1. 1 21 Electrochemistry

2. 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. 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. 4 Electrode Potentials for Other Half-Reactions Example 21-4: Will permanganate ions, MnO 4 -, 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. 5 Electrode Potentials for Other Half-Reactions Example 21-5: Will nitric acid, HNO 3, oxidize arsenous acid, H 3 AsO 3, in acidic solution? The reduction product of HNO 3 is NO in this reaction.

6. 6 The Nernst Equation Example 21-6: Calculate the potential for the Cu 2+ / Cu + electrode at 25 0 C when the concentration of Cu + ions is three times that of Cu 2+ ions.

7. 7 The Nernst Equation Example 21-7: Calculate the potential for the Cu 2+ /Cu + electrode at 25 0 C when the Cu + ion concentration is 1/3 of the Cu 2+ ion concentration.

8. 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 H 2 pressure is 0.50 atmosphere.

9. 9 The Nernst Equation Example 21-9: Calculate the initial potential of a cell that consists of an Fe 3+ /Fe 2+ electrode in which [Fe 3+ ]=1.0 x 10 -2 M and [Fe 2+ ]=0.1 M connected to a Sn 4+ /Sn 2+ electrode in which [Sn 4+ ]=1.0 M and [Sn 2+ ]=0.10 M. A wire and salt bridge complete the circuit.

10. 10 The Nernst Equation Calculate the E 0 cell by the usual procedure.

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

12. 12 The Nernst Equation

13. 13 Relationship of E 0 cell to  G 0 and K Example 21-10: Calculate the standard Gibbs free energy change,  G 0, at 25 0 C for the following reaction.

14. 14 Relationship of E 0 cell to  G 0 and K 1.Calculate E 0 cell using the appropriate half-reactions.

15. 15 Relationship of E 0 cell to  G 0 and K 2.Now that we know E 0 cell, we can calculate  G 0.

16. 16 Relationship of E 0 cell to  G 0 and K Example 21-11: Calculate the thermodynamic equilibrium constant for the reaction in example 21-10 at 25 0 C.

17. 17 Relationship of E 0 cell to  G 0 and K Example 21-12: Calculate the Gibbs Free Energy change,  G and the equilibrium constant at 25 0 C for the following reaction with the indicated concentrations.

18. 18 Relationship of E 0 cell to  G 0 and K 1.Calculate the standard cell potential E 0 cell.

19. 19 Relationship of E 0 cell to  G 0 and K 2.Use the Nernst equation to calculate E cell for the given concentrations.

20. 20 Relationship of E 0 cell to  G 0 and K

21. 21 Relationship of E 0 cell to  G 0 and K

22. 22 Relationship of E 0 cell to  G 0 and K E cell = +1.540 V, compared to E 0 cell = +1.562 V. We can use this information to calculate  G. The negative  G tells us that the reaction is spontaneous.

23. 23 Relationship of E 0 cell to  G 0 and K Equilibrium constants do not change with reactant concentration. We can use the value of E 0 cell at 25 0 C to get K.

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

25. 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.