CHAPTER 18
ELECTROCHEMISTRY
Review of Oxidation-Reduction Reactions
15.
Oxidation: increase in oxidation number; loss of electrons
Reduction: decrease in oxidation number; gain of electrons
16.
See Table 4.2 in Chapter 4 of the text for rules for assigning oxidation numbers.
a. H (+1), O (2), N (+5)
b. Cl (1), Cu (+2)
c. O (0)
d. H (+1), O (1)
e. H(+1), O (2), C (0)
f.
g. Pb (+2), O (2), S (+6)
h. O (2), Pb (+4)
i.
j.
Na (+1), O (2), C (+3)
Ag (0)
O (2), C (+4)
k. (NH4)2Ce(SO4)3 contains NH4+ ions and SO42 ions. Thus cerium exists as the Ce4+ ion.
H (+1), N (3), Ce (+4), S (+6), O (2)
l.
17.
O (2), Cr (+3)
The species oxidized shows an increase in oxidation numbers and is called the reducing agent.
The species reduced shows a decrease in oxidation numbers and is called the oxidizing agent.
The pertinent oxidation numbers are listed by the substance oxidized and the substance
reduced.
Substance
Substance
Redox? Ox. Agent
Red. Agent
Oxidized
Reduced
a. Yes
H2O
CH4
CH4 (C, 4  +2)
H2O (H, +1  0)
b. Yes
AgNO3
Cu
Cu (0  +2)
AgNO3 (Ag, +1  0)
c. Yes
HCl
Zn
Zn (0  +2)
HCl (H, +1  0)
d. No; there is no change in any of the oxidation numbers.
30.
a. (Cu  Cu2+ + 2 e) × 3
NO3 NO + 2 H2O
(3 e + 4 H + NO3  NO + 2 H2O) × 2

+
Adding the two balanced half-reactions so that electrons cancel:
CHAPTER 18
ELECTROCHEMISTRY
689
3 Cu  3 Cu2+ + 6 e
6 e + 8 H+ + 2 NO3  2 NO + 4 H2O
3 Cu(s) + 8 H+(aq) + 2 NO3(aq)  3 Cu2+(aq) + 2 NO(g) + 4 H2O(l)
b. (2 Cl  Cl2 + 2 e) × 3
Cr2O72  2 Cr3+ + 7 H2O
6 e + 14 H + Cr2O72  2 Cr3+ + 7 H2O

+
Add the two half-reactions with six electrons transferred:
6 Cl  3 Cl2 + 6 e
6 e + 14 H + Cr2O72  2 Cr3+ + 7 H2O

+
14 H+(aq) + Cr2O72(aq) + 6 Cl(aq)  3 Cl2(g) + 2 Cr3+(aq) + 7 H2O(l)
c.
Pb  PbSO4
Pb + H2SO4  PbSO4 + 2 H+
Pb + H2SO4  PbSO4 + 2 H+ + 2 e
PbO2  PbSO4
PbO2 + H2SO4 → PbSO4 + 2 H2O
2 e + 2 H+ + PbO2 + H2SO4 → PbSO4 + 2 H2O
Add the two half-reactions with two electrons transferred:
2 e + 2 H+ + PbO2 + H2SO4  PbSO4 + 2 H2O
Pb + H2SO4  PbSO4 + 2 H+ + 2 e
Pb(s) + 2 H2SO4(aq) + PbO2(s)  2 PbSO4(s) + 2 H2O(l)
This is the reaction that occurs in an automobile lead-storage battery.
d.
Mn2+  MnO4
(4 H2O + Mn2+  MnO4 + 8 H+ + 5 e) × 2
NaBiO3

Bi3+
+
Na+
3+
+
NaBiO3 → Bi + Na
6 H+ + NaBiO3  Bi3+ + Na+ + 3 H2O
(2 e + 6 H+ + NaBiO3  Bi3+ + Na+ + 3 H2O) × 5
10 e
8 H2O + 2 Mn2+  2 MnO4 + 16 H+ + 10 e
+ 30 H+ + 5 NaBiO3  5 Bi3+ + 5 Na+ + 15
H2O
8 H2O + 30 H+ + 2 Mn2+ + 5 NaBiO3  2 MnO4 + 5 Bi3+ + 5 Na+ + 15 H2O + 16 H+
Simplifying:
14 H+(aq) + 2 Mn2+(aq) + 5 NaBiO3(s) → 2 MnO4(aq) + 5 Bi3+(aq) + 5 Na+(aq) +
7 H2O(l)
e.
H3AsO4 AsH3
H3AsO4  AsH3 + 4 H2O
8 e + 8 H+ + H3AsO4  AsH3 + 4 H2O
8 e + 8 H+ + H3AsO4  AsH3 + 4 H2O
(Zn  Zn2+ + 2 e) × 4
CHAPTER 18
ELECTROCHEMISTRY
690
4 Zn  4 Zn2+ + 8 e
8 H+(aq) + H3AsO4(aq) + 4 Zn(s)  4 Zn2+(aq) + AsH3(g) + 4 H2O(l)
31.
Use the same method as with acidic solutions. After the final balanced equation, convert H+ to
OH as described in Section 18.1 of the text. The extra step involves converting H+ into H2O
by adding equal moles of OH to each side of the reaction. This converts the reaction to a basic
solution while still keeping it balanced.
a.
Al  Al(OH)4
4 H2O + Al  Al(OH)4 + 4 H+
4 H2O + Al  Al(OH)4 + 4 H+ + 3 e
MnO4  MnO2
3 e + 4 H+ + MnO4  MnO2 + 2 H2O
4 H2O + Al  Al(OH)4 + 4 H+ + 3 e
3 e + 4 H+ + MnO4  MnO2 + 2 H2O

