91. What is the value of the reaction quotient, Q, for the voltaic cell constructed from the
following two half-reactions when the Zn2+ concentration is 0.0120 M and the Ag+
concentration is 1.31 M?
A)
9.16 10–3
B)
6.99 10–3
C)
109
D)
1.25 10–2
E)
143
92. Given:
What is the cell potential at 25°C for the following cell?
Cr(s) | Cr3+(0.010 M) || Ag+(0.00020 M) | Ag(s)
A)
0.95 V
B)
2.09 V
C)
1.51 V
D)
1.36 V
E)
1.72 V
93. A cell consists of a magnesium electrode immersed in a solution of magnesium chloride and
a silver electrode immersed in a solution of silver nitrate. The two half-cells are connected
by means of a salt bridge. It is possible to increase the voltage of the cell by
A)
decreasing the concentration of Mg2+ and increasing the concentration of Ag+.
B)
adding sodium chloride to both half-cells.
C)
increasing the size of the Mg electrode and decreasing the size of the Ag electrode.
D)
decreasing the size of the Mg electrode and increasing the size of the Ag electrode.
E)
increasing the concentration of Mg2+.
94. A voltaic cell is made by placing an iron electrode in a compartment in which the Fe2+
concentration is 2.5 10–5 M and by placing a Pt electrode in the other compartment, in
which the H+ concentration is 3.70 M and = 1.00 atm. The Fe2+/Fe half-cell reduction
potential is –0.41 V, and the H+/H2 half-cell reduction potential is 0.00 V. What is E for the
cell at 25oC?
A)
0.10 V
B)
0.58 V
C)
0.27 V
D)
0.24 V
E)
0.56 V
95. The following cell is initially at standard-state conditions. Which of the following
statements is true after the cell is allowed to discharge (do work) for a period of time?
Zn2+(aq) + 2e– Zn(s); E° = –0.76 V
Cu2+(aq) + 2e– Cu(s); E° = 0.34 V
A)
Initially Ecell = – 1.10 V, and it will become more negative with time.
B)
Ecell does not change with time.
C)
Initially Ecell = – 1.10 V, and it will become more positive with time.
D)
Initially Ecell = + 1.10 V, and it will become more negative with time.
E)
Initially Ecell = + 1.10 V, and it will becomes more positive with time.
Test Bank General Chemistry, 10th edition 35
96. In order to determine the identity of a particular transition metal (M), a voltaic cell is
constructed at 25°C with the anode consisting of the transition metal as the electrode
immersed in a solution of 0.0977 M M(NO3)2, and the cathode consisting of a copper
electrode immersed in a 1.00 M Cu(NO3)2 solution. The two half-reactions are as follows:
M(s) M2+(aq) + 2e–
Cu2+(aq) + 2e– Cu(s)
The potential measured across the cell is 0.77 V. What is the identity of the metal M?
Reduction Half-Reaction
E° (V)
Cu2+(aq) + 2e– Cu(s)
+0.34
Ni2+(aq) + 2e– Ni(s)
–0.23
Cd2+(aq) + 2e– Cd(s)
–0.40
Zn2+(aq) + 2e– Zn(s)
–0.76
Mn2+(aq) + 2e– Mn(s)
–1.18
A)
Cd
B)
Ni
C)
Cu
D)
Mn
E)
Zn
97. What is the reduction potential for the half-reaction Al3+(aq) + 3e– Al(s) at 25°C if
[Al3+] = 0.44 M and E° = –1.66 V?
A)
–1.66 V
B)
–1.65 V
C)
–1.84 V
D)
–1.72 V
E)
–1.67 V
98. In order to determine the identity of a particular lanthanide metal (M), a voltaic cell is
constructed at 25°C with the anode consisting of the lanthanide metal as the electrode
immersed in a solution of 0.0873 M MCl3, and the cathode consisting of a copper electrode
immersed in a 1.00 M Cu(NO3)2 solution. The two half-reactions are as follows:
M(s) M3+(aq) + 3e–
Cu2+(aq) + 2e– Cu(s)
The potential measured across the cell is 2.68 V. What is the identity of the metal?
Reduction Half-Reaction
E° (V)
Cu2+(aq) + 2e– Cu(s)
0.34
Ce3+(aq) + 3e– Ce(s)
–2.336
Tm3+(aq) + 3e– Tm(s)
–2.319
Eu3+(aq) + 3e– Eu(s)
–1.991
Gd3+(aq) + 3e– Gd(s)
–2.279
Sm3+(aq) + 3e– Sm(s)
–2.304
A)
Sm
B)
Gd
C)
Tm
D)
Eu
E)
Ce
99. Which of the following statements is true concerning the electrochemical cell described
below at 25oC?
Ni | || | Ni
A)
The cell reaction is spontaneous with a cell potential of 20.8 mV.
B)
The cell reaction is nonspontaneous with a cell potential of –0.23 V.
C)
The cell reaction is nonspontaneous with a cell potential of –10.6 mV.
D)
The cell reaction is spontaneous with a cell potential of 10.6 mV.
E)
The cell reaction is nonspontaneous with a cell potential of –20.8 mV.
