative
atomic mass.
127
carbon
20% boron-10 and 80% boron-11; the relative atomic mass
(10.81 amu) is about 80% of the difference between the
two isotopes.
Br2(l)
hydrogen
The energy required to force a proton (positive charge)
into the nucleus is too great.
3.45
3.47
3.49
3.51
3.53
(a) CaSO4; (b) BaO; (c) (NH4)2SO4; (d) BaCO3;
(e) NaClO3
(a) Co2+, cobalt(II) chloride; (b) Pb4+, lead (IV) oxide or
plumbic oxide; (c) Cr3+, chromium(III) nitrate; (d) Fe3+,
iron(III) sulfate or ferric sulfate
(a) CoCl2; (b) Mn(NO3)2; (c) Cr2O3; (d) Cu3(PO4)2
ferrous nitrate and ferric nitrate
Chapter 3 – Chemical Compounds
3.1
3.3
3.5
3.7
3.9
3.11
3.13
3.15
3.17
3.19
3.21
3.23
3.25
Ca2+
Fe2+
K+
Cl–
CaCl2
calcium
chloride
FeCl2
iron(II)
chloride
KCl
potassium
chloride
–
CaO
calcium
oxide
FeO
iron(II)
oxide
K2O
potassium
oxide
–
Ca(NO3)2
calcium
nitrate
Fe(NO3)2
iron(II)
nitrate
KNO3
potassium
nitrate
SO32
–
CaSO3
calcium
sulfite
FeSO3
iron(II)
sulfite
K2SO3
potassium
sulfite
OH–
Ca(OH)2
calcium
hydroxide
Fe(OH)2
iron(II)
hydroxide
KOH
potassium
hydroxide
ClO3– Ca(ClO3)2
calcium
chlorate
Fe(ClO3)2
iron(II)
chlorate
KClO3
potassium
chlorate
O2
(a) formula unit; (b) strong electrolyte; (c) molecular compound; (d) acid; (e) nonelectrolyte; (f) oxoanion
(a) ionic; (b) ionic; (c) molecular
(a) molecular; (b) ionic; (c) molecular; (d) ionic
(a) molecular; (b) ionic; (c) ionic; (d) molecular
The ionic compound LiF would have the highest melting
point.
(a) Na+, sodium ion; (b) K+, potassium ion; (c) Rb+,
rubidium ion
(a) Ca2+, calcium ion; (b) N3−, nitride ion; (c) S2−,
sulfide ion
NO2−, nitrite ion
(a) sulfate ion; (b) hydroxide ion; (c) perchlorate ion
(a) N3−; (b) NO3−; (c) NO2−
(a) CO32−; (b) NH4+; (c) OH−; (d) MnO4−
SO3−, sulfite ion
IO3−, iodate ion
NO3
3.27
Mn2+
Al3+
NH4+
Cl–
MnCl2
manganese(II)
chloride
AlCl3
aluminum
chloride
NH4Cl
ammonium
chloride
–
MnO
manganese(II)
oxide
Al2O3
aluminum
oxide
(NH4)2O
ammonium
oxide
–
Mn(NO3)2
manganese(II)
nitrate
Al(NO3)3
aluminum
nitrate
NH4NO3
ammonium
nitrate
SO32
–
MnSO3
manganese(II)
sulfite
Al2(SO3)3
aluminum
sulfite
(NH4)2SO3
ammonium
sulfite
OH–
Mn(OH)2
manganese(II)
hydroxide
Al(OH)3
aluminum
hydroxide
NH4OH
ammonium
hydroxide
ClO3–
Mn(ClO3)2
manganese(II)
chlorate
Al(ClO3)3
aluminum
chlorate
NH4ClO3
ammonium
chlorate
O2
NO3
3.29
3.31
3.33
3.35
3.37
3.39
3.41
3.43
(a) BaCl2; (b) FeBr3; (c) Ca3(PO4)2; (d) Cr2(SO4)3
(a) K+ and Br−; (b) Ba2+ and Cl −; (c) Mg2+ and PO43−;
(d) Co2+ and NO3−
(a) Fe2+ will form FeO and Fe3+ will form Fe2O3. (b) Fe2+
will form FeCl2 and Fe3+ will form FeCl3.
(a) 2−; (b) 2+
For each “compound” written, the charges do not balance
(compounds have no net charge). (a) too many chloride
ions, NaCl; (b) not enough potassium ions, K2SO4;
(c) There should be three nitrate ions and one aluminum
ion, Al(NO3)3.
(a) magnesium chloride; (b) aluminum oxide; (c) sodium
sulfide; (d) potassium bromide; (e) sodium nitrate;
(f) sodium perchlorate
MnSO4 and CoCl2
(a) copper(I) oxide or cuprous oxide; (b) chromium(II)
chloride; (c) iron(III) phosphate or ferric phosphate;
(d) copper(II) sulfide or cupric sulfide
3.55
potassium
iron(III)
strontium
iodide
KI
FeI3
SrI3
oxide
K2O
Fe2O3
SrO
sulfate
K2SO4
Fe2(SO4)3
SrSO4
nitrite
KNO2
Fe(NO2)3
Sr(NO2)2
acetate
KCH3CO2
Fe(CH3CO2)3
Sr(CH3CO2)2
KClO
Fe(ClO)3
Sr(ClO)2
hypochlorite
Appendix D Answers to Selected Questions and Problems
aluminum
cobalt(II)
lead(IV)
3.113
iodide
AlI3
CoI2
PbI4
3.115
oxide
Al2O3
CoO
PbO2
sulfate
Al2(SO4)3
CoSO4
Pb(SO4)2
nitrite
Al(NO2)3
Co(NO2)2
Pb(NO2)4
acetate
Al(CH3CO2)3
Co(CH3CO2)2
Pb(CH3CO2)4
hypochlorite
Al(ClO)3
Co(ClO)2
Pb(ClO)4
3.57
AgCl
3.59
NF3, P4O10, C2H4Cl2
3.61
(a) phosphorus pentafluoride; (b) phosphorus trifluoride;
(c) carbon monoxide; (d) sulfur dioxide
3.63
(a) SF4; (b) C3O2; (c) ClO2; (d) SO2
3.65
The central image represents H3PO4.
3.67
(a) hydrofluoric acid; (b) nitric acid; (c) phosphorous acid
3.69
(a) HF; (b) H2SO3; (c) HClO4
3.71
hydrogen ions and nitrate ions
3.73
KI, Mg(NO3)2, NH4NO3
3.75
ionic; TiO2, ZnO
3.77
molecular; CO2 and N2O
3.79
potassium sulfide; sodium sulfate, sulfur dioxide
3.81
(a) nitrogen trioxide, molecular; (b) nitrate ion, an ion;
(c) potassium nitrate, ionic; (d) sodium nitride, ionic;
(e) aluminum chloride, ionic; (f) phosphorus trichloride,
molecular, (g) titanium(II) oxide, ionic; (h) magnesium
oxide, ionic
3.83
(a) Na2CO3; (b) NaHCO3; (c) H2CO3; (d) HF; (e) SO3;
(f) CuSO4; (g) H2SO4; (h) H2S
3.85
HCl(aq) is dissolved in water and ionized; HCl(g) is not
ionized; HCl(aq), hydrochloric acid; HCl(g) hydrogen
chloride
3.87
(a) No prefixes with ionic compounds. (b) Calcium’s
charge should not be stated. (c) Copper’s charge,
2+, should be stated. (d) Prefixes are used for
molecular compounds.
3.89
(a) Two potassium ions are needed to balance S2−. (b) Only
one Co2+ is needed to balance charge. (c) Nitride is N3−.
(d) The prefix tri is associated with iodide, NI3.
3.91
(a) one sodium ion, Na+, and one chloride ion, Cl−
(b) one magnesium ion, Mg2+, and two chloride ions, Cl−
(c) two sodium ions, Na+, and one sulfate ion, SO42−
(d) one calcium ion, Ca2+, and two nitrate ions, NO3−
3.93
(a) electrolyte; (b) electrolyte; (c) electrolyte; (d) nonelectrolyte
3.95
The ions of silver, zinc, and cadmium can each only have
one possible charge. They are assumed to have these
charges in the compounds they form.
3.97
(a) NO3−; (b) SO32−; (c) NH4+; (d) CO32−; (e) SO42−;
(f) NO2−; (g) ClO4−
3.99
(a) magnesium bromide; (b) hydrogen sulfide; (c) hydrosulfuric acid; (d) cobalt(III) chloride; (e) potassium
hydroxide; (f) silver bromide
3.101 (a) PbCl2; (b) Mg3(PO4)2; (c) NI3; (d) Fe2O3; (e) Ca3N2;
(f) Ba(OH)2; (g) Cl2O5; (h) NH4Cl
3.103 NaHCO3
3.105 Ca(ClO)2
3.107 H2O(l)
3.109 Cu, AgNO3, Ag, Cu(NO3)2
3.111 (a) NH3; (b) HNO3(aq); (c) HNO2(aq)
A-13
Hydrogen atoms attached to oxygen atoms are responsible
for the acidic properties.
All contain oxygen. Magnesium oxide, MgO, has metal
and nonmetal ions which makes it an ionic compound. It
is a solid at room temperature. Oxygen and carbon dioxide
(O2 and CO2) are both molecular substances and gases at
room temperature.
Chapter 4 – Chemical Composition
4.1
4.3
4.5
4.7
4.9
4.11
4.13
4.15
4.17
4.19
4.21
4.23
4.25
4.27
4.29
4.31
4.33
4.35
4.37
4.39
4.41
4.43
4.45
4.47
4.49
4.51
4.53
4.55
4.57
4.59
4.61
4.63
4.65
(a) mole; (b) Avogadro’s number; (c) empirical formula;
(d) solute; (e) molarity; (f) concentrated solution
40.0% calcium
60.0% carbon
0.226 g lithium
Chemical formulas must have whole-number subscripts.
(a) H2S; (b) N2O3; (c) CaCl2
H2SO4, SCl4, C2H4
CO2
A formula unit describes the simplest unit of an ionic or
network solid. Sodium chloride, NaCl; chlorine, Cl2;
methane, CH4; silicon dioxide, SiO2.
3.0 × 1023 molecules; 3.0 × 1023 N atoms;
9.0 × 1023 H atoms
3.011 × 1023 formula units
1 × 1023 atoms
6.022 × 1023 calcium ions
NaCl, 58.44 g/mol; CH4, 16.04 g/mol; Cl2, 70.90 g/mol;
SiO2, 60.09 g/mol
(a) 472.1 amu; (b) 172.18 amu; (c) 150.90 amu;
(d) 120.07 amu
(a) 253.8 g/mol; (b) 158.35 g/mol; (c) 56.10 g/mol
To measure out a useful number of atoms by counting
would not be possible because atoms are too small for us
to see and manipulate.
42.394 amu
35.9 g/mol
(a) 0.0999 mol; (b) 0.293 mol; (c) 0.127 mol;
(d) 0.0831 mol
(a) 0.556 mol; (b) 1.388 × 10-3 mol; (c) 733 mol;
(d) 5.72 × 10-8 mol
Na
(a) 428 g; (b) 177 g; (c) 436 g; (d) 2.20 × 102 g
5.0 × 102 g
(a) 1.76 mol; (b) 1.06 × 1024 molecules; (c) 1.06 × 1024 N
atoms; (d) 5.28 mol of H atoms
H2SO4 has the most atoms per mole because it has more
atoms per molecule. Na has the least atoms per mole.
1.7 × 1021 molecules
(a) 9.421 × 1023 formula units; (b) 1.581 × 1024 formula
units; (c) 8.356 × 1024 formula units; (d) 2.696 × 1024
formula units
(a) 8.364 × 1025 atoms; (b) 1.541 × 1024 ions;
(c) 5.619 × 1024 atoms; (d) 1.451 × 1024 ions
6.8 g SO2
NH3 (82%)
no
The molecular formula shows the exact numbers of each
atom present in a compound. The empirical formula
shows the relative amounts of each atom in a compound
expressed as small whole numbers.
H2O2 (empirical, HO); N2O4 (empirical, NO2)
A-14
4.67
4.69
4.71
4.73
4.75
4.77
4.79
4.81
4.83
4.85
4.87
4.89
4.91
4.93
4.95
4.97
4.99
4.101
4.103
4.105
4.107
4.109
4.111
4.113
4.115
4.117
4.119
4.121
4.123
4.125
4.127
4.129
Appendix D Answers to Selected Questions and Problems
(a) P2O5; (b) same as molecular; (c) same as molecular;
(d) C3H5O2
(a) C3H2Cl; (b) same as molecular; (c) same as molecular
both NO2 and N2O4
(a) Fe3O4; (b) C6H5NO2
C5H6O
C7H5N3O6
percent composition by mass and the molar mass
C3H6O3
C6H12O6
(a) 50.05% S, 49.95% O; (b) 47.27% Cu, 52.73% Cl;
(c) 42.07% Na, 18.89% P, 39.04% O; (d) 16.39% Mg,
18.89% N, 64.72% O
Chalcocite (Cu2S) and cuprite (CuO) are both relatively
high in copper content (79.85% and 79.88%, respectively).
A solution is a homogeneous mixture of two or more substances. Some common solutions: clear drinks (coffee and
tea), window cleaner, soapy water, tap water, air, brass (a
homogeneous mixture of copper and zinc).
Water is the solvent because it is present in the largest
amount. Calcium chloride, CaCl2, is the solute.
Concentrated means high in solute concentration, and
dilute means low in solute concentration.
Concentration describes the quantity of solute in a given
amount of solvent or solution.
solution A
(a) 2.03 M; (b) 0.540 M; (c) 3.20 M; (d) 3.21 M
0.0186 mol Na2SO4; 0.0372 mol Na+; 0.0186 mol SO42(a) 0.375 moles, 28.0 g; (b) 0.512 mol, 72.8 g
(a) 1.00 L; (b) 0.0833 L
40.0 mL
69.2 mL
(a) 0.01814 M; (b) 0.2974 M; (c) 0.04020 M
9.76 g
39.95 amu, 6.634 × 10-23 g
(a) 0.904 mol, 5.44 × 1023 atoms; (b) 0.741 mol,
4.46 × 1023 atoms; (c) 0.169 mol, 1.02 × 1023 atoms
7.493 × 10-3 mol
C6H11OBr
2.458 × 1023 oxygen atoms
1.31 × 1025 molecules CO2
(a) C8H8O3; (b) C8H8O3
Al2O3
Chapter 5 – Chemical Reactions and Equations
5.1
5.3
5.5
5.7
(a) single-displacement reaction; (b) anhydrous; (c) molecular equation; (d) decomposition reaction; (e) balanced
equation; (f) reactant; (g) spectator ion; (h) combustion;
(i) precipitate
The reactants are aluminum and oxygen gas. The product
is aluminum oxide.
Image A represents the reactants, and image C represents
the products.
The numbers of hydrogen atoms do not match. One
hydrogen molecule should be added to the reactant image.
5.9
The product image should show five XeF2 molecules and
one unreacted Xe atom.
5.11
5.13
5.15
5.17
5.19
5.21
5.23
5.25
5.27
5.29
5.31
Three major signs are visible: (1) a brown gas is formed,
(2) bubbles, (3) color change.
It is not a chemical change because no new substance is
formed.
New substances are formed, so a chemical reaction has
taken place.
No new chemical substances have formed, so no chemical
reaction has taken place.
