Chapter 14 REVIEW
Part 1
(Page 664) 1. A 2. B 3. D 4. C 5. 1, 3, 4, 7 6. 2, 3, 5, 7 7. 1.49 8. B 9. D 10. A 11. A 12. C 13. 1.16 14. 1.51 Solutions7. E°cell = +1.36 V – (–0.13 V) = 1.49 V
According to the redox table, the standard cell potential is +1.49 V.
13. Cu2+(aq) + 2 e– o Cu(s)
0.325 A m
3.00 h 63.55 g/mol
9.65 u 104 C/mol
ne = C 3600 s
0.325 3.00 h
s u u 1 mol4
1 h 9.65 10 C
§ ·
¨ ¸ u u
© ¹ = 0.0364 mol
nCu = 1
0.0364 mol
u 2 = 0.0182 mol mCu = 0.0182 mol 63.55 g
1 mol
u = 1.16 g
or
mCu = 0.325 C e 3600 s 3.00 h
1 s u u 1 mol e
1 h
u 4
9.65 10 C eu
1 mol Cu u 2 mol e
63.55 g Cu 1 molCu u
= 1.16 g
According to stoichiometry and Faraday’s law, the mass of copper produced is 1.16 g.
14. SOA OA
Co2+(aq), NO3–(aq), H2O(l) SRA
E°cell = –0.28 V – (+1.23 V) = –1.51 V
According to the redox table, a minimum applied voltage of +1.51 V is required.
Part 2
(Page 665)15. (a) Electrolytes are aqueous electrical conductors that are in contact with electrodes.
Electrolytes usually contain some reactants and are part of the internal circuit inside a voltaic cell.
(b) A salt bridge completes a circuit by allowing anions to flow towards the anode and cations to flow towards the cathode, in order to maintain a neutral electric charge in the solutions.
(c) Electrodes are solid conducting metals, or carbon, that provide a surface for the half- reactions to occur and a location to connect the external wires.
(d) A connecting wire connects the electrodes to voltmeters, power supplies, or each other, which ultimately allows electrons to flow from the anode to the cathode.
16. (a) Ni(s) | Ni2+(aq) || Cu2+(aq) | Cu(s) SRA SOA cathode: Cu2+(aq) + 2 e–o Cu(s)
anode: Ni(s) o Ni2+(aq) + 2 e– net: Cu2+(aq) + Ni(s) o Cu(s) + Ni2+(aq) (b) Zn(s) | Zn2+(aq) || Cr2O72–(aq), H+(aq) | C(s)
SRA SOA
cathode: Cr2O72– (aq) + 14 H+ + 6 e–o 2 Cr3+(s) + 7 H2O(l) anode: 3 [Zn(s) o Zn2+(aq) + 2 e–]
net: Cr2O72– (aq) + 14 H+ + 3 Zn(s) o 2 Cr3+(s) + 7 H2O(l) + 3 Zn2+(aq) (c) Pt(s) | H2(g) | OH–(aq) || Ag+(aq) | Ag(s)
SRA SOA
cathode: 2 [Ag+ + e–o Ag(s)]
anode: H2(g) + 2 OH–(aq) o 2 H2O(l) + 2 e–
net: 2 Ag+ + H2(g) + 2 OH–(aq) o 2 Ag(s) + 2 H2O(l) 17.
19. (a) One form of cathodic protection is an impressed current—a small electric current that is forced by an external battery to flow toward the steel object. This forces the steel object to become the cathode where only reduction of some entity and not oxidation of iron occurs.
The sacrificial anode is a metal, that is more easily oxidized than iron in the steel so that when it is connected to the steel, (such as by galvanization or by wire), it is oxidized instead of the iron in the steel.
(b) From an economic perspective it is very costly to replace corroded metals, generally much more so than protecting them, so it is important to protect them from corrosion.
From an ecological perspective, not protecting metals from corrosion could easily result in the spilling of toxic materials into the environment.
20.
21. (a) Ni(s) o Ni2+(aq) + 2 e– (b) Ni2+(aq) + 2 e–o Ni(s)
(c) There is no net reaction as all entities cancel.
(d) E°cell = –0.26 V – (–0.26 V) = 0.00 V
22. (a) [There are a number of choices for design and chemicals that will produce 1.25 V or slightly larger.]
The electrolytes are 1.0 mol/L at a temperature of 25 °C.
(b) The design is sufficient because the voltage supplied satisfies the 1.25 V needed for the clock.
cathode: 2 [Ag+(aq) + e–o Ag(s)] E°r = +0.80 V
anode: Fe(s) o Fe2+(aq) + 2 e– E°r = –0.45 V net: 2 Ag+(aq) + Fe(s) o 2 Ag(s) + Fe2+(aq) E°cell = +1.25 V E°cell = +0.80 V – (–0.45 V)
= +1.25 V
(c) There are a number of possible issues with this cell. Standard cells often produce less voltage than is predicted, due to impurities that exist in the reactants. Impurities that are found in these reactants include the presence of oxide coatings, the effects of the porous boundary, and electrical resistance in the external circuit. Another issue is that when the cell is used as designed, the aqueous electrolytes can be spilled. Since aqueous solutions are exposed to the air, they may evaporate, which would affect the concentrations of electrolytes. Furthermore, the cell is not very portable and is very large compared with that of a commercial cell.
23. Al3+(aq) + e– o Al(l)
1.00 h m
Al 1.00 h
m 3600 s
u 1 h
150 k C e u 1 s
1 mol e
u 4
9.65 10 C eu
1 mol Al u 3 mol e
26.98 g Al 1 mol Al u
Al
= 50.3 kg m
According to Faraday’s law and stoichiometry, the mass of aluminium produced in the cell would be 50.3 kg.
24. cathode: Cu2+(aq) + e– o Cu(s)
24.0 h 24.68 g – 21.12 g = 3.56 g
I 63.55 g/mol
9.65 u 104 C/mol
Cu 3.56 g
n 1 mol
63.55 g
u 0.0560 mol
e
0.0560 mol 2 0.112 mol
n u 1
0.121 mol
I 4 C
9.65 10
u u mol 1
24.0 h
§ ·
¨ ¸ u
© ¹ 3600 s
u 1 h
0.125A
or
3.56 g Cu
I 1 mol Cu
u 63.55 g Cu
2 mol e u 1 mol Cu
9.65 10 C4
1 mol e
u u 1
24.0 h
u 1 h
u 3600 s
= 0.125 A
According to Faraday’s law and stoichiometry, the average current produced in the cell is 0.125 A.
25. Some examples of metal corrosion are:
Ɣ corrosion of oil and/or gas pipelines and storage tanks (environmental, health, and safety issues);
Ɣ corrosion of steel bridges and towers (safety issues);
Ɣ corrosion of reinforcing steel in concrete structures (environmental and safety issues) and Ɣ corrosion of the bodies of motor vehicles, boats, and airplanes (environmental and safety
issues).
Methods used to prevent corrosion of iron include barrier methods, which involve adding protective coatings to the iron, and cathodic protection, in which the iron is forced to become the cathode by supplying it with electrons using an impressed current or by adding a
sacrificial anode.
Some examples of desirable metal corrosion are: self-protecting metals, e.g. aluminium and zinc; attractive patina forming on copper roofs and bronze statues; and corrosion of sacrificial anodes to protect buried pipes and tanks.