100. Which of the following statements about batteries is false?
A)
A fuel cell is a galvanic cell for which the reactants are continuously supplied.
B)
Lithium–iodine batteries have low current, but extremely long life.
C)
Lead storage batteries contain lead at the anode and lead coated with lead dioxide
at the cathode.
D)
It is not safe to recharge a nickel–cadmium battery.
E)
Dry cell batteries are used in flashlights and portable radios.
101. A piece of iron half-immersed in a sodium chloride solution will corrode more rapidly than
a piece of iron half-immersed in pure water, because
A)
the ions, which are not present in pure water, balance the accumulation of charge
in each of the half-cells.
B)
the chloride ions increase the pH of the solution.
C)
the sodium ions oxidize the iron atoms.
D)
the chloride ions form a precipitate with iron.
E)
the chloride ions oxidize the iron atoms.
102. Cathodic protection results when
A)
iron is amalgamated with mercury.
B)
iron is made amphoteric.
C)
iron is painted to protect it from corrosion.
D)
iron is tin-plated for use as a tin can.
E)
iron is attached to a more active metal.
103. Protection of iron (E(Fe2+/Fe) = −.23 V; (E(Fe3+/Fe) = −0.04 V) from corrosion can be
accomplished by making an electrical contact between iron and certain other metals.
Metal(s) that would provide protection is(are)
1.
Mg (E = −.38 V).
2.
Ni (E = −.23 V).
3.
Zn (E = −.76 V).
4.
Pb(E = −.13 V).
A)
1 only
B)
4 only
C)
1 and 3
D)
2 and 4
E)
3 and 4
104. What half-reaction occurs at the cathode during the electrolysis of molten potassium
bromide?
A)
K(s) → K+(l) + e–
B)
Br2(l) + 2e– → 2Br–(l)
C)
2Br–(l) → Br2(l) + 2e–
D)
K+(l) + e– → K(s)
E)
2H2O(l) + 2e– → H2(g) + 2OH–(l)
105. Molten magnesium chloride is electrolyzed using inert electrodes and reactions represented
by the following two half-reactions:
2Cl–(l) → Cl2(g) + 2e–
Mg2+(l) + 2e– → Mg(s)
Concerning this electrolysis, which of the following statements is true?
A)
The anions in the electrolyte undergo reduction.
B)
Mg2+ ions are reduced at the anode.
C)
Oxidation occurs at the cathode.
D)
Electrons pass through the metallic part of the circuit from Mg2+ ions to the Cl–
ion.
E)
Cl– ions are reducing agents.
106. What is one of the major products in the electrolysis of a saturated aqueous sodium chloride
solution?
Reduction Half-Reaction
E° (V)
Na+(aq) + e– Na(s)
–2.71
Mg2+(aq) + 2e– Mg(s)
–2.37
2H2O(l) + 2e– H2(g) + 2OH–(aq)
–0.83
O2(g) + 4H+(aq) + 4e– 2H2O(l)
1.23
Cl2(g) + 2e– 2Cl–(aq)
1.36
A)
Mg
B)
O2
C)
H2O
D)
Na
E)
Cl2
107. What is the half-reaction that occurs at the cathode during electrolysis of an aqueous
potassium iodide solution?
Reduction Half-Reaction
E° (V)
K+(aq) + e– K(s)
–2.93
Mg2+(aq) + 2e– Mg(s)
–2.37
2H2O(l) + 2e– H2(g) + 2OH–(aq)
–0.83
2H+(aq) + 2e– H2(g)
0.00
I2(s) + 2e– 2I–(aq)
0.54
O2(g) + 4H+(aq) + 4e– 2H2O(l)
1.23
A)
H2O(l) → ½ O2(g) + 2H+(aq) + 2e–
B)
2H2O(l) + 2e– → H2(g) + 2OH–(aq)
C)
K(s) → K+(aq) + e–
D)
K+(aq) + e– → K(s)
E)
I2(aq) + 2e– → 2I–(aq)
108. When an aqueous solution of lithium sulfate is electrolyzed, what are the expected products?
Reduction Half-Reaction
E° (V)
Li(s)
–3.04
2H2O(l) + 2e– H2(g) + 2OH–(aq)
–0.83
2H+(aq) + 2e– H2(g)
0.00
O2(g) + 4H+(aq) + 4e– 2H2O(l)
1.23
S2O82–(aq) + 2e– 2SO42–(aq)
2.01
A)
Li(s) and H2(g)
B)
H2(g), OH–(aq), O2(g), and H+(aq)
C)
O2(g), H+(aq), and Li(s)
D)
H2(g), OH–(aq), and Li(s)
E)
H2(g), OH–(aq), and S2O82–(aq)
Test Bank General Chemistry, 10th edition 40
109. When an aqueous solution of AgNO3 is electrolyzed, a gas is formed at the anode. What is
the identity of the gas?