A chemical equation is a chemist’s way of showing what
happens during a chemical reaction. It identifies the formulas for the reactants and products and demonstrates how
mass is conserved during the reaction.
(a) chemical reaction; (b) physical change; (c) chemical
reaction
Balancing a chemical equation demonstrates how mass is
conserved during a reaction. This makes the equation a
quantitative tool for determining the amount of reactant
used and product produced.
(a) NaH(s) + H2O(l)
H2(g) + NaOH(aq)
(b) 2Al(s) + 3Cl2(g)
2AlCl3(s)
N2(g) + 3Cl2(g)
2NCl3(g)
Image B: 2Mg(s) + O2(g)
2MgO(s)
5.33
5.35
5.37
5.39
(a) 2Al(s) + 3Cl2(g)
2AlCl3(s)
(b) Pb(NO3)2(aq) + K2CrO4(aq)
PbCrO4(s) + 2KNO3(aq)
(c) 2Li(s) + 2H2O(l)
2LiOH(aq) + H2(g)
(d) 2C6H14(g) + 19O2(g)
12CO2(g) + 14H2O(l)
(a) CuCl2(aq) + 2AgNO3(aq)
Cu(NO3)2(aq) + 2AgCl(s)
(b) S8(s) + 8O2(g)
8SO2(g)
(c) C3H8(g) + 5O2(g)
3CO2(g) + 4H2O(g)
Appendix D Answers to Selected Questions and Problems
5.41
5.43
5.45
5.47
5.49
5.51
5.53
5.55
5.57
5.59
5.61
5.63
5.65
5.67
5.69
5.71
5.73
5.75
5.77
5.79
5.81
Cu(NO3)2(aq) + 2Ag(s)
Cu(s) + 2AgNO3(aq)
decomposition: one reactant, more than one product
combination: more than one reactant, one product
single-displacement: element and compound as reactants,
different element and compound as products
double-displacement: two compounds as reactants, two
compounds as products
(a) combination; (b) single-displacement;
(c) decomposition
double-displacement
(a) combination; (b) single-displacement
(a) CaCl2(aq) + Na2SO4(aq)
CaSO4(s) + 2NaCl(aq): double-displacement
(b) Ba(s) + 2HCl(aq)
BaCl2(aq) + H2(g): singledisplacement
(c) N2(g) + 3H2(g)
2NH3(g): combination
(d) FeO(s) + CO(g)
Fe(s) + CO2(g): not classified
(e) CaO(s) + H2O(l)
Ca(OH)2(aq): combination
(f) Na2CrO4(aq) + Pb(NO3)2(aq)
PbCrO4(s) + 2NaNO3(aq): double-displacement
(g) 2KI(aq) + Cl2(g)
2KCl(aq) + I2(aq): single-displacement
(h) 2NaHCO3(s)
Na2CO3(s) + CO2(g) + H2O(g):
decomposition
NiCO3(s)
NiO(s) + CO2(g)
(a) CaCO3(s)
CaO(s) + CO2(g)
(b) CuSO4⋅5H2O(s)
CuSO4(s) + 5H2O(g)
2Mg(s) + O2(g)
2MgO(s)
(a) 3Ca(s) + N2(g)
Ca3N2(s)
(b) 2K(s) + Br2(l)
2KBr(s)
(c) 4Al(s) +3O2(g)
2Al2O3(s)
(a) Zn(s) + 2AgNO3(aq)
2Ag(s) + Zn(NO3)2(aq)
(b) 2Na(s) + FeCl2(s)
2NaCl(s) + Fe(s)
Zn(s) + SnCl2(aq)
ZnCl2(aq) + Sn(s)
(a) Ca reacts with water: Ca(s) + 2H2O(l)
Ca(OH)2(aq) + H2(g)
Ca reacts with HCl: Ca(s) + 2HCl(aq)
CaCl2(aq) + H2(g)
(b) Fe reacts with HCl: Fe(s) + 2HCl(aq)
FeCl2(aq) + H2(g)
(c) no reaction
(a) yes; (b) no; (c) no; (d) yes
(a) soluble; (b) soluble; (c) insoluble; (d) insoluble
(a) K2CO3(aq) + BaCl2(aq)
BaCO3(s) + 2KCl(aq)
(b) CaS(s) + Hg(NO3)2(aq)
Ca(NO3)2(aq) + HgS(s)
(c) Pb(NO3)2(aq) + K2SO4(aq)
PbSO4(s) + 2KNO3(aq)
(a) CaCO3(s) + H2SO4(aq)
CaSO4(s) + CO2(g) + H2O(l)
(b) SnCl2(aq) + 2AgNO3(aq)
Sn(NO3)2(aq) + 2AgCl(s)
lead(II) sulfate
CaCl2(aq) + K2CO3(aq)
CaCO3(s) + 2KCl(aq)
(a) precipitation of HgS(s)
(b) formation of the insoluble gas H2S(g)
(c) formation of the stable molecular compound H2O(l)
and precipitate BaSO4(s)
(a) H2S(aq) + Cu(OH)2(s)
CuS(s) + 2H2O(l)
(b) no reaction
(c) KHSO4(aq) + KOH(aq)
K2SO4(aq) + H2O(l)
5.83
5.85
5.87
5.89
5.91
5.93
5.95
5.97
5.99
5.101
5.103
5.105
5.107
5.109
5.111
5.113
5.115
5.117
5.119
A-15
O2(g)
(a) Cs2O(s); (b) PbO(s) or PbO2(s); (c) Al2O3(s);
(d) H2O(g); (e) CO(g) or CO2(g)
(a) CO(g) or CO2(g) and H2O(g); (b) CO2(g); (c) Al2O3(s);
(d) CO(g) or CO2(g) and H2O(g)
An electrolyte produces ions when dissolved in water. A
nonelectrolyte does not produce ions in water.
(a) and (b) are electrolytes; (c) nonelectrolyte
nonelectrolyte
(a) no ions; (b) no ions; (c) Na+(aq) and Cl-(aq)
Molecular: all substances written as complete chemical
formulas or atoms
Ionic: all soluble salts, strong acids and bases are written
as ions
Net ionic: all spectator ions are removed from the ionic
equation
ions that do not participate in a reaction
(a) 2NaCl(aq) + Ag2SO4(s)
Na2SO4(aq) + 2AgCl(s)
2Cl-(aq) + Ag2SO4(s)
2AgCl(s) + SO42-(aq)
(b) Cu(OH)2(s) + 2HCl(aq)
CuCl2(aq) + 2H2O(l)
Cu(OH)2(s) + 2H+(aq)
Cu2+(aq) + 2H2O(l)
(c) BaCl2(aq) + Ag2SO4(s)
BaSO4(s) + 2AgCl(s)
Ba2+(aq) + 2Cl-(aq) + Ag2SO4(s)
BaSO4(s) + 2AgCl(s)
(a) calcium carbonate, CaCO3
(b) Na2CO3(aq) + CaBr2(aq)
CaCO3(s) + 2NaBr(aq)
(c) 2Na+(aq) + CO32-(aq) + Ca2+(aq) + 2Br-(aq)
CaCO3(s) + 2Na+(aq) + 2Br-(aq)
(d) Na+(aq) and Br-(aq)
(e) CO32-(aq) + Ca2+(aq)
CaCO3(s)
Ag+(aq) + Cl-(aq)
AgCl(s)
2Na(s) + 2H2O(l)
2Na+(aq) + 2OH-(aq) + H2(g)
(a) Cu(s) + 2AgNO3(aq)
Cu(NO3)2(aq) + 2Ag(s)
Cu(s) + 2Ag+(aq)
Cu2+(aq) + 2Ag(s)
(b) FeO(aq) + 2HCl(aq)
FeCl2(aq) + H2O(l)
O2-(aq) + 2H+(aq)
H2O(l)
(a) Sr(NO3)2(aq) + H2SO4(aq)
SrSO4(s) + HNO3(aq)
Sr2+(aq) + SO42-(aq)
SrSO4(s)
(b) no reaction
(c) CuSO4(aq) + BaS(aq)
CuS(s) + BaSO4(s)
+
Cu2 (aq) + SO42-(aq) + Ba2+(aq) + S2-(aq)
CuS(s) + BaSO4(s)
(d) NaHCO3(aq) + CH3CO2H(aq)
CO2(g) + H2O(l) + NaCH3CO2(aq)
HCO3-(aq) + CH3CO2H(aq)
CO2(g) + H2O(l) + CH3CO2-(aq)
combination
(a) ZnSO4(aq) + Ba(NO3)2(aq)
BaSO4(s) + Zn(NO3)2(aq)
(b) 3Ca(NO3)2(aq) + 2K3PO4(aq)
Ca3(PO4)2(s) + 6KNO3(aq)
(c) ZnSO4(aq) + BaCl2(aq)
BaSO4(s) + ZnCl2(aq)
(d) 2KOH(aq) + MgCl2(aq)
2KCl(aq) + Mg(OH)2(s)
(e) CuSO4(aq) + BaS(aq)
CuS(s) + BaSO4(s)
2K(s) + 2H2O(l)
2KOH(aq) + H2(g)
(b)
Appendix D Answers to Selected Questions and Problems
A-16
5.121
5.123
Aqueous solutions of HCl and NaOH, HNO3 and KOH, HCl and KOH will give the desired net ionic equation.
No reaction occurs.
K+
5.125
I–
O
C
Solution of CO
Solution of KI
5.127
5.129
Any metal higher in activity can be used. (a) Al; (b) Zn; (c) Mg; (d) Sn
(a) Potassium chromate precipitates iron(III) chromate leaving Al3+ in solution. (b) Sodium sulfate precipitates barium sulfate leaving Mg2+ in solution. (c) Silver perchlorate precipitates silver chloride leaving NO3- in solution. (d) Barium sulfate is insoluble in
water; MgSO4 will dissolve in water.
Chapter 6 – Quantities in Chemical Reactions
6.1
6.3
6.5
6.7
6.9
6.11
(a) stoichiometry; (b) heat; (c) endothermic reaction; (d) specific heat; (e) actual yield
(a) C6H12O6 + 6O2
6CO2 + 6H2O; (b) 72 CO2 molecules; (c) 5 C6H12O6 molecules
The coefficients give the relative ratios of products used and reactants produced. They can represent the actual numbers of molecules or moles of each substance.
(a) 1 Mg2+(aq) and 2 NO3−(aq) ions (b) 1 mole of Mg2+(aq) and 2 moles of NO3−(aq) ions
The best representation is (d); however, (c) could also be chosen because it has the reactants and products in the correct proportions. In reaction (a), the reactants are not diatomic. In reaction (b), mass is not conserved (not balanced).
15mol O2
(a)
(b) 7.5 mol O2; (c) 2.9 mol C6H6
2 mol C6 H 6
6 mol HNO3
3mol H 2
2 mol Al
; (b)
; (c)
3mol H 2
2 mol Al
2 mol Al
6.13
(a)
6.15
(a) not conserved; (b) conserved; (c) conserved; (d) conserved; (e) moles of molecules may be conserved in some cases. The others
are conserved in all cases.
6.17
(a)
6.19
6.21
6.23
6.25
6.27
6.29
6.31
6.33
(b) 3.50 mol C, 3.50 mol CO, 1.50 mol N2, 2.50 mol H2O (c) 21.9 mol C, 21.9 mol CO, 9.38 mol N2, 15.6 mol H2O
(a) 249.70 g/mol; (b) 159.62 g/mol; (c) 0.639 g CuSO4
(a) 2.05 g; (b) 5.89 g; (c) 7.94 g
(a) 21.0 g CO; (b) 38.0 g I2
(a) 1.23 g Cl2; (b) 0.0460 g H2O; (c) 2.13 g O2; (d) 3.06 g Cl2
1.0 g CaCO3
the fuel
(a) 12 sandwiches; (b) bread is limiting; (c) 3 pieces of turkey
(a) F2; (b) N2; (c) N2 and F2 (both react completely)
6.35
(a)
3mol N 2
5mol H 2 O
7 mol C
7 mol CO
,
,
,
2 mol C7 H 5 (NO2 )3 2 mol C7 H 5 (NO2 )3 2 mol C7 H 5 (NO2 )3 2 mol C7 H 5 (NO2 )3
(b) oxygen
(c) hydrogen
6.37
2C2H2(g) +
7O2(g)
Initially mixed
6 molecules
18 molecules
How much reacts
4 molecules
14 molecules
Composition of
mixture
2 molecules
4 molecules
4CO2(g) + 6H2O(g)
0 molecules
0 molecules
8 molecules
12 molecules
Appendix D Answers to Selected Questions and Problems
6.39
6.41
6.43
6.45
6.47
6.49
6.51
6.53
6.55
6.57
6.59
6.61
6.63
6.65
6.67
6.69
6.71
6.73
6.75
6.77
6.79
6.81
6.83
6.85
6.87
6.89
6.91
6.93
6.95
6.97
6.99
6.101
6.103
6.105
6.107
6.109
6.111
6.113
6.115
6.117
6.119
6.121
2C2H10(g) + 13O2(g)
Initially mixed
3.10 mol
13.0 mol
How much reacts
2.00 mol
13.0 mol
Composition of
mixture
1.10 mol
0.0 mol
8CO2(g)
+
A-17
10H2O(g)
0.00 mol
0.00 mol
8.00 mol
10.0 mol
(a) P4; (b) both are consumed completely; (c) O2
(a) F2; (b) 12.5 g NF3
limiting reactant, NaHCO3; 4.2 g CO2 produced
2C2H6(g) + 7O2(g)
Initially mixed
0.260 g
1.00 g
How much reacts
0.260 g
0.968 g
Composition of
mixture
0.000 g
0.03 g
4CO2(g) + 6H2O(g)
0.00 g
0.00 g
0.761 g
0.467 g
(a) NaOH(aq) + HNO3(aq)
H2O(l) + NaNO3(aq) (b) HNO3 (c) basic
actual yield
An actual yield greater than the theoretical yield can be caused by contamination. The solid may have not been dry, causing the
apparent mass of product to be high.
(a) 7.44 g; (b) 8.95 g; (c) 83.1%
90.5%
0.47 mol I2
84%
The potential energy is converted to kinetic energy (including heat), some of which is transferred to the ground as heat.
The potential energy of the hydrogen and oxygen is used to provide electrical energy which runs electric motors and produces
kinetic energy.
The reaction is endothermic because the decrease in temperature of the surroundings indicates energy is absorbed by the reaction.
As the liquid evaporates, the molecules going into the gas state must absorb some energy from their surrounding. As a result, the
surroundings (water and your skin) lose energy and feel colder.
The products are lower in potential energy.
2.20 × 103 J
3.47 × 104 cal
0.209 Cal
114 Cal
lead
5470 J or 1310 cal
−210 J
A calorimeter should have good insulation and a way to accurately and precisely measure the temperature.
(a) The rock must have lost heat. (b) qrock = −1.3 × 103 J
(a) lower; (b) gain; (c) 2.8 kJ
63.5°C
Because a reaction is a process, you can only measure the effect it has on the surroundings.
The chemical reaction is the system, and the water and calorimeter are the surroundings.
