1. Chapter 21
Electrochemistry

2. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
Example 21-1: Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes.

3. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
Example 21-2: Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in example 21-1.

4. Electrode Potentials for Other Half-Reactions
Example 21-4: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution?
Thus permanganate ions will oxidize iron (II) ions to iron (III) and are reduced to manganese (II) ions in acidic solution.

5. Electrode Potentials for Other Half-Reactions
Example 21-5: Will nitric acid, HNO3, oxidize arsenous acid, H3AsO3, in acidic solution? The reduction product of HNO3 is NO in this reaction.

6. The Nernst Equation
Example 21-6: Calculate the potential for the Cu2+/ Cu+ electrode at 250C when the concentration of Cu+ ions is three times that of Cu2+ ions.

7. The Nernst Equation
Example 21-7: Calculate the potential for the Cu2+/Cu+ electrode at 250C when the Cu+ ion concentration is 1/3 of the Cu2+ ion concentration.

8. The Nernst Equation
Example 21-8: Calculate the electrode potential for a hydrogen electrode in which the [H+] is 1.0 x 10-3 M and the H2 pressure is 0.50 atmosphere.

9. The Nernst Equation
Example 21-9: Calculate the initial potential of a cell that consists of an Fe3+/Fe2+ electrode in which [Fe3+]=1.0 x 10-2 M and [Fe2+]=0.1 M connected to a Sn4+/Sn2+ electrode in which [Sn4+]=1.0 M and [Sn2+]=0.10 M . A wire and salt bridge complete the circuit.

10. The Nernst Equation
Calculate the E0 cell by the usual procedure.

11. The Nernst Equation
Substitute the ion concentrations into Q to calculate Ecell.

12. The Nernst Equation

13. Relationship of E0cell to ΔG0 and K
Example 21-10: Calculate the standard Gibbs free energy change, ΔG0 , at 250C for the following reaction.

14. Relationship of E0cell to ΔG0 and K
Calculate E0cell using the appropriate half-reactions.

15. Relationship of E0cell to ΔG0 and K
Now that we know E0cell , we can calculate ΔG0 .

16. Relationship of E0cell to ΔG0 and K
Example 21-11: Calculate the thermodynamic equilibrium constant for the reaction in example at 250C.

17. Relationship of E0cell to ΔG0 and K
Example 21-12: Calculate the Gibbs Free Energy change, G and the equilibrium constant at 250C for the following reaction with the indicated concentrations.

18. Relationship of E0cell to ΔG0 and K
Calculate the standard cell potential E0cell.

19. Relationship of E0cell to ΔG0 and K
Use the Nernst equation to calculate Ecell for the given concentrations.

20. Relationship of E0cell to ΔG0 and K

21. Relationship of E0cell to ΔG0 and K

22. Relationship of E0cell to ΔG0 and K
Ecell = V, compared to E0cell = V.
We can use this information to calculate ΔG.
The negative ΔG tells us that the reaction is spontaneous.

23. Relationship of E0cell to ΔG0 and K
Equilibrium constants do not change with reactant concentration.
We can use the value of E0cell at 250C to get K.

24. Synthesis Question
What are the explosive chemicals in the fuel cell that exploded aboard Apollo 13?

25. Synthesis Question
The Apollo 13 fuel cells contained hydrogen and oxygen. Both are explosive, especially when mixed. The oxygen tank aboard Apollo 13 exploded.