Multiple Choice
Identify the
letter of the choice that best completes the statement or answers the question.
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1.
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In
the following reaction,
2
Fe3+(aq) + Zn(s) →
2 Fe2+(aq) + Zn2+(aq)
a. | Fe3+(aq) is the reducing agent and Zn(s) is the oxidizing
agent. | b. | Zn(s) is the reducing agent and Fe3+(aq) is the
oxidizing agent. | c. | Fe3+(aq) is the reducing agent and
Fe2+(aq) is the oxidizing agent. | d. | Zn(s) is the
reducing agent and Zn2+(aq) is the oxidizing agent. | e. | Zn(s) is the
reducing agent and Fe2+(aq) is the oxidizing agent. | | |
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2.
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The
following reaction occurs spontaneously.
2 H+(aq) + Ca(s)
→ Ca2+(aq) +
H2(g)
Write the balanced reduction
half-reaction. a. | 2
H+(aq) + 2 e- → H2(g) | b. | 2
H+(aq) → H2(g) + 2 e- | c. | H2(g)
→2
H+(aq) + 2 e- | d. | Ca(s) + 2 e-
→
Ca2+(aq) | e. | Ca(s) → Ca2+(aq) + 2 e- | | |
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3.
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The
following reaction occurs spontaneously.
3 Cu2+(aq) + 2 Fe(s)
→ 2
Fe3+(aq) + 3 Cu(s)
Write the
balanced oxidation half-reaction. a. | 2 Fe(s) → 2 Fe3+(aq) + 6
e- | b. | 2 Fe(s) + 6 e-
→ 2
Fe3+(aq) | c. | 3 Cu2+(aq) + 6 e-
→ 3
Cu(s) | d. | 3 Cu2+(aq)
→ 3 Cu(s) + 6 e- | e. | 3 Cu(s)
→ 3
Cu2+(aq) + 6 e- | | |
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4.
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Write
a balanced half-reaction for the reduction of ClO3-(aq) to Cl2(g) in
an acidic solution. a. | 2 ClO3-(aq) + 6 H+(aq) + 10
e- → Cl2(g) + 6 OH-(aq) | b. | 2
ClO3-(aq) + 12 H+(aq) + 5 e-
→
Cl2(g) + 6 H2O(l) | c. | 2 ClO3-(aq) + 10 e-
→
Cl2(g) + 6 H2O(l) + 3 O2(g) | d. | 2 ClO3-(aq) + 12 H+(aq) + 10
e- → Cl2(g) + 6 H2O(l) | e. | 2 ClO3-(aq) + 18 H+(aq)
→
Cl2(g) + 6 H3O+(aq) | | |
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5.
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Write
a balanced half-reaction for the reduction of permanganate ion, MnO4-, to
MnO2 in a basic solution. a. | MnO4-(aq) + 4 OH-(aq) + 3
e- → MnO2(s) + 2 H2O(l) + 2
O2(g) | b. | MnO4-(aq) + 2 OH-(aq) + 3
e- → MnO2(s) + 2 HO2(aq) | c. | MnO4-(aq) + 3 e-
→MnO2(s) + O2(g) | d. | MnO4-(aq) + 2 H+(aq) + 3
e- → MnO2(s) + 2 OH-(aq) | e. | MnO4-(aq) + 2 H2O(l) + 3
e- → MnO2(s) + 4 OH-(aq) | | |
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6.
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All
of the following statements concerning voltaic cells are true EXCEPT a. | a salt bridge
allows cations and anions to move between the half-cells. | b. | electrons flow
from the cathode to the anode. | c. | reduction occurs at the cathode. | d. | a voltaic cell
can be used as a source of energy. | e. | a voltaic cell consists of two-half
cells. | | |
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7.
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What
is the correct cell notation for a voltaic cell based on the reaction
below?
Ni2+(aq) + Zn(s)
→Ni(s) +
Zn2+(aq)
a. | Zn(s) ï Zn2+(aq)  Ni2+(aq) ï Ni(s) | b. | Zn(s) 
Zn2+(aq), Ni2+(aq) ï Ni(s) | c. | Ni(s) 
Ni2+(aq), Zn2+(aq) 
Zn(s) | d. | Ni(s) ï Zn2+(aq)  Ni2+(aq) ï Zn(s) | e. | Ni(s) ï Ni2+(aq)  Zn2+(aq) ï Zn(s) | | |
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8.
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Write
a balanced net ionic equation for the overall reaction represented by the cell notation
below.