(a) exothermic; (b) -525 kJ/mol
(a)147 kJ; (b) -24.5 kJ/g
(a) qnut = −2.6 × 104 J so 2.6 × 104 J is released; (b) 6.3 × 103 cal or 6.3 Cal; (c) 3.2 Cal/g
−256 kJ
(a) 3.00 mol CO2, 1.50 mol N2, 0.250 mol O2, 2.50 mol H2O (b) 7.50 mol CO2, 3.75 mol N2, 0.625 mol O2, 6.25 mol H2O
(a) 30.6 AgNO3; (b) 25.8 g
9330 g O2
(a) 2K(s) + Cl2(g)
2KCl(s) (b) Cl2 (c) There should be some gray solid and a white solid (KCl) inside the container.
(d) 11 g (e) 4.5 g K
1mol O2
; (b) 7.9 g O2 and 8.9 g H2O; (c) 4.5 g H2O
2 mol H 2
31.8°C
0.450 J/(g °C); chromium
2.80 × 103 kJ/mol, 6.70 × 102 kcal/mol
(a)
A-18
6.123
6.125
Appendix D Answers to Selected Questions and Problems
54.8°C
0.450 J/(g °C); chromium
7.41
Their primary differences are size and energy. The 3p
orbital is larger and higher in energy.
Chapter 7 – Electron Structure of the Atom
7.1
7.3
(a) electromagnetic radiation; (b) frequency; (c) ionization
energy; (d) Hund’s rule; (e) electron configuration; (f) core
electron; (g) orbital; (h) continuous spectrum;
(i) isoelectronic
Infrared, microwave, and radio frequency
7.5
7.43
7.45
(a) 1; (b) 3; (c) 5; (d) 7; (e) 1; (f) 3
Si
B
P
Lower frequency (
) with higher
frequency (
) superimposed
7.7
7.9
7.11
7.13
7.15
7.17
7.19
7.21
7.23
7.25
7.27
7.29
7.31
7.33
7.35
7.37
7.39
blue, yellow, orange, red
As wavelength increases, the frequency decreases. As
wavelengths decrease, frequencies increase.
infrared
Gamma photons have the highest energy and highest frequency. Radio frequency waves are the longest.
4.00 × 1015 Hz, ultraviolet
4.27 × 10−19 J, visible light (blue)
white light
no
No. If they did, line spectra would not exist. The energy of
the electron is quantized—it can only have certain values.
To move to a higher orbit, the electron would have to
absorb energy. To go to a lower orbit, it would have to
release energy (e.g., a photon).
A single photon is released.
The n = 6 to n = 3 transition gives the highest-energy
photon.
The n = 5 to n = 3 transition gives the lowest-energy
photon and therefore the longest wavelength.
The four lines are a result of four different transitions in
the hydrogen atom. These transitions are
n = 6 to n = 2
violet (highest energy)
n = 5 to n = 2
blue
n = 4 to n = 2
green
n = 3 to n = 2
red (lowest energy)
4.58 × 10-19 J
Bohr’s orbits required that the orbits be a fixed distance
from the nucleus with the electron following a specific
pathway. Modern orbitals describe the region of space
surrounding the nucleus where we are most likely to find
the electron.
B
3p
2p
7.47
7.49
7.51
7.53
7.55
7.57
7.59
7.61
7.63
7.65
7.67
7.69
7.71
7.73
7.75
7.77
7.79
7.81
7.83
7.85
1s
2s
2p
3s
3p
1s
2s
2p
3s
3p
1s
2s
2p
3s
3p
(a) 1s22s22p63s23p2; (b) 1s22s1; (c) 1s22s22p63s2
All orbitals can hold 2 electrons (including a 2p orbital).
8 electrons
3
elements in group VIIA (17)
transition metals
silicon
(a) 1s22s22p63s1
(b) 1s22s22p63s23p64s23d5
(c) 1s22s22p63s23p64s23d104p4
Bromine should be 1s22s22p63s23p64s23d104p5. The 4p
orbital given in the problem has one too many electrons
and the 10 electrons in the d orbital are in the third principle energy level.
(a) chlorine; (b) cobalt; (c) cesium
(a) [Ne]3s1; (b) [Ar]4s23d5; (c) [Ar]4s23d104p4
bromine (1s22s22p63s23p64s23d104p5)
Valence electrons are the electrons in the highest principle
energy level.
No, the d-orbital electrons are always one energy level
lower than the valence electrons.
As you go left to right on the periodic table, one electron
is added each time the group number increases. The group
number is the number of electrons in the s and p orbitals of
the highest energy level. Electrons in the d orbitals are not
included since they are always one energy level below the
highest energy level.
(a) third valence level; 3 valence electrons
(b) third valence level; 6 valence electrons
(c) fourth valence level; 5 valence electrons
Cations always have fewer electrons than the element from
which they are formed. They are similar in that they possess the same core of electrons.
(a) 1s22s22p6; [Ne]; F −
(b) 1s22s22p6; [Ne]; N3−
(c) 1s22s22p63s23p63d10; [Ar]3d10; Zn2+
O2−, N3−, Na+
Valence electrons farther from the attraction of the
nucleus are easier to remove. This means that potassium (4s valence electrons) has a lower ionization energy
Appendix D Answers to Selected Questions and Problems
7.87
7.89
7.91
7.93
7.95
7.97
7.99
7.101
7.103
7.105
7.107
7.109
7.111
7.113
7.115
7.117
7.119
7.121
than sodium (3s valence electrons). Since less energy is
expended to remove potassium’s electron, more energy is
left to drive the reaction.
The first ionization energy of calcium is higher than that of
potassium. Since less energy is expended to remove potassium’s electron, more energy is left to drive the reaction.
In addition, calcium must lose a second electron before it
becomes stable. Even more energy is expended to remove
this electron.
P<S<O
Sodium’s valence electron is in a higher principle energy
level. This means it is not as close to the nucleus and is
therefore easier to remove.
The fluorine has more protons in its nucleus than oxygen.
As a result, the electrons of fluorine are more strongly
attracted to the nucleus; the ionization energy is higher.
First ionization, IE1, Mg(g)
Mg+(g) + 1e−
+
Second ionization, IE2, Mg (g)
Mg2+(g) + 1e−
Magnesium because the third electron is removed from the
core electrons.
(a) Fluorine has a very high ionization energy. (b) Elements do not lose core electrons when forming ions (under
normal conditions). (c) high ionization energy because Ne
has a core configuration
O<S<P
The valence electrons of chorine and sulfur are both in
principal energy level 3, but chlorine has more protons in
its nucleus to attract the electrons.
(a) Mg; (b) P3−
Both ions are isoelectronic with argon, but potassium ion
is larger because it has fewer protons in its nucleus.
For hydrogen the energies of the sublevels are all the same.
(a) [Xe]6s24f 145d106p3; (b) [Rn]; (c) [Rn]7s2
12 filled p orbitals
This means that there are 10 filled d orbitals.
Both are related to the attraction of electrons for the
nucleus. As attraction gets higher, electrons are more
tightly held. This results in greater ionization energy and
smaller radius.
The radius of the sodium ion is smaller and the chloride
ion greater than for their respective elements.
Neon’s red color indicates lower energy, lower frequency,
and longer wavelength. Argon’s blue color indicates higher
energy, higher frequency, and shorter wavelength.
8.3
8.5
8.7
8.9
8.11
8.13
8.15
8.17
8.19
8.21
(a) single bond; (b) alkane; (c) covalent bonding; (d) ionic
crystal; (e) octet rule; (f) polar covalent bond; (g) alkyne;
(h) electronegativity; (i) triple bond; (j) crystal lattice;
(k) Lewis formula; (l) chemical bond
A chemical bond is an attractive force between atoms or
ions in a substance.
Metals form ionic bonds with nonmetals.
HF, NCl3, CF4
(a) ionic; (b) covalent; (c) covalent; (d) ionic
A and D
CH4, SF4, HCl
(a) high; (b) high; (c) low; (d) high; (3) low
The electronegativity value decreases.
bonds between two identical atoms
(a) N = Br < Cl < O < F (highest electronegativity);
(b) H < C < N < O < F
-
+
+
-
8.23
(a) δ H-F δ ; (b) δ I-Cl δ ; (c) δ B-Hδ
8.25
8.27
(a) H−H < C−H < O−H < F−H
(b) O−Cl < C−Cl < F−Cl < HCl
The Lewis symbol can be used to determine the number
of valence electrons and the number of electrons gained or
lost to reach an octet configuration.
8.29
(a) C ; (b) I ; (c) Se ;
(d) Sr ; (e) Cs ; (f) Ar
8.31
(a) Cl – ; (b) Sc3+; (c) S
8.33
(a) Li+
(c) Ba2+
8.35
8.37
8.39
8.41
8.43
8.45
8.47
8.49
8.51
8.53
2–
; (d) Ba2+; (e) B3+
–
Cl
–
(b) Cl
Chapter 8 – Chemical Bonding
8.1
-
+
A-19
Ba2+
S
Cl
–
2–
The second electron removed from potassium, K, would be
taken from the core. These electrons are very strongly held
by the atom and are not easily removed.
When fluorine gains an electron and sodium loses one
electron, both have satisfied the octet rule. The ionic bond
formed during the process is very stable and the compound
is electrically neutral.
A crystal lattice is the repeating pattern observed in all
crystal structures. An ionic crystal is the structure that
forms when ions form crystals by minimizing the repulsive
energies of like-charged ions and maximizing the contact
of oppositely charged ions.
Six chloride ions surround each sodium ion.
The ratio of the cations and anions is different, so different
ionic crystals are formed.
F F (a) 8 valence electrons around each atom
O O
in O2 and F2; (b) The bond in O2 is a double bond. The
bond in F2 is a single bond.
Each hydrogen needs only one electron to satisfy its
valence shell. It does not have another electron to form
another bond.
(a) 1; (b) 3; (c) 1; (d) none
(a) boron
(b) oxygen
(c) nitrogen
(a) H C N
H
(b)
H C
C N
(c) H C
C H
H
(d) H
H
(e)
C
C
H
H
H H
H C
C H
H H
Appendix D Answers to Selected Questions and Problems
A-20
8.55
(a) O
–
O
N
8.75
8.77
8.79
O
O
S
S
(d) O
N O
(e)
8.57
8.59
8.61
8.63
O
O
N
8.83
2–
O
O
H C C C C H
O
O
(c)
8.81
2–
O
(b)
8.85
–
+
8.87
(a) different; (b) same; (c) same; (d) same; (e) same
Resonance descriptions are required when a molecule or
ion can be represented by two or more reasonable Lewis
structures that differ only in the positions of the bonding
and lone pair electrons. The positions of the atomic nuclei
do not change.
Resonance is exhibited in (c) and (d).
(a) resonance structures for NO2−
N O
–
N O
O
S
O
O
O
S
O
O
C O
C O
O
O
O
H O N O
8.65
8.67
8.69
8.71
8.73
O
S
O
O C O
2–
(e)
8.91
O
2–
O C O
O
2–
C O
O
O
H
C
C
C
C
H
C
C
H
H
8.95
8.97
8.99
8.101
8.103
8.105
8.107
H O N O
Hydrogen’s valence shell can only hold two electrons.
(a) octet obeyed; (b) expanded octet; (c) octet obeyed;
(d) octet obeyed
(a) S, expanded octet; (b) B, incomplete octet; (c) Xe,
expanded octet; (d) Cl, odd electron species
alkanes with only single bonds, alkenes with double
bonds, and alkynes with triple bonds
H
H
(d)
8.93
(c)
O C O
(d)
O
(c)
8.89
O
H H H
The geometric arrangement of electron pairs and atoms is
predicted by finding the geometry that gives the greatest
distance between the electron groups and atoms (Table
8.5). The shape is determined by the arrangement of the
bonded atoms (Table 8.6).
Shape is determined by first knowing the number of
bonded atoms and unshared electron pairs (Table 8.5).
Once that is determined, the shape is given by the arrangement of atoms (Table 8.6).
(a)
(b)
–
O
(b) O
(a) alcohol; (b) alkane; (c) ether; (d) alkene; (e) ketone
(a) alcohol; (b) alkyne; (c) aldehyde; (d) ether
(a) alkane; (b) amine; (c) ester
H H H O
8.109
8.111
8.113
(a) linear; (b) tetrahedral; (c) tetrahedral; (d) trigonal planar; (e) tetrahedral; (f) tetrahedral; (g) trigonal planar;
(h) tetrahedral
(a) linear, 180°; (b) trigonal pyramidal, 109.5°; (c) bent,
109.5° (tetrahedral); (d) bent, 120° (trigonal planar);
(e) bent, 109.5° (tetrahedral); (f) tetrahedral, 109.5°;
(g) trigonal planar, 120°; (h) bent, 109.5° (tetrahedral)
(a) 109.5°; (b) 109.5°; (c) not defined; (d) 180°; (e) 120°;
(f) 120°; (g) 109.5°
(a) BF3; (b) NH3; (c) SCl2
(a) BCl3; (b) H2O; (c) BeCl2; (d) NH4+
ClO3−
image B
Each nitrogen has one unshared electron pair.
The bond angles on the carbon are 109.5°. The bond
angles on the nitrogen are 120°.
A bond is polar if there is a difference in electronegativity between the two atoms that are bonded. A molecule is
polar if the polarity of the bonds to a central atom do not
cancel out.
The chlorines are all attracting the electrons of the carbon
with equal force. Because the forces are in opposite direction and the molecule is geometrically symmetric, the
polarity of the bonds cancels in the molecule.
(a) polar; (b) polar; (c) polar; (d) polar
(a) SO2 has an unshared pair of electrons on the central
atom and is not symmetrical, so the bond polarities do not
cancel. As a result SO2 is polar. CO2 is linear and contains
polar bonds but is symmetrical. The bond polarities in
CO2 cancel, so it is a nonpolar molecule. (b) SO2 has an
unshared pair of electrons on the central atom and is not
symmetrical, so the bond polarities do not cancel. As a
result SO2 is polar. SO3 is trigonal planar and contains
polar bonds but is symmetrical. The bond polaritites in
Appendix D
8.115
8.117
8.119
SO3 cancel, so it is a nonpolar molecule. (c) SeCl2 has
two unshared pairs of electrons on the central atom and is
not symmetrical, so the bond polarities do not cancel. As
a result SeCl2 is polar. BeCl2 is linear and contains polar
bonds but is symmetrical. The bond polarities in BeCl2
cancel, so it is a nonpolar molecule. (d) Both are tetrahedral, but in the case of CH3I, the bond polarities do not
cancel and the molecule is polar. The bonds in CH4 are all
equivalent, so the bond polarities cancel.
image B
The molecule CCl2F2 should be soluble in water.
image B
8.121
(a) Br (b) Pb (c) S
8.135
(d) Ca (e) Be (f) Xe
8.123
8.125
8.127
C
9.1
S
9.3
(c) F As F
F
(d) O C O
(e) C O
8.129
(a)
O
(a) Both structures are tetrahedral. The CCl4 structure is
nonpolar due to symmetry. However the different types
of atoms (with different electronegativities) bonded to the
carbon in CH2Cl2 causes this molecule to be polar.
(b) Both structures are tetrahedral and polar but the CH3F
molecule is more polar than the CH3Br molecule because
the C−F bond is more polar than the C−Br bond.
(c) Both structures are trigonal pyramidal and polar but
the NF3 molecule is more polar than the NH3 molecule
because the N−F bonds are more polar than the N−H
bonds.
(d) Both structures are bent and polar but the H2O molecule is more polar than the OF2 molecule because the
O−H bonds are more polar than the O−F bonds.
