1. Standard Reduction (Half-Cell) Potentials
Cell EMF
Standard Reduction (Half-Cell) Potentials
Consider Zn(s)  Zn2+(aq) + 2e-. We measure Ecell relative to the SHE (cathode):
Ecell = Ered(cathode) + Eox (anode)
0.76 V = 0 V + Eox(anode).
Therefore, Eox(anode) = V.
Standard reduction potentials must be written as reduction reactions:
Zn2+(aq) + 2e-  Zn(s), Ered = V.

2. 2Zn2+(aq) + 4e-  2Zn(s), Ered = -0.76 V.
Since Ered = V we conclude that the reduction of Zn2+ in the presence of the SHE is not spontaneous.
The oxidation of Zn with the SHE is spontaneous.
Changing the stoichiometric coefficient does not affect Ered.
Therefore,
2Zn2+(aq) + 4e-  2Zn(s), Ered = V.
Reactions with Ered > 0 are spontaneous reductions relative to the SHE.

3. Reactions with Ered < 0 are spontaneous oxidations relative to the SHE.
The larger the difference between Ered values, the larger Ecell.
In a voltaic (galvanic) cell (spontaneous) Ered(cathode) is more positive than Ered(anode).
Recall

4. Examples – Balance the following equations, and calculate E°cell for each:
Cr3+ + Cl2 (g)  Cr2O72- + Cl-
Cu2+ + Mg (s)  Mg2+ + Cu (s)
IO3- + Fe2+  Fe3+ + I2
Zn (s) + Ag+  Zn2+ + Ag

5. Oxidizing and Reducing Agents
The more positive Ered the stronger the oxidizing agent on the left.
The more negative Ered the stronger the reducing agent on the right.
A species on the higher to the left of the table of standard reduction potentials will spontaneously oxidize a species that is lower to the right in the table.
That is, F2 will oxidize H2 or Li; Ni2+ will oxidize Al(s).

7. Spontaneity of Redox Reactions
In a voltaic (galvanic) cell (spontaneous) Ered(cathode) is more positive than Ered(anode) since Or
More generally, for any electrochemical process
A positive E indicates a spontaneous process (galvanic cell).
A negative E indicates a nonspontaneous process.

8. Examples – Sketch the cells containing the following reactions
Examples – Sketch the cells containing the following reactions. Include E°cell, the direction of electron flow, direction of ion migration through salt bridge, and identify the anode and cathode:
Cr3+ + Cl2 (g)  Cr2O72- + Cl-
Cu2+ + Mg (s)  Mg2+ + Cu (s)
IO3- + Fe2+  Fe3+ + I2

9. EMF and Free-Energy Change
We can show that
G is the change in free-energy, n is the number of moles of electrons transferred, F is Faraday’s constant, and E is the emf of the cell.
We define
Since n and F are positive, if G > 0 then E < 0.

10. Examples – Balance, and calculate E°cell and ΔG° for each reaction:
Mg (s) + Au3+  Mg2+ + Au (s)
Al (s) + Ag+  Al3+ + Ag (s)
Cu2+ + Na (s)  Cu (s) + Na+

11. Cell EMF and Chemical Equilibrium
A system is at equilibrium when G = 0.
From the Nernst equation, at equilibrium and 298 K (E = 0 V and Q = Keq):

12. Batteries
A battery is a self-contained electrochemical power source with one or more voltaic cell.
When the cells are connected in series, greater emfs can be achieved.

13. Cathode: PbO2 on a metal grid in sulfuric acid:
Lead-Acid Battery
A 12 V car battery consists of 6 cathode/anode pairs each producing 2 V.
Cathode: PbO2 on a metal grid in sulfuric acid:
PbO2(s) + SO42-(aq) + 4H+(aq) + 2e-  PbSO4(s) + 2H2O(l)
Anode: Pb:
Pb(s) + SO42-(aq)  PbSO4(s) + 2e-

15. The overall electrochemical reaction is
PbO2(s) + Pb(s) + 2SO42-(aq) + 4H+(aq)  2PbSO4(s) + 2H2O(l)
for which
Ecell = Ered(cathode) + Eox (anode)
= ( V) + (0.356 V)
= V.
Wood or glass-fiber spacers are used to prevent the electrodes from touching.

16. 2NH4+(aq) + 2MnO2(s) + 2e-  Mn2O3(s) + 2NH3(aq) + 2H2O(l)
Alkaline Battery
Anode: Zn cap:
Zn(s)  Zn2+(aq) + 2e-
Cathode: MnO2, NH4Cl and C paste:
2NH4+(aq) + 2MnO2(s) + 2e-  Mn2O3(s) + 2NH3(aq) + 2H2O(l)
The graphite rod in the center is an inert cathode.
For an alkaline battery, NH4Cl is replaced with KOH.

17. Anode: Zn powder mixed in a gel:
Zn(s)  Zn2+(aq) + 2e-
Cathode: reduction of MnO2.

19. Direct production of electricity from fuels occurs in a fuel cell.
Fuel Cells
Direct production of electricity from fuels occurs in a fuel cell.
On Apollo moon flights, the H2-O2 fuel cell was the primary source of electricity.
Cathode: reduction of oxygen:
2H2O(l) + O2(g) + 4e-  4OH-(aq)
Anode:
2H2(g) + 4OH-(aq)  4H2O(l) + 4e-