Zn ï ZnCl2(aq)  HCl(aq) ï H2(g) ï Pt
a. | 2 H+(aq) + Zn(s)
→
H2(g) + Zn2+(aq) | b. | H2(g) + Zn2+(aq)
→ 2
H+(aq) + Zn(s) | c. | H2(g) + ZnCl2(aq)
→ 2 HCl(aq) +
Zn(s) | d. | Pt(s) + Zn2+(aq)
→
Pt2+(aq) + Zn(s) | e. | Pt(s) + ZnCl2(aq)
→
PtCl2(aq) + Zn(s) | | |
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9.
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Use
the standard reduction potentials below to determine which element or ion is the best oxidizing
agent.
Br2(l) + 2
e- → 2 Br-(aq) E° = +1.08
V
Hg2+(aq) + e- → Hg(l) E° = +0.86 V
Ni2+(aq) + 2 e- → Ni(s) E° = -0.25 V
a. | Br2 | b. | Br- | c. | Hg2+ | d. | Ni2+ | e. | Ni | | |
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10.
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Consider the following half-reactions:
Cu2+(aq) + 2 e-
→ Cu(s)
E° = +0.34
V
Sn2+(aq) + 2 e- → Sn(s) E° = -0.14 V
Fe2+(aq) + 2 e- → Fe(s) E° = -0.44 V
Al3+(aq) + 3 e- → Al(s) E° = -1.66 V
Mg2+(aq) + 2 e- → Mg(s) E° = -2.37 V
Which of the above metals or metal ions will reduce
Fe2+(aq)? a. | Cu(s) and Sn(s) | b. | Cu2+(aq) and Sn2+(aq) | c. | Al3+(aq) and Mg2+(aq) | d. | Al(s) and
Mg(s) | e. | Sn(s) and Al3+(aq) | | |
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11.
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Given
the following two half-reactions, determine which overall reaction is spontaneous and calculate its
standard cell potential.
Sn4+(aq) + 2 e- → Sn2+(aq) E° = +0.15
V
Ag+(aq) + e- → Ag(s) E° = +0.80 V
a. | Sn4+(aq) + 2 Ag(s) → Sn2+(aq) + 2 Ag+(aq)  = -0.65 V | b. | Sn4+(aq) + 2 Ag(s) → Sn2+(aq) + 2 Ag+(aq)  = +0.95 V | c. | Sn2+(aq) + 2 Ag+(aq)
→ Sn4+(aq) + 2 Ag(s)  = -0.65 V | d. | Sn2+(aq) + 2 Ag+(aq)
→
Sn4+(aq) + 2 Ag(s)  =
+0.65 V | e. | Sn2+(aq) + 2 Ag+(aq)
→
Sn4+(aq) + 2 Ag(s)  =
+0.95 V | | |
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12.
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Calculate  for
the reaction below,
Tl+(aq) + 2 Cl-(aq)
→2 Tl(s) +
Cl2(g)
given the following standard reduction
potentials.
Cl2(g) + 2 e-
→ 2
Cl-(aq) E° = +1.36 V
Tl+(aq) + e- → Tl(s) E° = -0.34 V
a. | -2.04
V | b. | -1.70
V | c. | 1.02
V | d. | +1.70
V | e. | +2.04
V | | |
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13.
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Calculate  for
the reaction below,
5 Cd(s) + 2 MnO4-(aq) + 16
H+(aq) →5 Cd2+(aq) + 2 Mn2+(aq) + 8
H2O(l)
given the following standard reduction
potentials.
MnO4-(aq) + 8 H+(aq)
+ 5 e- → Mn2+(aq) + 4 H2O(l)
E° = +1.52
V
Cd2+(aq) + 2 e- → Cd(s) E° = -0.40 V
a. | +1.04
V | b. | +1.12V | c. | +1.92 V | d. | +5.04
V | e. | +8.40
V | | |
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14.
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Calculate  for
the electrochemical cell below,
Pb(s)
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PbSO4(s) ï Pb2+(aq) 
Ag+(aq) ï Ag(s)
given the following standard reduction
potentials.
Ag+(aq) + e-
→ Ag(s)
E° = +0.799
V
PbSO4(s) + 2 e- → Pb(s) + SO42-(aq) E° = -0.356
V
a. | -1.954
V | b. | -1.155
V | c. | -0.443
V | d. | +0.443
V | e. | +1.155
V | | |
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15.
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A
Faraday, F, is defined as a. | the charge on a single electron. | b. | the charge, in
coulombs, carried by one mole of electrons. | c. | the voltage
required to reduce one mole of reactant. | d. | the current required to reduce one mole of
reactant. | e. | the charge passed by one ampere of current in one
second. | | |
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16.
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Calculate E for the following electrochemical cell at 25 °C
Pt(s) ï H2(g, 1.00 atm) ï
H+(aq, 1.00 M) 
Cu2+(aq, 0.315 M) ï Cu(s)
given the following standard reduction
potentials.