Chapter 9 – The Gaseous State
F > Cl > Br > I
(a) covalent; (b) ionic; (c) covalent; (d) covalent; (e) ionic
(a) H N N H
(b) S
9.5
–
(a) effusion; (b) Boyle’s law; (c) combined gas law;
(d) ideal gas; (e) pressure; (f) Dalton’s law of partial pressures; (g) molar volume; (h) ideal gas constant, R
Y axis units and label are missing, x axis units are missing, the points should not be connected by lines (a straight
best-fit line that comes closest to all points should be
used), gridlines should be uniformly labeled (for example,
10, 20, 30, 40, 50, etc.), the x axis maximum should be 60
instead of 120, a graph title is missing.
There is a trend, but it is not linear.
Boiling Point vs. Number of Carbon Atoms
O Cl O
175
O
125
N O
O
(c)
O
–
–
N C O
O
(e)
F
8.133
–
N C O
(d) H C O
8.131
–
F
B
Boiling point (°C)
–
–
H C O
N C O
75
25
–25
–75
–125
–175
–
0
1
2
O
9.7
F
(a) tetrahedral; (b) trigonal planar; (c) bent; (d) trigonal
pyramidal; (e) tetrahedral; (f) bent; (g) tetrahedral;
(h) trigonal planar; (i) bent; (j) bent
(a) BeCl2 is nonpolar because the molecule is symmetrical.
OCl2 is polar. The molecule has polar bonds and is bent
(not symmetrical).
(b) PH3 is polar. The bonds are polar and the molecule is
not symmetrical. BH3 is nonpolar because the molecule is
symmetrical.
(c) BCl3 is nonpolar because the molecule is symmetrical.
AsCl3 is polar because the bonds are polar and the molecule is not symmetrical.
(d) SiH4 is nonpolar because the molecule is symmetrical.
NH3 is polar because the bonds are polar and the molecule
is not symmetrical.
3 4 5 6 7 8 9 10 11
Number of carbon atoms
Pressure and volume are inversely related.
Inches Hg vs. Inches Gas
Inches Hg
(b) N O
A-21
Answers to Selected Questions and Problems
100
90
80
70
60
50
40
30
20
10
0
0
10
20
30
40
Inches gas
50
60
A-22
9.9
9.11
9.13
9.15
9.17
9.19
Appendix D Answers to Selected Questions and Problems
The vapor pressure is 21 torr and 50 torr at 22°C and 38°C,
respectively.
(a) x = 0.4; (b) x = 0.206; (c) x = ±4; (d) x = 2
(a) 100°C; (b) 26.7°C; (c); 0°C; −40.0°C
(a) 24.4; (b) 0.116; (c) 17.2; (d) 4.70
When compared to other states of matter, gases have low
densities and are very compressible. They also take the
shape of any container they are put in.
The density of warm air is lower than the density of cold
air.
9.47
9.49
9.51
9.53
9.55
9.21
9.23
9.25
Gas pressure is the amount of force exerted by the gas
divided by the area over which the force is exerted.
Absolute pressure is measured with a device called a
barometer. A tire gauge measures the pressure above atmospheric pressure.
9.27
9.57
9.59
9.61
9.63
9.65
9.67
9.69
9.71
9.73
9.29
9.31
9.33
9.35
(a) 0.980 atm; (b) 935 torr; (c) 0.119 atm; (d) 643 Pa;
(e) 1.01 × 103 mm Hg; (f) 563 torr; (g) 1.06 × 105 Pa;
(h) 293 mm Hg
1.024 × 105 Pa
As pressure increases, the volume decreases at constant
temperature.
The particles collide more frequently with the container
walls causing an increase in pressure. Since the number of
moles has not changed, we expect to find twice as many
molecules in the same space. The velocity of the molecules
should not change.
9.75
9.77
9.79
9.81
9.83
9.85
9.87
9.89
9.91
9.93
9.95
9.97
9.99
9.101
9.103
9.37
9.39
9.41
9.43
9.45
(a) 3.60 L; (b) 26.7 mL; (c) 0.392 mL
(a) 161 torr; (b) 0.100 torr; (c) 17.9 atm
182 atm
0.577 L
Charles’s law states that if the pressure is constant, the
volume and temperature (in kelvin) are directly proportional to each other (volume increases when temperature
increases; volume decreases when temperature decreases).
The velocity of the gas molecules increases as temperature
increases. The volume of the container will increase as the
temperature increases.
(a) 5.41 L; (b) 671 mL; (c) 75.5 L
(a) 273°C; (b) -265°C; (c) −18°C
0.169 mL
Nothing happens to the particles if the temperature, pressure, and volume are constant. If the tank is sealed so that
no gas molecules are allowed to escape, then nothing happens to the pressure. If the tank is opened, then the molecules of gas in the tank will leave the tank. Particles leave
the tank to maintain an equilibrium pressure with the air
outside the tank.
(a) 418 torr; (b) 673 torr; (c) 4.44 atm
(a) −91.3°C; (b) 750°C; (c) 437°C
7.71 atm
(a) 1.2 L; (b) −195.3°C; (c) 9.10 × 102 torr
1.61 L
Guy-Lussac’s law states that volumes of gas react in
simple whole-number ratios when the volumes of reactants
and products are measured at the same temperature and
pressure.
22.414 L/mol
(a) 0.385 mol, 6.17 g; (b) 1.562 × 10-2 mol, 2.050 g;
(c) 2.15 mol, 60.1 g
The number of particles is the same. The balloon containing argon has the greater mass, and therefore greater density since the volumes are the same.
(a) 0.34 mol, 7.6 L; (b) 1.5 mol, 34 L; (c) 2.70 mol, 60.5 L
2380 balloons
11.0 L
Any gas whose behavior is described by the five postulates
of kinetic molecular theory is an ideal gas.
(a) 0.34 mol, 0.69 L; (b) 1.5 mol, 3.1 L; (c) 2.70 mol,
5.51 L
(a) 0.245 mol, 3.92 g; (b) 4.33 × 10−3 mol, 8.74 × 10−3 g;
(c) 0.288 mol, 8.06 g
The balloon filled with CO2 sinks because the density of
CO2 is greater than the density of air.
(a) 0.7600 g/L; (b) 1.250 g/L; (c) 1.964 g/L
(a) 0.673 g/L; (b) 1.11 g/L; (c) 1.74 g/L
There are 0.208 moles of each gas. (a) 0.42 g; (b) 3.3 g;
(c) 13 g
Dalton’s law of partial pressures states that the total pressure in a container is the sum of the pressures exerted by
each of the individual gases in a mixture of gases.
708 torr
(a) 2.78 mol; (b) 2.08 mol; (c) 2.15 atm; (d) 1.60 atm
(a) 40.01 g/mol; (b) 30.03 g/mol; (c) 45.99 g/mol;
(d) 20.3 g/mol; (d) 16.0 g/mol
The five postulates of kinetic molecular theory are
1) Gases are composed of small particles widely separated.
2) Particles behave independently of each other.
3) Gas particles move in straight lines.
4) Gas pressure results from force exerted by the molecules in the container. The force is the sum of the forces
exerted by each molecule as it bounces off the walls of the
container.
5) The average kinetic energy of the gas particles depends
only on the absolute temperature.
Appendix D Answers to Selected Questions and Problems
9.105
9.107
9.109
9.111
9.113
9.115
9.117
9.119
9.121
9.123
9.125
9.127
As the temperature is increased, the kinetic energy of the particles increases. Recall that kinetic energy is the energy of motion.
Increasing temperature increases molecular velocity. The particles strike the walls of the container with greater force and, therefore,
the pressure is higher.
Pressure depends on the force of collisions and the frequency of collisions in a given area. The density of gas particles is inversely
proportional to the volume. If the volume decreases, the density increases. Since the density of gas particles increases with decreasing volume, the frequency of collisions in a given area has increased. This means that the pressure has increased.
(lowest velocity) CO2 < CH4 < He < H2 (highest velocity)
Molecules with lower masses have higher average velocities and as a result diffuse faster.
helium
24 L H2
48.0 L CO2 and 76.0 L O2
1.68 × 103 g
The balloon expands or contracts so that the external pressure and the internal pressure are the same.
The density of a gas decreases if it is heated and the volume is allowed to expand. The air in the hot-air balloon is less dense than
the surrounding air, so the balloon floats.
1100 K
What will happen if . . .
Macroscopic View
Microscopic View
Temperature increases and
pressure remains constant?
Balloon gets bigger
Molecules move faster
Molecules collide harder
Collision frequency increases
Molecules are further apart
Temperature decreases and
pressure remains constant?
Balloon gets smaller
Molecules move slower
Less energetic collisions
Collision frequency decreases
Molecules are closer together
Balloon probably
gets a little bigger
Molecules move slower
Less energetic collisions
Lower collision frequency
10,000 feet (temperature
decreases and pressure
decreases)
9.129
9.131
9.133
9.135
9.137
9.139
9.141
9.143
A-23
1.75 L
1.21 L
27.7 L
(a) 5.00 L; (b) 3.33 L; (c) 1.67 L
228 K
(a) If the pressure and temperature are constant, increasing the moles of gas increases the volume occupied by the gas. (b) pressure
and temperature
131 g/mol
0.718 L
Chapter 10 – The Liquid and Solid States
10.1
(a) vapor pressure; (b) melting point; (c) equilibrium; (d) evaporation; (e) induced dipole; (f) alloy; (g) boiling point;
(h) sublimation; (i) molecular solid; (j) intermolecular force; (k) London dispersion force; (l) melting point; (m) crystal
10.3
Particle spacing
Intermolecular attraction
Kinetic energy
10.5
10.7
10.9
10.11
Liquid
Gas
Dense or closely spaced
Molecules far apart
Moderate
Weak
Low
High
A liquid is fluid because the molecules can move past each other. Attractive forces hold the particles in solids together in a rigid,
three-dimensional array.
The substance is in the solid state. If the substance were a liquid or gas, it would not maintain its shape. In addition, if the substance
was not a solid, the images on the surface would not maintain their shape.
The image is a molecular-level representation of the liquid state. The liquid state and solid state both have dense particle spacing.
Basically, the molecules are always touching each other. However, in the image, the molecules do not show any clear pattern. This
is a characteristic of the liquid state.
The solid, liquid, and gas states are represented in circles 1, 3, and 2, respectively. The phase changes are deposition, vaporization,
and melting for phase changes A, B, and C, respectively.
A-24
10.13
10.15
Appendix D Answers to Selected Questions and Problems
The vapor pressure of a liquid increases as the temperature
is increased. A liquid boils when its vapor pressure is equal
to the pressure above the liquid. The temperature at which
this occurs is the boiling point. The normal boiling point is
the temperature at which the vapor pressure of the liquid is
equal to 1 atm.
In order to convert a substance from the liquid to the gas
state, the energy of the molecules must increase enough to
allow them to overcome their attractions for each other and
move about independently. Because the molecules must
absorb energy to go into the gas state, the process is endothermic (i.e., absorbs energy from the surroundings).
10.57
10.59
10.61
10.17
The vapor pressures of pure liquids depend on the strength
of the intermolecular forces between the particles of the
liquids.
When a molecule is polar, it means that some portion of
the molecule appears permanently slightly negative to
other molecules and another portion is slightly positive.
This comes from the unequal sharing of electrons (Chapter
8). The oppositely signed charges on different molecules
attract each other (dipole-dipole interaction), so the molecules stick together. The instantaneous and induced dipoles
in nonpolar molecules are not permanent.
As molar mass increases the importance of London dispersion forces increases.
10.63
10.19
10.21
10.23
10.25
10.27
10.29
10.31
10.33
10.35
10.37
10.39
10.41
10.43
10.45
10.47
10.49
10.51
10.53
10.55
solid state to the gas state, sublimation; endothermic
All molecules and atoms have some degree of attraction
for each other. As a gas is cooled, the molecules are losing
kinetic energy. Eventually the attractive forces are greater
than the kinetic energy of the molecules and the molecules
begin to coalesce (stick together) to form the liquid state.
Evaporation of water is endothermic. The energy water
needs to evaporate comes from the surrounding air, cooling it. The cooled air can be used to cool a house (if it is
circulated by a fan).
Water boils when it has been heated to the point that its
vapor pressure equals the atmospheric pressure. At high elevation, the atmospheric pressure can be significantly lower
than 1 atm (about 0.6 to 0.7 atm at 12,000 ft). Since the
water boils at a lower temperature, the pasta cooks slower.
As a liquid is heated the molecules of the liquid acquire
more energy. The percent of molecules on the liquid surface with enough energy to go into the gas phase increases.
This causes an increase in the vapor pressure.
Gallium is a liquid at 100°C. At 15°C gallium is a solid.
BC
237 kJ
41.1 kJ
1.69 × 103 kJ
124 kJ
An attractive force between individual particles (molecules
or atoms) of a substance is called an intermolecular force.
gas state
Ar
NO, NF3, and CH3Cl
A bond is an attractive force between two atoms in a molecule. A hydrogen bond is not a true bond because it usually takes place between two atoms on different molecules.
In comparison to covalent bonds, the strength of attraction
is much weaker and the electrons are not shared.
Only B can hydrogen-bond.
(a) and (b)
(b)
Dispersion
Force
(a) C6H6
x
(b) NH3
x
(c) CS2
x
(d) CHCl3
x
10.65
10.67
10.69
10.71
10.73
10.75
10.77
10.79
10.81
10.83
10.85
Dispersion
Force
(a) Kr
x
(b) CO
x
(c) CH4
x
(d) NH3
x
DipoleDipole
x
Hydrogen
Bonding
x
x
DipoleDipole
Hydrogen
Bonding
x
x
x
(lowest boiling point) CH3OH < CH3CH2OH <
CH3CH2CH2OH (highest boiling point)
NO2
Hydrogen, H2, is larger than (but not more massive than)
helium. The size of the hydrogen molecule allows the
electrons to be moved around more easily and therefore
the dispersion force is greater. More energy is required to
vaporize liquid hydrogen than helium.
HCl is polar, while F2 is nonpolar.
(a) CH3OH; (b) CH3OH; (c) NH3
(lowest) He < N2 < H2O < I2 (highest)
Both molecules are polar, but water has very strong hydrogen bonding while acetone has dipole-dipole interactions.
As a result, water has a higher boiling point.
On a molecular level, the particles of a substance in the
liquid state are higher in energy than for the same substance in the solid state and the molecules have freedom
to move apart from each other. In addition, the liquid state
does not have a long range structure like the crystal lattice
of many solids. On a macroscopic level, liquids flow and
can take the shapes of their containers. Liquids tend to be
less dense than solids.
Viscosity is the resistance of a liquid to flow. The more
viscous a liquid, the slower it will flow when it is poured.
All molecules of a liquid experience attractive forces to
other molecules in the liquid. Molecules in the middle
experience forces that are equal in all directions. On the
surface, the molecules only experience attractive forces
Appendix D Answers to Selected Questions and Problems
10.87
10.89
10.91
10.93
10.95
10.97
10.99
10.101
10.103
10.105
10.107
10.109
10.111
10.113
10.115
10.117
10.119
10.121
10.123
10.125
10.127
10.129
to the sides and “downward” (toward the center of the liquid). Surface tension causes liquids to bead up on solid surfaces and is
also why liquids tend to form spheres. In addition, surface tension is related to capillary action and the formation of a meniscus
(curved surface) of a liquid in a container.
There is an imbalance of forces on the molecules at the surface of a liquid. As the surface area of the liquid is minimized, the
imbalance of energy is also minimized. Substances assume the shape with the least surface area for the given volume. As a result,
liquids have a tendency to form spherical shapes to minimize this energy difference.