Reduction Half-Reaction
E° (V)
2H2O(l) + 2e– H2(g) + 2OH–(aq)
–0.83
2H+(aq) + 2e– H2(g)
0.00
Ag+(aq) + e– Ag(s)
0.80
NO3–(aq) + 4H+(aq) + 3e– NO(g) + 2H2O(l)
0.96
O2(g) + 4H+(aq) + 4e– 2H2O(l)
1.23
A)
H2O
B)
NO
C)
Ag
D)
O2
E)
H2
110. What reaction occurs at the anode during the electrolysis of aqueous CuSO4?
Reduction Half-Reaction
E° (V)
2H2O(l) + 2e– H2(g) + 2OH–(aq)
–0.83
2H+(aq) + 2e– H2(g)
0.00
Cu2+(aq) + 2e– Cu(s)
0.34
O2(g) + 4H+(aq) + 4e– 2H2O(l)
1.23
A)
2H2O(l) → O2(g) + 4H+(aq) + 4e–
B)
Cu(s) → Cu2+(aq) + 2e–
C)
2H2O(l) + 2e– → H2(g) + 2OH–(aq)
D)
Cu2+(aq) + 2e– → Cu(s)
E)
2H+(aq) + 2e– → H2(g)
111. Which of the following statements is true concerning the electrolysis of a 1.0 M aqueous
solution of NaI?
Reduction Half-Reaction
E° (V)
Na+(aq) + e– Na(s)
–2.71
2H2O(l) + 2e– H2(g) + 2OH–(aq)
–0.83
2H+(aq) + 2e– H2(g)
0.00
I2(s) + 2e– 2I–(aq)
0.54
O2(g) + 4H+(aq) + 4e– 2H2O(l)
1.23
Test Bank General Chemistry, 10th edition 41
A)
The solution becomes more basic.
B)
Sodium is deposited at the cathode.
C)
Iodine is formed at the cathode.
D)
Oxygen is evolved at the anode.
E)
Hydrogen is evolved at the anode.
112. When Au is obtained by electrolysis from NaAu(CN)2, what is the minimum number of
coulombs required to produce 1.06 mol of gold?
A)
C
B)
C
C)
C
D)
C
E)
C
113. How many faradays are involved in the conversion of a mole of to ?
A)
5
B)
1
C)
2
D)
3
E)
4
114. How many faradays are required to convert a mole of AsO3– ions to AsH4+ ions?
A)
4
B)
7
C)
8
D)
6
E)
5
Test Bank General Chemistry, 10th edition 42
115. If an electrolysis plant operates its electrolytic cells at a total current of 1.0 106 A, how
long will it take to produce one metric ton (one million grams) of Mg(s) from seawater
containing Mg2+? (1 faraday = 96,485 coulombs)
A)
66 min
B)
0.55 year
C)
2.2 h
D)
2.2 days
E)
1.1 h
116. Gold (atomic mass = 197) is plated from a solution of chlorauric acid, HAuCl4; it deposits
on the cathode. Calculate the time it takes to deposit 0.64 g of gold, passing a current of 0.10
A. (1 faraday = 96,485 coulombs)
A)
54 min
B)
0.87 h
C)
0.29 h
D)
2.6 h
E)
none of these
117. Copper is electroplated from CuSO4 solution. A constant current of 4.39 A is applied by an
external power supply. How long will it take to deposit 1.00 ×102 g of Cu? The atomic mass
of copper is 63.546.
A)
19.2 h
B)
13.17 s
C)
1.48 days
D)
9.6 min
E)
2.74 h
118. How many moles of electrons are produced from a current of 15.6 A in 3.20 h?
A)
3.35 mol
B)
9.33 103 mol
C)
5.17 10–4 mol
D)
1.86 mol
E)
49.9 mol
119. In the electrolysis of an acid solution, oxygen can be produced by the following half-
reaction:
2H2O(l) → O2(g) + 4H+(aq) + 4e–
How many moles of O2 can be produced from a solution that was electrolyzed for 4.50 h
using a current of 5.80 A?
A)
mol
B)
mol
C)
mol
D)
0.243 mol
E)
0.974 mol
120. What mass of chromium could be deposited by electrolysis of an aqueous solution of
Cr2(SO4)3 for 160 min using a constant current of 15.0 A? (F = 96485 C/mol)
A)
0.431 g
B)
25.9 g
C)
232.8 g
D)
0.187 g
E)
38.8 g
121. A current of 15.0 A is passed through molten magnesium chloride for 15.0 h. How many
moles of magnesium metal could be produced via this electrolysis?
A)
0.0700 mol
B)
4.20 mol
C)
0.37 mol
D)
0.22 mol
E)
8.40 mol
122. A solution of is electrolytically reduced to Mn3+. A current of 7.79 A is passed
through the solution for 15.0 min. What is the number of moles of Mn3+ produced in this
process? (1 faraday = 96,486 coulombs)
A)
0.218 mol
B)
1.61 10–3 mol
C)
0.0242 mol
D)
0.0727 mol
E)
4.04 10–4 mol
123. Electrolysis of a molten salt with the formula MCl, using a current of 3.86 A for 16.2 min,
deposits 1.52 g of metal. Identify the metal. (1 faraday = 96,485 coulombs)
A)
Na
B)
Li
C)
Ca
D)
Rb
E)
K