H14 Week 1
1. Balance the following redox reactions:
a) ___ Cd (s) + ___ AgCl(s) → ___ Cd2+ (aq) + Ag(s) + Cl- (aq)
b) ___H+ (aq) + ___MnO4− (aq) + ___Cl− (aq) → ___Mn2+ (aq) + ___Cl2(g) + ___H2O (l)
2. A constant current of 0.751 7 A is used to deposit 0.348 3 g of Co2+ as solid Co on a cathode surface. How many
minutes are needed for complete reduction?
3. A voltaic (galvanic) cell is constructed with an iron rod in a 1.0 M Fe2+ solution as the right-hand electrode, and
a silver rod in a 1.0 M Ag+ solution as the left-hand electrode.
a) Draw the complete cell.
b) Electrons in the external circuit flow from the ______________ (Ag or Fe) electrode to the _______________
(Ag or Fe) electrode.
c) Negative ions move in the salt bridge from the __________________ (Ag+/Ag or Fe2+/Fe) half-cell to the
___________________ (Ag+/Ag or Fe2+/Fe) half-cell.
d) The half-reaction at the anode is:
e) The half-reaction at the cathode is:
f) Write the balanced spontaneous reaction for this cell.
g) What is the value of E ° for the reaction in part (g)?
4. Consider the following table of standard reduction potentials.
Reduction Half-Reaction
E° (V)
A+ + e− → A
0.80
B2+ + 2e− → B
0.38
C2 + 2e− → 2C−
0.17
D3+ + 3e− → D
−1.36
a) Which substance is the strongest oxidizing agent?
b) Which substance is the strongest reducing agent?
c) Which substances can be oxidized by B2+?
d) Which substances can be reduced by C2?
e) Write a balanced equation for the overall cell reaction that produces a voltage of +1.53 V under standard
conditions.
f)
Write a balanced equation for the overall cell reaction that produces a voltage of –0.63 V under standard
conditions.
5. Consider a galvanic cell based on the reaction: Al3+(aq) + Mg(s) → Al(s) + Mg2+(aq)
Give the balanced cell reaction and calculate E° for the cell.
6. Determine Eocell and Ecell based on the following half-reactions:
VO2+ + 2H+ + e- → VO2+ + H2O E°= 1.00 V
Zn2+ + 2e- → Zn
E° = -0.76V
Where:
T = 25°C, [VO2+] = 2.0 M, [H+] = 0.50 M, [VO2+] = 1.0 x 10-2 M, [Zn2+] = 1.0 x 10-1 M
7. Consider the cell:
Pt(s) | H2(g, 0.100 bar) | H+(aq, pH = 2.54) || Cl−(aq, 0.200 M) | Hg2Cl2(s) | Hg() | Pt(s)
Use the following information to calculate the cell potential.
Hg2Cl2(s) + 2e− → 2Hg() + 2Cl−(aq)
E ° = 0.268 V
2H+(aq) + 2e− → H2(g)
E ° = 0.000 V
8. For each of the following reactions, calculate 𝐸 𝑜 𝑎𝑛𝑑 𝛥𝐺 𝑜
a) Ag+(aq) + Fe2+(aq) → Ag(s) + Fe3+(aq)
b) 2Al3+(aq) + 3Ni(s) → 2Al(s) + 3Ni2+(aq)
Table of Standard Reduction Potentials
Reduction Half-Reactions
F2(g) + 2e− → 2F−(aq)
O3(g) + 2H+(aq) + 2e− → O2(g) + H2O()
S2O82−(aq) + 2e− → 2SO42−(aq)
Co3+(aq) + e− → Co2+(aq)
H2O2(aq) + 2H+(aq) + 2e− → 2H2O()
Ce4+(aq) + e− → Ce3+(aq)
MnO4−(aq) + 8H+(aq) + 5e− → Mn2+(aq) + 4H2O(l)
Cl2(g) + 2e− → 2Cl−(aq)
Cr2O72−(aq) + 14H+(aq) + 6e− → 2Cr3+(aq) + 7H2O(l)
Potential, E°
(V)
2.890
2.075
2.01
1.92
1.763
O2(g) + 4H+(aq) + 4e− → 2H2O()
1.72
1.507
1.360 4
1.36
1.229 1
Br2() + 2e− → 2Br−(aq)
1.078
NO3−(aq) + 4H+(aq) + 3e− → NO(g) + 2H2O()
0.955
2Hg2+(aq) + 2e− → Hg22+(aq)
0.908
0.852
2+
−
Hg (aq) + 2e → Hg()
Ag+(aq) + e− → Ag(s)
Hg22+(aq) + 2e− → 2Hg()
Fe3+(aq) + e− → Fe2+(aq)
I2(s) + 2e− → 2I−(aq)
Cu+(aq) + e− → Cu(s)
Cu2+(aq) + 2e− → Cu(s)
AgCl(s) + e− → Ag(s) + Cl−(aq)
Cu2+(aq) + e− → Cu+(aq)
Sn4+(aq) + 2e− → Sn2+(aq)
2H+(aq) + 2e− → H2(g)
Fe3+(aq) + 3e− → Fe(s)
Pb2+(aq) + 2e− → Pb(s)
Sn2+(aq) + 2e− → Sn(s)
Ni2+(aq) + 2e− → Ni(s)
Cd2+(aq) + 2e− → Cd(s)
Fe2+(aq) + 2e− → Fe(s)
Cr3+(aq) + 3e− → Cr(s)
Zn2+(aq) + 2e− → Zn(s)
0.799 3
0.796
2H2O() + 2e− → H2(g) + 2OH−(aq)
0.771
0.535
0.518
0.339
0.222
0.161
0.139
0.000
−0.036
−0.126
−0.141
−0.236
−0.402
−0.44
−0.74
−0.762
−0.828 0
Al3+(aq) + 3e− → Al(s)
Mg2+(aq) + 2e− → Mg(s)
Na+(aq) + e− → Na(s)
Ca2+(aq) + 2e− → Ca(s)
K+(aq) + e− → K(s)
Li+(aq) + e− → Li(s)
−1.677
−2.360
−2.714 3
−2.868
−2.936
−3.040