When ice forms on bodies of water such as lakes and rivers, it stays at the surface and provides an insulating layer. Energy loss
through evaporation is minimized and less water freezes than would if the ice did not form at the top. Ice would continually form
and sink. Many bodies of water would freeze completely making life in them impossible during the coldest months.
Crystals are defined as solids that, on a molecular or atomic level, have a regular repeating pattern of atoms. See Figures 10.31 and
10.37.
Crystal structures help us understand the macroscopic properties of the solid state (i.e., conductivity, hardness, brittleness).
Close packing is a common structure of metals.
Amorphous solids have a tendency to flow because the molecules are not locked into a crystal lattice.
Solids are classified by the strength of attraction between the subunits of the solid. This allows us to broadly classify the characteristics of the solid substances.
ionic solid
There are two reasons. First, the attractive forces between metal atoms are relatively weak. This allows them to be separated from
each other. Secondly, the electrons in metal atoms are delocalized among all their neighboring atoms. Atoms can be moved with
relatively little effect on the charge distribution. In ionic solids, cations and anions cannot be separated easily, so particles do not
move apart easily. This causes ionic solids to be brittle.
Both ionic and metallic solids form crystal lattices. Ionic crystals resemble metal crystals with smaller ions fitting into holes
between larger ions.
Molecular solids are held together either by dipole-dipole interactions, hydrogen bonding, or London dispersion forces. Ionic solids
are held together by interionic attractions (i.e., forces of attraction between oppositely charged ions). Because ionic interactions are
much stronger than intermolecular forces, the ionic solids tend to melt at much higher temperatures.
(a) covalent bonds; (b) London dispersion forces; (c) London dispersion forces, dipole-dipole interactions, and hydrogen bonding;
(d) covalent bonds
(a) ionic; (b) molecular; (c) metallic; (d) network; (e) molecular
(b) and (e) are molecular solids
molecular solid
SiF4
Below its freezing point, ice is lost by sublimation.
CH4 (−162°C), C2H6 (−88°C), C3H8 (−42°C), C4H10 (0°C); London dispersion forces increase with molecular size.
A glass plate (SiO2) is very polar. Polar substances are attracted to polar substances. Water is attracted to glass, so it spreads out,
but mercury is repelled by the glass and tends to bead up.
(c) is an ionic solid
(a) Cl2; (b) CH3SH; (c) NF3
Substance
(solid type)
Electrical
Conductivity
Hardness
Melting Point
Vapor Pressure
NaCl (ionic)
conductive in
molten state
hard
high
low
SiC
(network)
not conductive
hardest
highest
lowest
SiCl4
(covalent)
not conductive
softest
lowest
highest
high
low
Fe (metallic)
always conductive soft*
*Metals are soft relative to network or ionic solids.
10.131 melting; freezing; sublimation; deposition; vaporization; condensation
10.133
A-25
Appendix D Answers to Selected Questions and Problems
A-26
Chapter 11 – Solutions
11.1
11.3
11.5
11.7
11.9
11.11
11.13
(a) saturated solution; (b) aqueous solution; (c) molarity;
(d) solvent; (e) entropy; (f) solubility; (g) parts per million;
(h) miscible; (i) osmosis; (j) colligative property;
(k) percent by volume; (l) mass/volume percent
Solute
Solvent
(a) sodium chloride (and other salts)
water
(b) carbon
iron
(c) oxygen (O2)
water
Strong electrolytes are substances that fully ionize when
placed in solution. Strong acids, strong bases, and soluble
ionic compounds are all strong electrolytes.
nonelectrolyte
(a) H+(aq) and Br− (aq); (b) NH4+(aq) and Cl−(aq);
(c) CH3CH2CH2CH2OH(aq)
A
HO
(a) Ca(OH)2(s) 2
Ca2+(aq) + 2OH−(aq)
(b) N2(g)
H2O
N2(aq)
HO
11.15
11.17
11.19
11.21
11.23
11.25
11.27
11.29
11.31
11.33
11.35
11.37
11.39
CH3OH(aq)
(c) CH3OH(l) 2
In the solvent, hydrogen-bonding and dipole-dipole interactions are broken. In the solute, ionic bonds are broken.
Ion-dipole interactions are formed in the solution.
In both the solute and the solvent, London dispersion
forces must be overcome. The new forces are also London
dispersion forces.
The attraction between the products (dissolved ions) and
water molecules is weaker than the forces of attraction in
the pure substances.
Energy and entropy drive the formation of solutions.
The intermolecular attractions between water molecules
and oil molecules are not the same (water is polar and
cooking oil is nonpolar). Since they are not “like,” they are
immiscible.
Both molecules are polar and can engage in hydrogen
bonding.
(a) insoluble; (b) soluble; (c) soluble
Ion-ion interactions are very strong. There is not enough
energy released or entropy gained in the formation of
ion-solvent interactions to make the dissolution process
favorable.
The ions in NaCl are attracted by ionic bonding. The
most significant attractive force between water molecules
is hydrogen bonding. Although disrupting the attractive
forces in this solute and solvent requires considerable
energy, formation of the ion-dipole interactions between
solute and solvent particles releases enough energy to
allow the ionic compound to dissolve.
The solubility of gases such as O2 decreases as temperature increases.
The solubility of oxygen and nitrogen gases should be
lower at high altitude. The solubility of these gases is
directly proportional to their pressures above the solution
and these pressures are lower at high altitude.
The insoluble gases that cause the bends can be redissolved into the blood by putting the divers in pressurized
chambers. This increases the solubility of the gas in the
blood. Decreasing the chamber pressure slowly (decompression) allows the gases (mostly excess nitrogen) to
leave through the lungs.
11.41
11.43
11.45
11.47
11.49
11.51
11.53
11.55
11.57
11.59
11.61
11.63
11.65
11.67
11.69
11.71
11.73
11.75
11.77
11.79
11.81
11.83
11.85
11.87
11.89
11.91
11.93
11.95
17.0 g KCl
On a macroscopic level, you can’t dissolve more solid in a
saturated solution, but you can dissolve more in an unsaturated solution. On a molecular level, the ions in a saturated
solution are in equilibrium with the solid.
You could add a little solid at a time until no more solid
dissolves and you see traces of undissolved solid in the
solution.
The solution will be saturated and 0.85 g Ca(OH)2(s) will
be left undissolved.
14.3%
105 g KI
5.423%
Percent by volume is often used for mixtures of liquids.
This is done because liquids are usually measured by
volume.
30.4%
(a) 2.2 ppb; (b) Since the maximum contamination level is
5 ppb, the water is safe to drink.
0.015 mol
2.41 m
3.0 M; 3.0 m
9.6 × 10−4%; 9.6 ppm; 9600 ppb
(a) 9.0 ppb; yes, the concentration is below the EPA
standard; (b) 9.0 × 10−6 mg/mL; (c) 9.0 × 10−4 mg;
(d) 4.3 × 10−9 mol
(a) 0.030 L KI; (b) 6.9 g PbI2
0.30 mol Na2CO3
57.14 mL
0.8812 M H2SO4
(a) 10.00 mL; (b) 52.50 mL; (c) 37.5 mL
A colligative property is a physical property of a solvent
that varies with the number of solute particles dissolved
and not on the identity of those solute particles.
Osmosis occurs when two solutions of different solute
particle concentrations are separated by a membrane that
allows the solvent but not the solute molecules to pass
(semipermeable membrane). The volume of the more
concentrated solution increases because solvent molecules
migrate through the membrane to the more concentrated
solution.
Reverse osmosis occurs when solvent is forced to travel in
a direction opposite of the direction it would travel during
osmosis.
Water will flow out of the blood cell, and the cell will collapse or shrink.
(a) decrease; (b) increase; (c) decrease
The boiling point of the water is elevated slightly because
of the presence of the sugar.
(a) 1.3°C; (b) 101.3°C
The sodium chloride solution will have the higher boiling
point because NaCl produces two solute particles when it
dissolves in water.
Appendix D Answers to Selected Questions and Problems
11.99
11.101
11.103
11.105
11.107
11.109
11.111
11.113
11.115
11.117
11.119
Solution B has the highest concentration of dissolved
particles and therefore has the lowest freezing point. Glucose is molecular, and images B and C show molecular
compounds. Sodium chloride produces two ions of slightly
different size as shown in image A.
(a) 2.0 mol; (b) 5.0 mol; (c) 4.0 mol
The osmotic pressure of MgCl2 will be greater than that
of NaNO3 because each formula unit of MgCl2 produces
three ions when it dissolves while NaNO3 produces only
two ions.
Vitamin D is nonpolar.
The breaking of the crystal lattice and disruption of the
intermolecular attractions between water molecules require
energy. The formation of ion-dipole interactions between
water and the ions releases energy.
182 g
8.90 M
Volume changes with temperature. If the concentration
units are based only on mass or number of moles, then
they will not change with temperature. Percent by mass (a)
and molality (d) do not change with temperature.
The density of a dilute solution is approximately the density of the solvent. In aqueous solutions, for example, one
liter of water has a mass of one kilogram. So the values for
molarity (moles/liter) and molality (moles/kilogram) will
be very similar.
(a) 10.5 g BaSO4; (b) 2.33 g BaSO4; (c) 2.33 g BaSO4
2.0 × 103 g
solution B
Chapter 12 – Reaction Rates and Chemical
Equilibria
12.1
12.3
12.5
12.7
12.9
12.11
12.13
12.15
12.17
(a) activation energy, Ea; (b) intermediate;
(c) homogeneous equilibrium; (d) Le Chatelier’s principle;
(e) collision theory; (f) equilibrium constant expression
increasing reactant concentrations, increasing temperature,
or providing a catalyst
grind the calcium carbonate to a fine powder; increase the
temperature; increase the concentration of HCl
catalyst
No. If the reactants do not collide with sufficient energy
or with the proper orientation, then the activated complex
cannot be formed.
(a) improper orientation; (b) proper orientation;
(c) improper orientation
activated complex
minimum energy required to break bonds in reactants
before reactants can be converted into products
For similar reactions, the higher the activation energy, the
slower the reaction.
12.19
‡
Ea
Energy
11.97
A-27
EReactants
EProducts
Reaction progress
12.21
12.23
12.25
12.27
12.29
12.31
12.33
12.35
12.37
12.39
12.41
12.43
12.45
12.47
12.49
12.51
12.53
12.55
Increasing temperature increases the kinetic energy and
velocity of the molecules reacting. The fraction of colliding molecules with energy greater than the activation
energy increases when temperature is increased.
increasing temperature
At room temperature, molecules of propane and oxygen
do not collide with sufficient energy to cause a reaction.
However, by providing a spark (heat), the kinetic energy of
the molecules is increased. Combustion reactions are exothermic, so the reaction is sustained by the heat produced
by the reaction.
Roasts are generally thicker and more massive than steaks.
For the same amount of heat, the steak gets hotter faster
and therefore cooks faster.
A catalyst provides a lower-energy route (pathway) to the
same products.
Catalytic converters use palladium and platinum catalysts
(Figure 12.11).
active site
Catalysts are only temporarily changed during the course
of a reaction (e.g., during the binding of the substrate).
Chlorine atoms in the upper atmosphere act as catalysts for
the decomposition of ozone. Since they act catalytically,
they are not lost until they diffuse into outer space (which
is a slow process) or are removed by some other reaction.
intermediate
(a) NO3 is an intermediate; (b) N2O5 + NO
3NO2
(a) H+ and Br− are catalysts and Br2 is an intermediate.
(b) 2H2O2
2H2O + O2
When a reaction reaches equilibrium, the concentrations of
reactants and products stop changing.
A chemical equilibrium is a state of a reaction where
the rate of the forward reaction is equal to the rate of the
reverse reaction.
The system has reached equilibrium in images D and E.
(a) yes; (b) no; (c) yes; (d) yes
You could measure the mass of the bromine liquid, but that
would be tricky because it is very reactive and volatile.
A better option is to measure the pressure of the gas produced. When the partial pressure has stabilized, the system
is at equilibrium. Finally, bromine gas has a dark brown
appearance. By measuring how much light the gas absorbs,
we can determine if its concentration remains constant and
the system is at equilibrium.
The position of the equilibrium tells us whether a reaction
is product-favored or reactant-favored.
A-28
12.57
Appendix D Answers to Selected Questions and Problems
12.59
The coefficients of the reactants and products are used as
exponents in the equilibrium constant expression.
The brackets represent molar concentrations.
12.61
(a) Keq =
12.93
2
HF
H2 F2
In the following table, the larger arrows (⇑) indicate the
change made to the equilibrium reaction and the smaller
arrows indicate what the concentration of the products
and reactants do following the change. Finally, the central
arrow indicates which direction the equilibrium shifts.
NO(g)
(b) Keq =
(c) Keq =
12.63
12.65
12.67
12.69
12.71
12.73
12.75
12.77
12.79
12.81
12.83
12.85
CS2 H2
4
CH4 H2S
NO2
N2O4
2
2
(a) C(g)
A(g) + B(g)
(b) 2A(g)
4B(g) + C(g) + D(g)
(a) The equations are inverses of each other.
(b) 2NO2
2NO + O2; 2NO + O2
2NO2
0.25
An equilibrium constant only changes with temperature.
(a) 3.1 × 10−3; (b) Because the value of Keq is much less
than one, the position of equilibrium is described as
“reactant-favored.”
(a) 1.3 × 10−9; (b) The reaction is reactant-favored since
Keq < 1.
The reaction is at equilibrium, so shifting will not take
place.
The reaction will proceed in the reverse direction.
The reaction will proceed in the forward direction.
Heterogeneous reactions are those in which the products
and reactants are not all in the same physical state.
The concentrations of pure liquids and solids do not
change during chemical reactions; only the amount of
the pure substance changes. Because the concentrations
of pure substances do not change during reactions, their
values are essentially incorporated into the equilibrium
constant.
(a) Keq = NH3 HCl
12.95
12.97
12.99
12.101
12.103
12.105
12.107
12.109
(b) Keq = Ca2+ CO32–
+
–
(c) Keq = H3O F
HF
12.87
12.89
12.91
0.27 M
Le Chatelier’s principle describes how reactions that are at
equilibrium respond to changes in reaction conditions. If
stress is put on a reaction, the reaction responds to counteract that stress.
(a) The reaction shifts to the left. (b) The concentrations of
the products, CO2 and H2, will decrease and the reactant,
H2O, will increase in concentration.
12.111
12.113
SO3(g)
NO2
SO2
(a)
⇑
↓
→
↑
↑
(b)
↓
⇑
→
↑
↑
(c)
↓
↓
→
⇓
↑
(d)
↑
↑
←
⇑
↓
(a) This has no effect since the concentration of the solid
does not change when more solid is added. (b) This
decreases the concentration of Pb2+ by shifting the equilibrium to the left. (c) This causes an increase in Pb2+
concentration. When the concentration of lead ion is
increased, the equilibrium shifts to the left. However, it
never shifts enough to completely remove all the added
material.
(a) left; (b) right; (c) no effect
(a) no; (b) yes
(a) left; (b) right
If the equilibrium shifts to the right, the concentration
of products at equilibrium is higher. The value of Keq
increases.
(a) decrease; (b) increase
exothermic
An increase in the concentration of NO requires that the
equilibrium shift to the right. (a) Water is in the liquid
state; removing it does not affect the equilibrium.