Cu2+(aq) + 2 e-
→ Cu(s)
E° = +0.337
V
2 H+(aq) + 2 e- →H2(g) E° = 0.000
V
a. | +0.307
V | b. | +0.322
V | c. | +0.337
V | d. | +0.351
V | e. | +0.367
V | | |
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17.
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Calculate E for the following electrochemical cell at 25 °C
Pt(s) ï Sn2+(aq, 0.50 M), Sn4+(aq, 0.50 M)  I-(aq, 0.15 M) ï AgI(s)
ï
Ag(s)
given the following standard reduction
potentials.
AgI(s) + e-
→ Ag(s) +
I-(aq) E° = -0.15 V
Sn4+(aq) + 2 e- → Sn2+(aq) E° = +0.15
V
a. | -0.35
V | b. | -0.32
V | c. | -0.30
V | d. | -0.25
V | e. | +0.05
V | | |
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18.
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Calculate the cell potential, at 25 °C, based upon the overall reaction
3 Cu2+(aq) + 2 Al(s)
→ 3 Cu(s) + 2
Al3+(aq)
if [Cu2+] = 0.75 M and [Al3+] =
0.0010 M. The standard reduction potentials are as follows:
Cu2+(aq) + e-
→ Cu(s)
E° = +0.34
V
Al3+(aq) + 3 e- →Al(s) E° = -1.66 V
a. | -2.11
V | b. | -1.26
V | c. | +1.94
V | d. | +2.06
V | e. | +2.11
V | | |
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19.
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Which
factor will decrease the potential of the following electrochemical cell?
Pt ï Sn4+(aq, 1.0 M), Sn2+(aq, 1.0 M)  Cu2+(aq, 0.200 M) ï
Cu
a. | switching from a
platinum to a graphite anode | b. | decreasing the size of the cathode | c. | increasing the
concentration of Cu2+ | d. | decreasing the concentration of
Sn4+ | e. | increasing the temperature of the
cell | | |
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20.
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 for the following galvanic cell is +0.254
V.
Hg22+(aq) + 2 I-(aq)
→ 2
Hg(l) +
I2(s)
What is ΔG° for this reaction? a. | -49.0
kJ | b. | -24.5
kJ | c. | +24.5
kJ | d. | +49.0
kJ | e. | +197
kJ | | |
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21.
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Calculate ΔG° for the disproportionation reaction of Cu+ at 25
°C,
2 Cu+(aq)
→
Cu2+(aq) + Cu(s)
given the
following thermodynamic information.
Cu+(aq) + e- → Cu(s) E° = +0.518 V
Cu2+(aq) + 2 e- → Cu(s) E° = +0.337 V
a. | -165
kJ | b. | -135
kJ | c. | -34.9
kJ | d. | +17.5
kJ | e. | +135
kJ | | |
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22.
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If
Δ G° for the following reaction is -324 kJ, calculate  .
Cr2O72-(aq) + 2 Fe(s) + 14
H+(aq) → 2 Cr3+(aq) + 2 Fe3+(aq) + 7
H2O(l)
a. | -3.36 V | b. | -1.12
V | c. | +0.0201
V | d. | +0.560
V | e. | +1.68
V | | |
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23.
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Calculate the equilibrium constant for the following reaction at 25 °C,
Co2+(aq) + 2 Cr2+(aq)
→ Co(s) + 2
Cr3+(aq)
given the following thermodynamic
information.
Co2+(aq) + 2 e-
→ Co(s)
E° = -0.28
V
Cr3+(aq) + e- → Cr2+(aq) E° = -0.41
V
a. | 5
ｘ
10-24 | b. | 7
ｘ 10-11 | c. | 4
ｘ
10-5 | d. | 2
ｘ 104 | e. | 1
ｘ
1010 | | |
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24.
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Given
the following standard reduction potentials,
Fe2+(aq) + 2 e-
→ Fe(s)
E° = -0.440
V
FeS(s) + 2 e- → Fe(s) + S2-(aq) E° = -1.010
V
determine the Ksp for FeS(s) at 25 °C. a. | 9
ｘ
10-50 | b. | 5
ｘ 10-20 | c. | 1
ｘ
10-14 | d. | 2
ｘ 10-10 | e. | 2
ｘ
1020 | | |
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25.
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What
charge, in coulombs, is required to deposit 0.301 g Cu(s) from a solution of
Cu2+(aq)? a. | 9.82
ｘ 10-8 C | b. | 3.97
ｘ10-4
C | c. | 228
C | d. | 457
C | e. | 914
C | | |
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