Electrochemical Cells
1. Using reduction potentials answer the following.
a. What is the strongest oxidizing agent?
b. What is the strongest reducing agent?
c. Will Cl2 be able to reduce Cr3+?
d. Will Pb2+ oxidize Ni?
e. Will Au dissolve in an HCl solution?
f. Will Zn dissolve in an HCl solution?
g. What can oxidize Al but not Fe?
h. What can reduce I2 but not H+?
2. Consider the following reaction at 75 °C.
Pb2+(aq) + 2 Cr2+(aq)  Pb(s) + 2 Cr3+(aq)
a. Calculate the standard potential
b. Calculate the standard change in free energy
c. Calculate the equilibrium constant
d. If the above cell is allowed to operate spontaneously, will the voltage
increase, decrease or remain the same?
e. If the above cell is allowed to operate spontaneously, what will happen to
the concentrations at the electrodes?
f. Calculate the initial cell voltage for the above reaction if the initial
concentrations are [Pb2+] = 0.25 M, [Cr2+] = 0.20M and [Cr3+] = 0.005 M.
3.
Consider the following cell: Cu (s) | Cu2+ (0.001 M) || Br2 (l) | Br-(aq) (? M)
a. Assign the electrodes
b. Determine the direction of electron flow and the direction of the current
c. Describe the flow of ions in the salt bridge
d. Calculate the standard potential
e. At 25 °C the measured cell voltage is 0.975 V. Calculate the concentration
of the bromide ion.
4.
Consider the following cell: Al(s) | Al3+ (1.0 M) | | Pb2+ (1.0 M) | Pb(s)
Calculate the cell potential after the reaction has operated long enough for the
[Al3+] to have changed by 0.3 M at 25 °C.
5. Consider a cell with 0.0001 M Fe2+/Fe (s) and 10 M Fe2+/Fe(s). Determine the
direction of electron flow, assign the electrodes and calculate the cell voltage at
25 °C. What are the concentrations at equilibrium?
6. What mass of Co forms from a solution of Co2+ when a current of 15 amps is
applied for 1.15 hours?
7. A solution of iron chloride underwent electrolysis for 2 hrs at 10 amps yielding
20.84 g of iron. What is the molecular formula for the iron chloride?
8. Consider a cell with 0.0001 M Fe2+/Fe (s) and 10 M Fe2+/Fe(s). Determine the
direction of electron flow, assign the electrodes and calculate the cell voltage at
25 °C.
9. What volumes of H2 and O2 at STP are produced from the electrolysis of water by
a current of 3.5 amps in 15 minutes?
Hierarchy for Assigning Oxidation Numbers/States
1.
2.
3.
4.
5.
6.
7.
8.
9.
Atoms in their elemental form = 0
Monoatomic ions = charge
The sum of the oxidation numbers for compounds = 0
The sum of the oxidation numbers for polyatomic ions = net charge
Hydrogen in a compound or polyatomic ion = +1
Fluorine in a compound = -1
Oxygen in a compound or polyatomic ion = -2
Chlorine, Bromine and Iodine = -1
Nitrogen in a compound or polyatomic ion = -3
Red-ox Reactions
1. A reaction where there is a transfer of electrons
2. Oxidation is when an element loses electrons and is noted by an increase in the oxidation
number
3. Reduction is when an element gains electron and is noted by a decrease in the oxidation
number
4. The number of electrons gained must equal the number of electrons lost
5. An oxidizing agent is the substance getting reduced
6. A reducing agent is the substance getting oxidized
Electrochemical Cells
1.
Galvanic or Voltaic – spontaneously produce current (ε > 0)
2. Electrolytic or Electrolysis – require a current to run a non-spontaneous redox (ε < 0)
3. Anode – where oxidation occurs – if the electrode is an active metal it will corrode as the
cell discharges – anions in the salt bridge flow to this side
4. Cathode – where reduction occurs – if a metal ion is reduced to it’s ground state it will
plate out on the electrode – cations in the salt bridge flow to the side
5. For all cells electrons flow from anode to cathode or the current flows from cathode to
anode
Balancing Red-ox Reactions
1. Separate into two half reactions
2. Balance all elements EXCEPT H and O
3. Balance the oxygen by adding H20
4. Balance the hydrogen by adding H+
5. Balance the net charges by adding electrons (e-) to the more positive side
6. Make the electrons lost equal to the electrons gained in each half reaction by
multiplying the half reactions
7. Add two half reactions back together canceling out electrons, H+ and H2O
8. If acidic stop. If basic add an OH- to each side for every H+. The OH- cancels out
the H+ making water, which then needs to be adjusted for.
Table of Standard Reduction Potentials
Half-Reaction
F2 + 2 e -  2 F H2O2 + 2 H+ + 2 e-  2 H2O
MnO4- + 8 H+ + 5 e-  Mn2+ + 4 H2O
Au3+ + 3 e-  Au
Cl2 + 2 e-  2 ClCr2O72- + 14 H+ + 6 e-  2 Cr3+ + 7 H2O
O2 + 4 H+ + 4 e-  2 H2O
Br2 + 2 e-  2 BrNO3- + 4 H+ + 3 e-  NO + 2 H2O
Ag+ + e-  Ag
Fe3+ + e-  Fe2+
I2 + 2 e-  2 ICu+ + e-  Cu
O2 + 2 H2O + 4 e-  4 OHCu2+ + 2 e-  Cu
Cu2+ + e-  Cu+
2 H+ + 2 e-  H2
Pb2+ + 2 e-  Pb
Sn2+ + 2 e-  Sn
Ni2+ + 2 e-  Ni
Co2+ + 2 e-  Co
Fe2+ + 2 e-  Fe
Cr3+ + e-  Cr2+
Cr3+ + 3 e-  Cr
Zn2+ + 2 e-  Zn
2 H2O + 2 e-  H2 + 2 OHCr2+ + 2 e-  Cr
Mn2+ + 2 e-  Mn
Al3+ + 3 e-  Al
Sc3+ + 3 e-  Sc
Mg2+ + 2 e-  Mg
Na+ + e-  Na
Ca2+ + 2 e-  Ca
Ba2+ + 2 e-  Ba
K+ + e-  K
Li+ + e-  Li
ε°(volts)
2.87
1.78
1.49
1.42
1.36
1.33
1.23
1.06
0.96
0.80
0.77
0.53
0.52
0.40
0.34
0.16
0.00
-0.13
-0.14
-0.23
-0.28
-0.41
-0.42
-0.74
-0.76
-0.83
-0.91
-1.03
-1.71
-2.08
-2.37
-2.71
-2.76
-2.90
-2.93
-3.04