(b) There are fewer moles of gas on the right. The reaction
shifts right increasing the concentration of NO. (c) Since
the reaction is exothermic, the reaction will shift right with
decreasing temperature producing a higher concentration
of NO. (d) Adding oxygen causes the equilibrium to shift
right producing a higher concentration of NO. (e) Catalysts
do not have an effect on equilibrium concentrations.
The reaction that produces light slows down, and the reactants take longer to be consumed. The lower light output
of the cold light sticks is a result of fewer photons being
produced per second (lower reaction rate).
It provides a different reaction mechanism with lower activation energy. In general it splits one difficult step (high
activation energy) into two or more steps that have lower
activation energies.
12.115 (a) Keq =
H2 I2
HI
2
(b) Keq = 0.0199
(c) The position of the equilibrium favors the reactants.
12.117 The effect of temperature depends on whether the reaction
is endothermic or exothermic.
12.119 The reaction must be exothermic since the reaction shifted
to the right to form the powder.
A-29
Appendix D Answers to Selected Questions and Problems
12.121 The simplest treatment would be to increase the concentration of oxygen that the patient is breathing. By increasing
the O2 concentration, the second equilibrium will shift to
the left and the carbon monoxide can be displaced and
exhaled through the blood.
Chapter 13 – Acids and Bases
13.1
13.3
13.5
13.7
13.9
13.11
13.13
13.15
13.17
13.19
13.21
13.23
13.25
13.27
13.29
13.31
(a) strong base; (b) Brønsted-Lowry theory; (c) conjugate
acid; (d) polyprotic acid; (e) acidic solution; (f) weak base;
(g) amphoteric substance; (h) ion-product constant of
water, Kw; (i) self-ionization
(a) −9; (b) −11.0; (c) 3.87; (d) 5; (e) 0.0
(a) 16; (b) 6.3 × 10−7; (c) 1
Acids have a sour taste (but you should never taste anything to see if it’s an acid!), are corrosive to many metals,
turn blue litmus red, and neutralize bases.
Common “household” bases and their uses:
Formula
Name
Use
Mg(OH)2
magnesium
hydroxide
antacid
Al(OH)3
aluminum
hydroxide
antacid
NH3
ammonia
cleaning solutions,
smelling salts
Ca(OH)2
calcium hydroxide
concrete
An Arrhenius acid produces hydrogen ions, H+(aq), when
dissolved in water. An Arrhenius base produces hydroxide
ions, OH−(aq), when dissolved in water.
Brønsted-Lowry theory refers to the transfer of hydrogen
ions from an acid to a base. Arrhenius similarly describes
the release of hydrogen ions into the solution, but identifies OH− as a component of the base. In Brønsted-Lowry
theory, a base is any substance that can accept a hydrogen
ion (including OH−).
(a) acid; (b) base; (c) acid
(a) NO2−; (b) F −; (c) H2BO3−
(a) H2O; (b) C6H5NH3+; (c) H2CO3
(a) HCl(g) + H2O(l)
H3O+(aq) + Cl−(aq)
(b) HClO4(l) + H2O(l)
H3O+(aq) + ClO4−(aq)
−
(c) CH3CO2 (aq) + H2O(l)
HCH3CO2(aq) + OH−(aq)
Both water (a) and hydrogen sulfite (b) are amphoteric.
Carbonic acid has two acidic hydrogen atoms.
Strong acids and bases ionize or dissociate completely
(100%). For weak acids and bases, the percent of molecules that ionize is usually much lower (typically a few
percent).
(a) strong acid; H2SO4(aq) + H2O(l)
HSO4-(aq) + H3O+(aq)
(b) strong base; Ca(OH)2(aq)
Ca2+(aq) + 2OH−(aq)
(c) weak base; Na2CO3(aq) + H2O(l)
HCO3−(aq) + 2Na+(aq) + OH−(aq)
(d) weak acid; H3C6H5O7(aq) + H2O(l)
H3O+(aq) + H2C5O7−(aq)
(e) weak base; C6H5NH2 + H2O(l)
C6H5NH3+(aq) + OH−(aq)
A is NH3, B is HF, and C is HCl.
KF; HF is a weak acid, so its conjugate base, F−, is a weak
base.
13.35 Acid A is a stronger acid. The stronger the acid, the greater
the extent of ionization.
13.37 0.50 M HCO2H
13.39 NaCH3CO2
13.41 In a strong acid solution the acid molecules are fully ionized. It does not necessarily mean that the concentration is
high. A concentrated acid solution has a high percentage
of acid, but this does not necessarily mean that the acid is
ionized to any great extent.
13.43 A polyprotic acid is an acid that contains more than one
acidic hydrogen atom.
13.45 Hydrosulfuric acid is a weak acid that has two acidic
hydrogen atoms. The corresponding acid-base reactions
are.
H2S(aq) + H2O(l)
HS−(aq) + H3O+(aq) Ka1 = 8.9 × 10−8
−
HS (aq) + H2O(l)
S2−(aq) + H3O+(aq)
Ka2 = 1.0 × 10−19
13.47 The following substances can all be found in a solution of
oxalic acid: H2C2O4, HC2O4−, C2O42−, H3O+, H2O.
The substance with the lowest concentration is the oxalate
ion C2O42−.
13.49 Like other equilibrium constants, it is constant only under
conditions of constant temperature.
13.51 In pure water, [OH−] = [H3O+] = 1.0 × 10-7 M.
13.53 (a) basic; (b) acidic; (c) basic
13.55 (a) 1.0 × 10−11 M, basic; (b) 1.0 × 10−3 M, acidic;
(c) 3.1 × 10−7 M, acidic
13.57 (a) 0.010 M; (b) 0.020 M; (c) 6.7 × 10−13 M
13.59 A decrease in pH represents an increase in the acidity.
13.61 A 10-fold increase in hydrogen ion concentration is 1 pH
unit change. The difference in pH units is 1.
13.63 Practically speaking, most solutions of acids and bases are
lower in concentration than 1 M so the pH scale typically
ranges between 0 and 14. However, the pH values can
exceed this range.
13.65 No. A solution is defined as acidic if the pH is less than 7.
13.67 (a) 3.00, acidic; (b) 13.00, basic; (c) 9.47, basic
13.69 (a) 10.00, basic; (b) 7.00, neutral; (c) 4.91, acidic
13.71 (a) 2.00, acidic; (b) 1.70, acidic; (c) 12.18, basic
13.73 Underlined values were given in the problem.
13.33
[H3O+]
[OH-]
pH
pOH
Acidic or
Basic?
(a)
1.0 × 10−5
1.0 × 10−9
5.00
9.00
acidic
(b)
1.0 × 10−10
1.0 × 10−4
10.00
4.00
basic
(c)
1.0 × 10−6
1.0 × 10−8
6.00
8.00
acidic
(d)
2.9 × 10−9
3.5 × 10−6
8.54
5.46
basic
(e)
1.1 × 10−5
9.0 × 10−10
4.95
9.05
acidic
13.75
13.77
13.79
pH = 2.30, pOH =11.70. The pH and pOH always add to
14.00 at 25°C.
(a) 1.0 × 10−5 M, acidic; (b) 1.0 × 10−12 M, basic;
(c) 1.3 × 10−6 M, acidic
(a) 1.0 × 10−11 M, basic; (b) 4.0 × 10−8 M, basic;
(c) 1.3 × 10−2 M, acidic
A-30
13.81
13.83
13.85
13.87
13.89
13.91
13.93
13.95
13.97
13.99
13.101
13.103
13.105
13.107
13.109
13.111
13.113
13.115
13.117
Appendix D Answers to Selected Questions and Problems
(a) 1.0 × 10−13 M; (b) 3.2 × 10-4 M; (c) 4.0 × 10-11 M
Acetic acid is a weak acid, so the concentration of hydronium ion would be less than 0.010 M, and the pH would be
higher than 2.0.
0.1 M
A pH meter typically gives measurements to two decimal
places (Figure 13.17).
Acid-base indicators are typically colored compounds that
are also weak acids or bases. The colors of these compounds depend on whether they are in their acid or base
forms.
The color range of an indicator indicates the pH values
over which the indicator changes color. The hue of the
indicator can be used to estimate the pH of a solution. The
range over which an indicator changes color is about 1.5
pH units.
A selection of indicators would include thymolphthalein.
It should be red with a small amount of yellow-orange.
Red (or the presence of red color) indicates the pH is less
than 2.5. Yellow indicates a pH between about 2.5 and 8.0.
Blue (or a blue hue) indicates pH values above 8.0. It will
be completely blue above pH 9.5.
Universal indicator is a mixture of different indicators. The
color of an unknown solution with the indicator present
gives a good estimate of the pH. It is similar to pH paper
in that pH paper contains universal indicator.
The solution would be yellow at the start of the titration
and turn blue at the endpoint.
A buffer contains both a weak acid and its conjugate base
(or a weak base and its conjugate acid). As a result, when
small amounts of acid or base are added, the buffer can
absorb the excess hydronium or hydroxide with only a
small change in pH.
The buffer system of the blood helps to maintain a pH
between 7.35 and 7.45.
any salt containing CH3CO2− (e.g., NaCH3CO2)
(a) Hydrogen phosphate, HPO42−, is the acid and the conjugate base is phosphate, PO43−.
HPO42−(aq) + H2O(l)
PO43−(aq) + H3O+(aq)
(b) When acid is added, hydronium, H3O+, is produced
and the pH momentarily drops (more acidic). However,
the excess hydronium quickly reacts with the PO43−
shifting the equilibrium to the left to produce the weak
acid HPO42−.
Solutions described in (b) and (c) produce buffered solutions.
As the concentration of carbon dioxide increases, the equilibrium of the reaction would shift to the right. The formation of carbonic acid would cause a decrease in the pH of
the extracellular fluid.
Ammonia, NH3, has a lone electron pair in its Lewis
structure. Methane does not have a lone electron pair,
so it cannot accept a hydrogen ion (proton) in an acidbase reaction.
When NaF dissolves in water, sodium and fluoride ions are
produced. The fluoride ions react with water to produce a
weak base HF by the reaction:
F −(aq) + H2O(l)
HF(aq) + OH−(aq)
13.119
13.121
13.123
13.125
In general the conjugate base (F−) of a weak acid will
react with water to produce hydroxide ions. The molecular
species HF is stable, so its formation drives the formation
of hydroxide. This solution will be basic. The sodium ion
does not react with water.
No. A neutral solution has a pH = pOH = 7.00 (at 25°C).
If sodium hydroxide is added, the hydroxide ion concentration will have to increase. The solution will have a pH
greater than 7.00.
Both H2SO4 and HNO3 are strong acids. Both will dissociate completely for the first ionization. However, H2SO4 is
a polyprotic acid. The additional hydronium ion produced
by the second ionization will make the pH of the sulfuric
acid lower.
(a) HOCl(aq) + H2O(l)
OCl−(aq) + H3O+(aq).
(b) The HCl reacts with the water to produce hydronium
ion: HCl(aq) + H2O(l)
Cl−(aq) + H3O+(aq). The
+
added H3O causes the equilibrium to shift to the left and
some of the excess hydronium is consumed. The HOCl
concentration will increase and the OCl− concentration
will go down. While the pH may not drop below 7.00, any
time an acid is added to a solution, the solution becomes
more acidic (pH is lowered).
Sulfuric acid is a strong acid and is completely ionized.
H3O+
SO42–
HSO4–
13.127 sodium formate and formic acid
Chapter 14 – Oxidation-Reduction Reactions
14.1
14.3
14.5
14.7
14.9
(a) electrolytic cell; (b) cathode; (c) oxidation; (d) halfreaction; (e) electrochemistry; (f) electrode; (g) oxidizing
agent; (h) oxidation-reduction reaction; (i) corrosion;
(j) oxidation number
reaction where electrons are transferred
when there is an increase in charge (oxidation number)
(a) Mg(s) + Cu(NO3)2(aq)
Cu(s) + Mg(NO3)2(aq)
(b) The charge of magnesium increases by two; 2 electrons
transferred; electrons are lost.
(c) The charge of copper decreases by two; 2 electrons
transferred; electrons are gained.
(a) Mg; (b) Sn2+; (c) Sn2+; (d) Mg
(e)
Mg
Mg2+
SO42–
2+
Sn Sn
14.11
An oxidation number is a bookkeeping method for tracking where electrons go during a chemical reaction.
Appendix D Answers to Selected Questions and Problems
14.13
14.15
14.17
14.19
14.21
14.23
14.25
14.27
14.29
14.31
14.33
14.35
14.37
(a) 0; (b) 0; (c) 3+; (d) 1−
S = 6+, O = 2−
(a) 5+; (b) 3+; (c) 4+
(a) 5+; (b) 5+; (c) 5+; (d) 1+; (e) 3−; (f) 3+
S = 6+, O = 2−
(a) 7+; (b) 5+; (c) 1−; (d) 3+; (e) 5+
(a) Cl = 4+, O = 2−; (b) Ca = 2+, F = 1−; (c) H = 1+,
Te = 4+, O = 2−; (d) Na = 1+, H = 1−
(a) N = 3+, O = 2−; (b) Cr = 6+, O = 2−; (c) Ag =
1+, Cl = 1−; (d) S = 4+, O = 2−; (e) C = 4+, O = 2−
(a) not redox; (b) redox, N2 is the oxidizing agent, H2 is
the reducing agent; (c) not redox; (d) not redox; (e) redox,
O2 is the oxidizing agent, C2H6 is the reducing agent
(a) N2(g) + O2(g)
2NO(g). (b) The oxidation number increases. Each nitrogen atom loses two electrons
to change from an oxidation number of 0 to 2+. (c) The
oxidation number decreases. Each oxygen atom gains two
electrons to change from an oxidation number of 0 to 2−.
(a) V2+; (b) Cr2O72−; (c) Cr2O72−; (d) V2+
(a) I2 is both the oxidizing agent and reducing agent,
1 electron transferred.
(b) Cr = reducing agent, H+ = oxidizing agent, 2 electrons transferred.
(c) Cr2O72− is both the oxidizing and reducing agent, 12
electrons transferred.
(d) Fe3+ = oxidizing agent, Al = reducing agent, 3 electrons transferred.
In the anode half-cell the half-reaction is
Fe(s)
Fe2+(aq) + 2e−
In the cathode half-cell the half-reaction is
Ni2+(aq) + 2e−
Ni(s)
Voltmeter
Anode
Fe
Salt
bridge
14.43
14.45
14.47
14.49
14.51
14.53
14.55
14.57
Ni
Cathode
14.59
14.61
Fe(NO3)2(aq)
14.39
Ni(NO3)2(aq)
The iron electrode will lose mass as Fe is oxidized and
becomes Fe2+(aq). The nickel electrode will gain mass as
Ni2+ is reduced to Ni(s).
14.63
14.65
14.67
Iron electrode
14.41
Nickel electrode
The Pb(s) gets oxidized and the PbO2(s) is reduced. The
oxidizing agent is PbO2(s) and the reducing agent is Pb(s).