2 H2O(l) + Al(s) + MnO4(aq)  Al(OH)4(aq) + MnO2(s)
H+ doesn’t appear in the final balanced reaction, so we are done.
b.
Cl2  Cl
2 e + Cl2  2 Cl
Cl2  OCl
2 H2O + Cl2  2 OCl + 4 H+ + 2 e

2 e + Cl2  2 Cl
2 H2O + Cl2  2 OCl + 4 H+ + 2 e
2 H2O + 2 Cl2  2 Cl + 2 OCl + 4 H+
Now convert to a basic solution. Add 4 OH to both sides of the equation. The 4 OH will
react with the 4 H+ on the product side to give 4 H2O. After this step, cancel identical
species on both sides (2 H2O). Applying these steps gives: 4 OH + 2 Cl2 
2 Cl + 2 OCl + 2 H2O, which can be further simplified to:
2 OH(aq) + Cl2(g)  Cl(aq) + OCl(aq) + H2O(l)
c.
NO2  NH3
6 e + 7 H + NO2  NH3 + 2 H2O

+
Al  AlO2
(2 H2O + Al  AlO2 + 4 H+ + 3 e) × 2
Common factor is a transfer of 6 e.
6e + 7 H+ + NO2  NH3 + 2 H2O
4 H2O + 2 Al  2
e
AlO2
+
8
H+

OH + 2 H2O + NO2 + 2 Al  NH3 + 2 AlO2 + H+ + OH
Reducing gives OH(aq) + H2O(l) + NO2(aq) + 2 Al(s)  NH3(g) + 2 AlO2(aq).
+
6
CHAPTER 18
ELECTROCHEMISTRY
691
Exercises
Galvanic Cells, Cell Potentials, Standard Reduction Potentials, and Free Energy
(Here, if you see the oxidation is represented by –E0 at most places and since you learned to
look at the red.potential, take the E0 values for these with the opposite signs if that helps)
35.
The reducing agent causes reduction to occur; it does this by containing the species which is
oxidized. Oxidation occurs at the anode, so the reducing agent will be in the anode compartment. The oxidizing agent causes oxidation to occur; it does this by containing the species
which is reduced. Reduction occurs at the cathode, so the oxidizing agent will be in the cathode
compartment. Electron flow is always from the anode compartment to the cathode
compartment.
48.
b.
Au3+ + 3 e  Au (reduction) (cathode)
(Ag  Ag+ + e) × 3(oxidation) (anode)
E° = 1.50 V
E° = 0.80 V
Au3+(aq) + 3 Ag(s)  Au(s) + 3 Ag+(aq)
E ocell = 0.70 V
Spontaneous (E ocell  0)
(Here in this chapter, the oxidation is represented by –E0 values and since you learned to
look at the red.potential, take the E0 values for these with the opposite signs if that helps)
60.
a. H2  2 H+ + 2 e− E° = 0.00 V; Ni2+ + 2 e−  Ni E° = 0.23 V


= 0.23 V; no, H2 cannot reduce Ni2+ to Ni at standard conditions ( E cell
< 0).
E cell
b. Fe2+  Fe3+ + e− E° = 0.77 V; VO2+ + 2 H+ + e−  VO2+ + H2O
E° = 1.00 V

= 1.00  0.77 = 0.23 V; yes, Fe2+ can reduce VO2+ at standard conditions.
E cell
c. Fe2+  Fe3+ + e−
E° = 0.77 V; Cr3+ + e−  Cr2+
E° = -0.50 V