A-31
oxidation half-reaction: Cd(s) + 2OH−(aq)
Cd(OH)2(s) + 2e−
reduction half-reaction: 2e− + NiO2(s) + 2H2O(l)
Ni(OH)2(s) + 2OH−(aq)
−
3+
(a) 3e + Fe (aq)
Fe(s)
(b) Zn(s)
Zn2+(aq) + 2e−
(c) 2Cl−(aq)
Cl2(g) + 2e−
(d) Fe2+(aq)
Fe3+(aq) + e−
(a) Zn(s) + 2NO3−(aq)
Zn(NO3)2(aq) + 2e−
−
3e + Fe(NO3)3(aq)
Fe(s) + 3NO3−(aq)
3Zn(s) + 2Fe(NO3)3(aq)
3Zn(NO3)2(aq) + 2Fe(s)
(b) Mn(s) + 2Cl−(aq)
MnCl2(aq) + 2e−
2e− + 2HCl(aq)
H2(g) + 2Cl−(aq)
Mn(s) + 2HCl(aq)
MnCl2(aq) + H2(g)
Fe2(SO4)3(aq) + 2KI(aq)
2FeSO4(aq) + K2SO4(aq) + I2(aq)
(a) Ba(s)
Ba2+(aq) + 2e−
(b) e− + H+(aq) + HNO2(aq)
NO(g) + H2O(l)
(c) 2e− + 2H+(aq) + H2O2(aq)
2H2O(l)
(d) 2Cr3+(aq) + 7H2O(l)
Cr2O72−(aq) + 14H+ + 6e−
(a) 3OH−(aq) + La(s)
La(OH)3(s) + 3e−
−
−
(b) 2e + H2O(l) + NO3 (aq)
NO2−(aq) + 2OH−(aq)
(c) H2O2(aq) + 2OH−(aq)
O2(g) + 2e− + 2H2O(l)
(d) 8e− + 3H2O(l) + Cl2O7(aq)
2ClO2−(aq) + 6OH−(aq)
(a) 8H+(aq) + 3H2S(aq) + Cr2O72−(aq)
3S(s) + 2Cr3+(aq) + 7H2O(l)
+
2+
(b) 4H (aq) + 5V (aq) + 3MnO4−(aq)
5VO2+(aq) + 3Mn2+(aq) + 2H2O(l)
(c) 6H+(aq) + 6Fe2+(aq) + ClO3−(aq)
6Fe3+(aq) + Cl−(aq) + 3H2O(l)
−
(a) 2NH3(aq) + ClO (aq)
N2H4(aq) + Cl−(aq) + H2O(l)
(b) 2Cr(OH)4−(aq) + 3HO2−(aq)
2CrO42−(aq) + 5H2O(l) + OH−(aq)
−
(c) 2OH (aq) + Br2(aq)
Br−(aq) + BrO−(aq) + H2O(l)
24H+(aq) + 5C6H12O6(aq) + 24NO3−(aq)
30CO2(g) + 12N2(g) + 42H2O(l)
The anode could be any metal that is more active than iron
(e.g., aluminum or zinc). The supporting electrolyte could
be any soluble salt that does not precipitate with Fe3+ (e.g.,
potassium nitrate since all nitrates are soluble). The anodic
half-cell would also contain a salt that includes the metal
used for the anode (e.g., aluminum or zinc nitrate).
Au < Bi < Ni < Zn < Al < Ca
Electrolysis is the process of using electricity to make a
nonspontaneous redox reaction occur.
(a) Reduction will begin to take place at the cathode to
form sodium metal.
(b) Oxidation will begin to take place at the anode to form
I2(g).
(c) Na+(l) + e−
Na(l) (reduction)
2I−(l) → I2(g) + 2e− (oxidation)
(d) 2Na+ + 2I−(l)
2Na(l) + I2(g)
Appendix D Answers to Selected Questions and Problems
A-32
14.69
Anode
Cathode
15.7
15.17
Of the elements with atomic number less than 84, technetium (Tc), promethium (Pm), and astatine (At) have
no stable isotopes. No isotopes of atomic number 84 or
greater are stable.
Dense materials are better at shielding gamma radiation.
Lead, iron, and aluminum (metals) can be used, but concrete and water are effective only if thick layers are used.
Nuclear decay reactions and bombardment reactions result
in the production of radiation. In addition, new isotopes
are produced by both.
During nuclear decay alpha particles, beta particles,
gamma rays, and positrons are produced by the nucleus.
During nuclear reactions, nuclear charge and mass numbers are conserved.
222
86 Rn
15.19
99
43 Tc
15.21
15.23
15.25
15
7N
beta particle
7
4 Be
15.27
(a) 238
92U
15.9
I2(g)
Na(l)
Anode reaction
14.71
14.73
14.75
14.77
14.79
14.81
14.83
14.85
14.87
14.89
Cathode reaction
Magnesium is more active than iron. The magnesium will
reduce the O2(g) and keep the Fe in the reduced state.
When the Mg(s) is used up, the iron will begin to corrode.
One of the methods of preventing corrosion is to prevent
contact with oxygen and moisture. The tin coating serves
as a barrier to moisture and oxygen. When the tin coating
is scratched, the more active metal, iron, is more readily
oxidized.
The chromium should corrode first because it is a more
active metal than iron.
(a) 3−; (b) 2−; (c) 3+; (d) 1−; (e) 5+; (f) 3+
(a) 2H2O(l) + 2OH−(aq) + 2CoCl2(s) + Na2O2(aq)
2Co(OH)3(s) + 4Cl−(aq) + 2Na+(aq)
The oxidizing agent is Na2O2 (oxygen is being reduced
from 1− to 2−). The reducing agent is CoCl2 (cobalt is
being oxidized from 2+ to 3+).
(b) 2OH−(aq) + Bi2O3(s) + 2ClO−(aq)
2BiO3−(aq) + 2Cl−(aq) + H2O(l)
The oxidizing agent is ClO− (chlorine is being reduced
from 1+ to 1−).
The reducing agent is Bi2O3 (bismuth is being oxidized
from 3+ to 5+).
NH4+(aq) + NO3−(aq)
N2O(g) + 2H2O(l)
The oxidizing agent is NO3− (N is being reduced from 5+
to 1+).
The reducing agent is NH4+ (N is being oxidized from 3−
to 1+).
(a) Al; (b) Zn; (c) Mn; (d) Mg
Brass is an alloy of copper and zinc. Cu doesn’t react with
acid. There must be a protein in sweat that facilitates the
oxidation of the copper to form Cu2+. There are several
compounds of copper that are green (including CuCO3).
2H+(aq) + Cu(s) + H2SO4(aq)
Cu2+(aq) + SO2(g) + 2H2O(l)
+
(a) reducing agent, Tl ; oxidizing agent, Ce4+;
oxidation half-reaction: Tl+(aq)
Tl3+(aq) + 2e4+
reduction half-reaction: 2Ce (aq) + 2e2Ce3+(aq)
(b) reducing agent, Br ; oxidizing agent, NO3
oxidation half-reaction: 2Br-(aq)
Br2(l) + 2ereduction half-reaction: 2NO3 (aq) + 4H+(aq) + 2e2NO2(g) + 2H2O(l)
Chapter 15 – Nuclear Chemistry
15.1
15.3
15.5
(a) critical mass; (b) fission; (c) alpha particle; (d) radioactive; (e) gamma radiation; (f) nuclear bombardment;
(g) nucleon; (h) radiation
A substance that contains unstable nuclides of an element
will exhibit radioactivity by the emission of radiation.
Radioactivity produces photons (usually gamma rays) and
particulate radiation (i.e., beta rays, positrons, neutrons,
and alpha particles).
15.11
15.13
15.15
234
90Th
(b) 234
90Th
(c) 234
92U
234
91Pa
4
2α
0 (d) 210
83Bi+ -1e
15.29
(a) 227
87 Fr
(b) 189 F
+ 42α
+ -10β-
+ 230
90 Th
210
82Pb
4
2α
+ 223
85 At
+ +01β+ + 188O
15.33
0
14
(c) 146 C
-1β + 7N
In a nuclear bombardment reaction a particle accelerator is
used to accelerate particles or nuclides to very high velocities. The beam of particles produced is aimed at a target
element. If the particle strikes the target with sufficient
energy to fuse the target and beam particles, a newly created
nucleus is formed. The newly created nucleus spontaneously
decays to produce new substances and decay particles.
257
103 Lr
15.35
(a) 105 B + 21 H
15.31
4
(b) 209
83 Bi + 2α
(c) 126 C + 11 p
15.37
1
(a) 96
42 Mo + 1 p
64
(b) 209
83 Bi + 28 Ni
15.39
15.41
15.43
15.45
15.47
15.49
15.51
11
6C
+ 10 n
211
85 At
+ 210n
13
7N
96
43 Tc
+ 10 n
272
111Rg
+ 10 n
approximately 1:1
If the mass numbers of all the stable isotopes are graphed
versus their atomic number, a pattern emerges (Figure
15.14). The band of stability is the outline of that pattern.
The isotope 168 O should be the most stable. It has a N/Z
ratio of 1.
(a) beta; (b) alpha and beta; (c) alpha and beta;
(d) positron; (e) beta
234
542 α + 2-10β- + 214
91Pa
83 Bi
In a sample of radioactive material, the half-life is the
time it takes for half of the radioactive nuclei to decay. For
nuclear decay, each half-life has exactly the same duration.
Film, Geiger-Müller counter, and a scintillation counter
can be used to detect radiation.
A-33
Appendix D Answers to Selected Questions and Problems
15.53
15.55
15.57
15.59
15.61
15.63
15.65
15.67
15.69
15.71
15.73
15.75
15.77
15.79
15.81
15.83
15.85
15.87
15.89
15.91
15.93
One-sixteenth (6.25%) of the original material will be
present.
three half-lives or 42 days
22.5 hours (three half-lives)
5730 years
Twenty-four hours represents four half-lives. The half-life
is 6 hours.
Radiation-powered devices can last much longer than
battery-powered devices.
99m
Tc is used because it produces radiation that can be
used to image tumors and its short half-life ensures that it
won’t stay in the body for a long period of time.
Positron emission occurs for N/Z ratios < 1. 116 C has a low
N/Z ratio, whereas 146 C has a high N/Z ratio.
Radiation can (1) have no effect, (2) cause reparable damage, (3) cause damage leading to formation of mutant cells
(cancer cells), (4) kill cells (subsequently absorbed by the
body). Mutant cells can also be recognized as foreign entities and be destroyed by an immune response.
This is a difficult question to answer. How you answer
the question stems from your view of our governmental
system and your perception of the risk of radon and the
benefit of regulations. Clearly, the best protection against
unnecessary radon exposure is knowledge of what causes
enhanced levels of radon and what measures can be used
to reduce radon exposure. The Internet can be a good
source of information if you use it carefully (i.e.,
www.epa.gov/iedweb00/radon/index.html).
You would most likely find 210
82 Pb.
Fission occurs when unstable nuclei decay into smaller
nuclei and decay products. Fusion results when smaller
particles combine to make larger (“heavier”) nuclei.
three
In principle, a chain reaction needs the terminating step of
the reaction to be a necessary component of the initiating
step. For example, if the products of a reaction (initiating
step) start a second reaction (terminating step) and the
products of the second reaction start the first reaction, you
have a chain reaction.
A fission reactor has fuel rods (containing the radioactive
material), control rods (to absorb the excess neutrons),
a moderator (i.e., water, heavy water, or graphite), and
energy converters (convert water to steam which powers
electric generators).
The water is not directly converted to steam by the nuclear
reaction as a safety precaution. The steam that would be
produced by the nuclear reaction would be highly radioactive and is also made of poisonous heavy water (water with
deuterium rather than hydrogen is highly toxic). If there
were a leak of the steam, the released products would be
very dangerous.
The waste contains a mixture of fission by-products as
well as materials contaminated by radioactive waste (such
as the spent fuel rods).
Isotopes of hydrogen (H-1, H-2, H-3) and lithium are most
likely to be used.
2
3
4
1
1H + 1H
2 He + 0 n
1
1H
3
+ 21 H
2 He
In radioactive dating, a sample is taken and analyzed for
the quantity of a particular radioisotope. Based on the
amount of the isotope, and assuming you know how much
of the isotope was initially present in the sample, you can
calculate the sample age (using the half-life). For example,
1
if the radioisotope is at —4 the original concentration, two
1
1
half-lives have passed (i.e., —2 × —2).
15.95
15.97
1
The sample contains about —
64th the carbon-14 in a modern
sample. This is about 6 half-lives or 34,000 years.
1
1
2
0 +
1H + 1H
1H + +1β
2
1H
+ 11H
3
2He
4
0 +
+ 32He
2He + +1β
15.99 The initial product of the bombardment reaction is car13
bon-13 (94 Be + 42He
6 C). Loss of a neutron would
produce carbon-12 (N/Z =1):
9
4
12
1
4 Be + 2He
6 C + 0 n.
15.101 With a half-life of 138.4 days, at the end of one year
approximately 3 half-lives have already passed. Only oneeighth of the material would be remaining. You would be
better off buying brushes every year (or as you needed
them).
15.103 Even though the element chlorine has only been known
to exist since 1774, chloride-containing compounds were
used well before the discovery of chlorine.
1
1H
Chapter 16 – Organic Chemistry
16.1
16.3
16.5
16.7
(a) isomer; (b) alcohol; (c) ester; (d) aldehyde; (e) cyclic
hydrocarbon; (f) saturated hydrocarbon; (g) alkyl group;
(h) cycloalkane; (i) aliphatic hydrocarbon; (j) amino acid;
(k) organic chemistry; (l) unsaturated hydrocarbon
(a) seven single bonds; (b) one double bond, four single
bonds ; (c) one triple bond, two single bonds
(a) amine; (b) ketone; (c) alcohol; (d) alkene; (e) carboxylic acid; (f) alkyne; (g) aldehyde; (h) ether; (i) ester
molecular formula: C20H28O
structural formula:
H
H
H HH H
H C C H H C H H H C H H
H
C HC
C
C
C
C
C
C
H C
C
C
O
C
H H
H
H
H
H C
C
C
C H
H
H H H
line structure:
O
16.9
(a) C7H14O; (b) C8H18
16.11
(a)
(b)
O
(c)
(d)
O
C
H
Appendix D Answers to Selected Questions and Problems
A-34
16.13
(e)
16.35
(a) (CH3)2CHCH2CH3; (b) (CH2OH)2CHOH;
(c) CH3OCH2CH3
line structures:
16.37
(a)
(b)
16.39
OH
HO
16.43
HO
(b) H
H
16.19
16.21
16.23
16.25
16.27
16.29
3-methyl-1-butene
H
(a)
H
16.17
16.45
O
(c)
16.15
16.41
(a) CH3CH(CH3)CH3; (b) CH3CH(CH3)C(CH2CH3)2CH2CH3
(c) CH3CH(CH3)CH2CH(CH3)CH2CH2CH2CH3
(d) CH3CH2C(CH2CH3)(CH3)CH2CH2CH3
heat
(a) CH3CH2CH3 + Cl2
CH3CHClCH3 + HCl
heat
or CH3CH2CH3 + Cl2
CH3CH2CH2Cl + HCl
heat
(b) (CH3)3CCH2CH2CH3 + 11O2
7CO2 + 8H2O
heat
(c) CH3CH3 + Br2
CH3CH2Br + HBr
An alkene is a hydrocarbon containing one or more carbon-carbon double bonds.
(a) 5-methyl-3-heptene; (b) 2,3-dimethyl-2-hexene;
(c) trans-3,4-dimethyl-3-hexene
only (c) because it can be made to be asymmetric along
the carbon-carbon double bond axis
(a)
C H
H
C
C
(b)
(c)
H
H
H
H
(c) Cl
Cl
C C
H
H
H
H
An alkane hydrocarbon has only carbon-carbon single
bonds, an alkene contains at least one carbon-carbon
double bond, and an alkyne has at least one carbon-carbon
triple bond.
(a) unsaturated; (b) unsaturated; (c) unsaturated
Aromatic hydrocarbons are cyclic hydrocarbons that
can be drawn with alternating double and single bonds
between the carbon atoms. When the bonds are distributed
in this way, the electrons are said to be delocalized.
The London dispersion forces must be stronger in butane
than in 2-methyl propane (isobutane). London dispersion
forces increase with the length of the molecule because the
molecules have more opportunities to interact. Because the
London dispersion forces are stronger in butane, it boils at
a higher temperature.
(a) CH3CH2CH3; (b) CH3(CH2)4CH3
(a) ethane; (b) n-octane or octane; (c) propane;
(d) n-nonane or nonane
There are five isomers of C6H14.
propyne
(d)
16.47
trans-2-butene
cis-3-hexene
(a) CH3CH=CHCH3 + HBr
(b) CH3CH2CH=CHCH3 + Br2
(c)
16.49
16.51
16.53
+
H2
Aromatic hydrocarbons are cyclic and have alternating
single and double bonds (delocalized electrons).
(a) 1,2-dibromobenzene; (b) 1-chloro-4-ethylbenzene
H H
(a)
H
H
C
C
C
C
H
H
H
C
C
C
C
C
H
H
H H H
1,2,4-trimethylbenzene
H
H
(b)
C C
H C
C Cl
C C
H
H
chlorobenzene
FeCl3
16.55
16.57
16.59
16.61
C6H6 + Cl2
C6H5Cl + HCl
Both (b) and (c) are alcohols.
There are four isomers.
(a) 2-methyl-1-propanol; (b) 2-butanol
16.63
(a)
OH
(b)
16.33
(a) 2-methylpentane; (b) 2-methylbutane;
(c) 2-methylbutane (same as previous structure)
(a) 3-methylpentane; (b) 2,2-dimethylpropane;
(c) 2-methylbutane
CH3CH2CHBrCHBrCH3
Pd
3-hexanol
16.31
CH3CHBrCH2CH3
OH
2,2-dimethyl-3-pentanol
Appendix D Answers to Selected Questions and Problems
16.65
16.67
16.69
16.71
16.73
A-35
The primary difference is that the intermolecular forces between ethanol (CH3CH2OH) molecules are hydrogen bonds, while those
between propane molecules are London dispersion forces. The stronger intermolecular forces between ethanol molecules means
that a higher temperature (energy) is needed to raise the vapor pressure of ethanol to 1 atm.
O
Ethers contain the functional group −O− attached to two alkyl groups
.
R
R
Only (a) CH3OCH3 is an ether.
(a) ethylpropyl ether; (b) dimethyl ether
O
(a)
(b)
O
ethyl methyl ether
(
)
ethyl phenyl ether
16.75
Water is a hydrogen-bonding solvent, so substances that are hydrogen-bonding, polar molecules will be more soluble in water.
Ethers (CH3OCH3) do not directly hydrogen-bond because the oxygen is attached to carbon atoms on either side. Alcohols are
more water soluble because the alcohol group −OH can hydrogen-bond.
O
16.77
16.79
16.81
An aldehyde has the functional group R−CHO
where R can be an organic group or a hydrogen atom.
C
R
H
Compounds (a), (c), and (d) are aldehydes.
Both compounds are polar because of the C=O bond, but ethanal (CH3CHO) has stronger London dispersion forces because it has
a larger molecular mass. As a result ethanal has a higher boiling point.
( )
(
O
C
)
H where R is an organic group.
16.83
A carboxylic acid contains the functional group R−COOH or R−CO2H
16.85
One route to the formation of esters is the condensation of a carboxylic acid and an alcohol. Methyl propionate (CH3CH2CO2CH3)
can be formed from methanol (CH3OH) and propionic acid (CH3CH2CO2H).
Saturated fatty acids are high-molar-mass carboxylic acids which do not have double or triple bonds in the hydrocarbon chain. As
a result, it is no longer possible to add any more hydrogen atoms to the hydrocarbon chain and they are said to be saturated with
hydrogen.
Acetic acid, CH3CO2H, should have a higher boiling point than propanal, CH3CH2CHO, because acetic acid can hydrogen-bond.
Aldehydes are polar but do not hydrogen-bond.
N
An amine is a derivative of ammonia, NH3, with one or more of the H atoms replaced by organic groups R
H .
H
NH
(b)
(c)
(a)
2
N
N
N
or
H
16.87
16.89
16.91
16.93
16.95
16.97
16.99
R
O
( )
Methylamine, CH3NH2, can hydrogen-bond, so it has relatively strong intermolecular forces resulting in a higher boiling point.
Ethane, CH3CH3, is nonpolar and cannot hydrogen-bond.
(CH3CH2)2NH + HCl(aq)
(CH3CH2)2NH2+(aq) + Cl−(aq)
A simple organization of hydrocarbons is given in the chart below.
Hydrocarbons
Straight or branched
Saturated
Alkanes
Cyclic
Unsaturated
Alkenes
Alkynes
Cycloalkanes
Cycloalkenes
Nonaromatic
Conjugated
Aromatic
A-36
Appendix D Answers to Selected Questions and Problems
16.101 Alkenes and alkynes differ primarily in three ways: carbon-carbon bond type, bond angles, and number of hydrogen atoms in the
structure.
Alkanes
Alkenes
Alkynes
Bond type
Single
Double bonds
Triple bonds
Bond angle
109.5°
120°
180°
#H = 2n + 2
#H = 2n
#H = 2n − 2
General formula
(straight or
branched)
#H = number of hydrogens; n = number of carbons. The formula assumes only one double or triple bond per structure.
16.103 Aldehydes and ketones are similar functional groups.
O
O
. The aldehyde functional group is
.
C
C
R
H
R
R
(a) 2,4-dimethylpentane; (b) 1,3-dimethylcyclopentane; (c) 2,2,3,3-tetramethylbutane
One possible test is to compare the solubility of each substance in an alcohol (such as ethanol). However, this test is not very specific. Addition of bromine, Br2, which is a reddish brown solution would be a better alternative. Bromine will react with an alkene
but not with an alcohol. If an alkene is present, the bromine solution will lose color as the bromine reacts with the alkene.
The trend in boiling points is directly related to the strength of the intermolecular forces in each substance. The weakest attractive forces are found in propane which is nonpolar. Dimethyl ether is polar, but ethanol is both polar and has hydrogen bonding.
Because hydrogen bonds are so strong, ethanol has the highest boiling point.
(a) CH3COCH3; (b) CH3CH2COCH3; (c) CH3CH2CHO; (d) CHOCHCHCH3; (e) CH3CHO
(a) CH3CH2CH2CO2CH2CH3; (b) CH3CH2CH2CH2CO2H; (c) CH3CO2CH3; (d) HCO2H; (e) CH3CH2CO2H
(a) dimethylamine; (b) ethyldimethylamine (the official IUPAC name is N,N-dimethyl-1-aminoethane, which is why we use the
common name); (c) aniline
(a) CH3C(CH3)2CH2CH(CH3)2; (b) CH3CHC(CH2CH3)2; (c) CH2CHCHCH2; (d) C4H8; (e) C6H5Cl
(a) alkane; (b) ether; (c) carboxylic acid; (d) aldehyde; (e) ester; (f) amine; (g) ketone
Only (c) 2-butene and (d) 2-pentene can exist as cis and trans isomers.
(c)
(d)
A ketone has the general formula
16.105
16.107
16.109
16.111
16.113
16.115
16.117
16.119
16.121
16.123 (a) isomer; (b) isomer; (c) same compound
Chapter 17 – Biochemistry
17.1
17.3
17.5
(a) transfer RNA; (b) secondary structure; (c) nucleotide; (d) nucleic acids; (e) lipid; (f) primary structure; (g) deoxyribonucleic
acid; (h) fatty acid; (i) active site; (j) carbohydrate; (k) peptide; (l) replication; (m) translation
regulation of cell process, control reaction rates, transport substances in and out of cells, fight disease
COOH
H2N C H
R
17.7
COOH
H2N C H
H
COOH
COOH
H2N C H
H3C
Gly
CH
H2N C H
CH3
Val
CH2
OH
Ser
17.9
Alanine has a methyl group and valine has an isopropyl group.
O H
O
17.11
Ala-Gly H2N CH C
N
CH3
CH
OH ; The methyl group is on the N-terminal amino acid.
H
O H
Gly-Ala H2N CH C
H
C
N
O
CH
C
OH ; The methyl group is on the C-terminal amino acid.
CH3
The peptide bond is boxed and the side chains are circled.
A-37
Appendix D Answers to Selected Questions and Problems
17.13
O H
H2N CH C
N
CH
CH3
17.15
O
H
O
H
O
H
O
H
O
H
O
H
O
C
N CH
C
N CH
C
N CH
C
N CH
C
N CH
C
N CH
C
CH3
CH3
O
H2N CH C
CH3
O
OH +
OH
H2N CH C
R2
17.17
N
CH C
OH
CH
O
N
CH C
CH2
CH3
CH
CH3
H2N CH C
CH
OH
CH3
OH
O H
O
H2N CH C N
CH C
CH3
CH3
CH
CH3
17.23
17.25
17.27
CH C
O H
OH
H2N CH C
H
H
N
O
CH C
OH
CH3
Ser-Gly-Cys; Ser-Cys-Gly; Cys-Ser-Gly; Cys-Gly-Ser; Gly-Ser-Cys; Gly-Cys-Ser
legumes, nuts, and seeds
O
O
O
H2N CH C
H2N CH C
OH
H
Gly
OH
H2N CH C
CH CH3
CH2
CH3
OH
Val
17.29
CH3
Gly-Ala
O
CH3
OH
CH3
Ala-Gly
O H
N
OH
CH3
O
OH +
CH3
H2N CH C
O
CH C
CH3
O
17.21
CH3
R2
O H
H2N CH C
CH2
H2N CH C
CH3
Leu-Ala
O
CH3
N
R1
Ala-Leu
O H
H2N CH C
CH3
O H
H2N CH C
R1
17.19
CH3
OH
Ser
The abbreviated formula of the peptide is Gly-Ala-Val. Upon complete hydrolysis, the structures of the amino acids are
O
O
O
H2N CH C
OH
H2N CH C
H
H2N CH C
OH
CH CH3
CH3
Gly
CH3
Ala
Val
17.31
O H
C
O
H2N CH C
CH3
N
N
O H
CH C
CH2
OH
N
O
CH C
H
OH
OH
Appendix D Answers to Selected Questions and Problems
A-38
17.33
17.35
17.37
17.39
17.41
17.43
17.45
17.47
17.49
17.51
17.53
17.55
17.57
17.59
17.61
The primary structure is the amino acid sequence. The secondary structure results from intramolecular attractions such as hydrogen bonding and produces structural units such as the alpha helix and beta pleated sheets. The tertiary structure is the three-dimensional shape obtained from the folding of the peptide chain. Folding is caused by the intramolecular interactions of the secondary
structural units. The quaternary structure is used to describe the shape of a protein when it is composed of more than one peptide
chain. The quaternary structure is established through intermolecular interactions.
In myoglobin, a globular protein, the protein chain folds back on itself. Wool, a fibrous protein, is primarily composed of chains of
alpha helices.
tertiary structure
Hydrogen bonding holds together the twisted-helical shape of the peptide chain.
Fibrous proteins have long rod-like shapes. In globular proteins, the peptide chain folds back onto itself giving the protein a
“glob”-like shape.
Denaturation disrupts the attractive forces which give rise to secondary, tertiary, and quaternary structures and causes the “unfolding” of the protein.
Ribonucleic acids and deoxyribonucleic acids. In deoxyribonucleic acids, the ribose unit lacks the hydroxide group at the
2 position.
A nucleic acid is composed of a chain of nucleotides.
ribose; adenine; RNA
The shapes and sizes of these pairs (A-T and C-G) complement each other and provide for the maximum degree of hydrogen
bonding.
AATCGGTCACGAT
UGAUUCUAG
Genetic information is encoded as complementary strands of nucleic acids. When genetic information is replicated, it copies one
strand using complementary base pairing of DNA nucleotides.
AAA
AUG-CUU-CCC-CAA-CGA-UUU-CCC-CGA
Met - Leu - Pro - Gln - Arg - Phe - Pro - Arg
(so the answer is Pro and Arg)
17.63
CH2OH
CHO
H
C
OH
HO
C
H
H
C
OH
H
C
OH
OH C
CH 2OH
H
C
H
O
H
OH
C
H
C
C
O
H
OH
H
OH
H
C
C
H
OH C
OH
C
H
B-glucose
Open chain
17.65
17.67
17.69
17.71
17.73
CH2OH
H
C
OH
OH
A-glucose
glucose, C6H12O6; sucrose, C12H22O11
structure A (glucose)
monosaccharide
Sucrose yields glucose and fructose. Starch yields only glucose.
In starches, the main chain repeats with the glucose molecules maintaining the same orientation in the chains. Amylopectin also
branches (which is not observed in cellulose).
CH2OH
H
C
C
CH2OH
O
H
H
C
C
C
O
H
OH
H
H
OH
H
C
C
C
C
H
OH
H
OH
O
H
C
amylose
Appendix D Answers to Selected Questions and Problems
A-39
In cellulose, the glucose units invert after every ether bond (circled). Note the position of the – CH2OH group.
CH2OH
H
C
C
O
O
H
OH
H
C
C
H
OH
C
C
H
H
H
OH
C
C
OH
H
H
C
O
CH2OH
H
H
C
C
O
CH2OH
C
O
H
OH
H
C
C
H
OH
O
C
C
H
H
H
OH
C
C
OH
H
H
C
O
H
C
CH2OH
cellulose
17.75
17.77
17.79
17.81
17.83
17.85
17.87
Before we can taste the sugars in starch, the glucose monomers must be cleaved from the terminal of the amylose or amylopectin.
The fatty acids in fat molecules tend to be completely or almost completely saturated (few or no C=C double bonds). The fatty
acids in oil molecules are generally polyunsaturated and contain several C=C double bonds on each fatty acid chain.
triacyl glycerol (tristearate)
O
H2C
C O
O
HC
C O
O
H2C C O
It would be a fat.
phospholipid
They are both lipids, and all steroids are derived from
cholesterol.
O
O H
O H
H2N CH C
17.95
C
N
CH
CH2
CH2
OH
OH
OH
O H
H2N CH C
17.93
CH
CH2
17.89
17.91
N
N
O H
CH
CH CH3
CH2
CH3
SH
C
N
C
OH
O
CH
C
OH
CH2
OH
Proteins perform many functions; one of the most important is to catalyze the biochemical reactions. They are also responsible for
the machinery of the cell which regulates cellular processes such as transcription, translation, immune response, and energy production. Carbohydrates provide energy (i.e., glucose and metabolism) and play important structural roles in plants (cellulose).
DNA and RNA each have a sugar molecule, base, and a phosphate unit. The sugar in DNA is deoxyribose. The DNA contains all of
the genetic information. RNA, derived from DNA through transcription, contains only information for a specific protein.
Genetic therapy can be used to produce novel substances from organisms. An example is human insulin production from bacteria.
Plants can be bred to tolerate various chemical agents (i.e., herbicides) using this technique.