= 0.50  0.77 = 1.27 V
E cell

No, Fe2+ cannot reduce Cr3+ to Cr2+ at standard conditions( E cell
< 0).
61.
Cl2 + 2 e  2 Cl
Pb2+ + 2 e  Pb
Na+ + e  Na
E° = 1.36 V
E° = 0.13 V
E° = 2.71 V
Ag+ + e  Ag
Zn2+ + 2 e  Zn
E° = 0.80 V
E° = 0.76 V
a. Oxidizing agents (species reduced) are on the left side of the preceding reduction halfreactions. Of the species available, Ag+ would be the best oxidizing agent since it has the
CHAPTER 18
ELECTROCHEMISTRY
692
largest E° value. Note that Cl2 is a better oxidizing agent than Ag+, but it is not one of the
choices listed.
b. Reducing agents (species oxidized) are on the right side of the reduction half-reactions. Of
the species available, Zn would be the best reducing agent since it has the largest E° value.
c.
SO42 + 4 H+ + 2 e → H2SO3 + H2O E° = 0.20 V; SO42 can oxidize Pb and Zn at
standard conditions. When SO42 is coupled with these reagents, E ocell is positive.
d. Al → Al3+ + 3 e E° = 1.66 V; Al can reduce Ag+ and Zn2+ at standard conditions
because E ocell > 0.
The Nernst Equation
H2O2 + 2 H+ + 2 e  2 H2O
(Ag  Ag+ + e) × 2
67.
E° = 1.78 V
E° = 0.80 V
H2O2(aq) + 2 H+(aq) + 2 Ag(s)  2 H2O(l) + 2 Ag+(aq)
E ocell = 0.98 V
a. A galvanic cell is based on spontaneous redox reactions. At standard conditions, this
reaction produces a voltage of 0.98 V. Any change in concentration that increases the
tendency of the forward reaction to occur will increase the cell potential. Conversely, any
change in concentration that decreases the tendency of the forward reaction to occur
(increases the tendency of the reverse reaction to occur) will decrease the cell potential.
Using Le Chatelier’s principle, increasing the reactant concentrations of H2O2 and H+ from
1.0 to 2.0 M will drive the forward reaction further to right (will further increase the
tendency of the forward reaction to occur). Therefore, Ecell will be greater than E ocell .
b. Here, we decreased the reactant concentration of H+ and increased the product concentration of Ag+ from the standard conditions. This decreases the tendency of the forward
reaction to occur, which will decrease Ecell as compared to E ocell (Ecell < E ocell ).
77.
a.
Fe2+ + 2 e−  Fe
Zn  Zn2+ + 2 e−
Fe2+(aq) + Zn(s)  Zn2+(aq) + Fe(s)
b.
E ocell = 0.32 V = 0.32 J/C
o
= (2 mol e)(96,485 C/mol e)(0.32 J/C) = 6.2 × 104 J = 62 kJ
G o   nFE cell
o
E cell

c.
E° = 0.44 V
E° = 0.76 V
0.0591
nE o
2(0.32)

 10.83, K = 1010.83 = 6.8 × 1010
log K, log K 
0.0591
0.0591
n
o
E cell  E cell

0.0591
0.0591
[ Zn 2 ]
log
log Q = 0.32 V 
n
n
[Fe 2 ]
CHAPTER 18
ELECTROCHEMISTRY
Ecell = 0.32 
693
0.0591
0.10
= 0.32  0.12 = 0.20 V
log
2
1.0  10 5
Electrolysis
93.
a. Al3+ + 3 e  Al; 3 mol e are needed to produce 1 mol Al from Al3+.

1.0 × 103 g Al × 1 mol Al  3 mol e  96,485C
26.98 g Al
b. 1.0 g Ni ×
mol Al
mol e

1s =
100 .0 C
1.07 × 105 s
= 30. hours
1 mol Ni
2 mol e 
96,485 C
1s
= 33 s




58.69 g Ni
mol Ni
100 .0 C
mol e
1 mol e 
96,485 C
1s
= 4.8 × 103 s = 1.3 hours



mol Ag
100 .0 C
mol e
The oxidation state of bismuth in BiO+ is +3 because oxygen has a 2 oxidation state in this
ion. Therefore, 3 moles of electrons are required to reduce the bismuth in BiO+ to Bi(s).
c. 5.0 mol Ag ×
94.
10.0 g Bi ×
95.
15 A =
1 mol Bi
3 mol e 
96,485 C
1s
= 554 s = 9.23 min




209 .0 g Bi
mol Bi
25.0 C
mol e
15 C 60 s 60 min
= 5.4 × 104 C of charge passed in 1 hour


s
min
h
a. 5.4 × 104 C ×
1 mol e 
1 mol Co 58.93 g Co
= 16 g Co


96,485 C 2 mol e 
mol Co
b. 5.4 × 104 C ×
1 mol e 
1 mol Hf 178 .5 g Hf
= 25 g Hf


96,485 C 4 mol e 
mol Hf
c. 2 I  I2 + 2 e; 5.4 × 104 C ×
1 mol I 2
253 .8 g I 2
1 mol e 
= 71 g I2



96,485 C 2 mol e
mol I 2
d. CrO3(l)  Cr6+ + 3 O2; 6 mol e are needed to produce 1 mol Cr from molten CrO3.
5.4 × 104 C ×
96.
1 mol e 
1 mol Cr 52.00 g Cr


= 4.9 g Cr
96,485 C 6 mol e 
mol Cr
Al is in the +3 oxidation in Al2O3, so 3 mol e are needed to convert Al3+ into Al(s).
2.00 h ×
60 min
60 s 1.00  10 6 C 1 mol e 
1 mol Al 26.98 g Al





= 6.71 × 105 g
h
min
s
96,485 C 3 mol e 
mol Al