1. Chapter 12 Oxidation-Reduction Reactions
In chapter 5, we learned how to recognize an oxidation-reduction reaction. In this chapter we introduce electrode potentials and the Nernst equation. These allow for a quantitative treatment of electrochemistry.

2. Common Oxidation-Reduction Reactions
Oxidation-reduction reactions used for heat or work:
Combustion
Metabolic
Corrosion
Photosynthesis

3. Common Oxidation-Reduction Reactions
Oxidation-reduction reactions involve the transfer of electrons.
Element or compound that gains electrons undergoes reduction.
Element or compound that loses electrons undergoes oxidation.

4. Common Oxidation-Reduction Reactions
Consider this reaction:
2Na + Cl2 → 2NaCl
The Na has been oxidized.
The Cl2 has been reduced.

5. Common Oxidation-Reduction Reactions
This reaction, 2Na + Cl2 → 2NaCl, can be written as the sum of two half-reactions:
2Na ⇄ 2Na+ + 2e- oxidation
Cl2 + 2e- ⇄ 2Cl- reduction

6. Common Oxidation-Reduction Reactions
The addition of oxygen atoms or hydrogen atoms to an element or compound is also classified as an oxidation-reduction reaction.

7. Common Oxidation-Reduction Reactions
CO2 + H2 ⇄ CO + H2O
The H2 is oxidized and the CO2 reduced.
C2H4 + H2 ⇄ C2H6
The C2H4 is reduced and the H2 oxidized.

8. Common Oxidation-Reduction Reactions
In order to determine that an oxidation-reduction reaction has occurred, we must be able to assign oxidation numbers or oxidation states.

9. Determining Oxidation Numbers
Introduced in Chapter 5, section 16
Two methods:
One based on Lewis structure
Good for organic compounds
Other based on set of rules
e. g., elements = 0, monatomic ions = charge
Review section 16, chapter 5

10. Recognizing Oxidation-Reduction Reactions
After all the oxidation numbers in a chemical reaction have been determined, look for changes.
Oxidation occurs when the oxidation number of an atom increases.
Reduction occurs when the oxidation number of an atom decreases.

11. Recognizing Oxidation-Reduction Reactions
In an oxidation-reduction reaction, both oxidation and reduction must occur.
If one species is being oxidized, another must be reduced.
In biochemical reactions, often only the oxidation or reduction reaction is shown explicitly.

12. Recognizing Oxidation-Reduction Reactions
Figure 12.1

13. Recognizing Oxidation-Reduction Reactions
Organic reactions can be classified by examining the Lewis structures.
If the number of C-H bonds decreases, the molecule is being oxidized.
If the number of C-O bonds increases, the molecule is being oxidized.
Conversely, if the number of C-H bonds increases, the molecule is being reduced.

14. Voltaic Cells Also known as galvanic cells
Physically separate the half-reactions
Force electrons to travel through an external circuit connecting the two half-reactions
Battery!

15. Voltaic Cells
Figure 12.2

16. Voltaic Cells
As H+ ions leave the solution on the right, K+ ions fill in to keep the solution electrically neutral.
Salt Bridge
The voltage required to prevent the flow of electrons is measured with a voltmeter.
This voltage is called the cell potential.

17. Voltaic Cells Oxidation takes place at the anode.
Reduction takes place at the cathode.
In Figure 12.2, the half-reactions involve two electrons.
The half-reactions are added to produce the overall reaction.

18. Voltaic Cells
What if the half reactions do not have the same number of electrons?
Figure 12.3

19. Voltaic Cells

20. Standard Cell Potentials
The relative half-reactions from Figure 12.2 are
Zn ⇄ Zn+2 +2e- E° =
2H+ + 2e- ⇄ H2 E° =
Figure 12.2

21. Standard Cell Potentials
The overall standard cell potential, E°, for the cell is the sum of the two half-reaction E°.
Expect the reaction to go as written if the overall E° >0.

22. Oxidizing and Reducing Agents
Reducing agent donates electrons: its oxidation number increases.
Oxidizing agent accepts electrons: its oxidation number decreases.

23. Oxidizing and Reducing Agents
In Figure 12.2, zinc metal is the reducing agent.
Hydrogen ions are the oxidizing agent.
Figure 12.2

24. Oxidizing and Reducing Agents
As with acids and bases, there are conjugate oxidizing and reducing agents.
When Zn is oxidized to Zn+2, Zn+2 becomes the conjugate oxidizing agent because its oxidation number drops in the reverse reaction:
Zn ⇄ Zn+2 + 2e-

25. Oxidizing and Reducing Agents
Strong reducing agents produce weak conjugate oxidizing agents.
Strong oxidizing agents produce weak conjugate reducing agents.

26. Relative Strengths of Oxidizing Agents and Reducing Agents
Oxidation-reduction reactions should occur when they convert the stronger of a pair of oxidizing agents and the stronger of a pair of reducing agents into a weaker oxidizing agent and a weaker reducing agent.

27. Relative Strengths of Oxidizing Agents and Reducing Agents
Table 12.1

28. Relative Strengths of Oxidizing Agents and Reducing Agents
Standard electrode potentials, E°red
Half-reactions written as reductions
Standard means gases at 1 bar, solutions at 1 M
When written as oxidations, the sign on E°red is reversed.

29. Batteries Alkaline dry cells, ubiquitous Lead-Acid, cars
NiCd, rechargeable
NiMH, hybrids
Lithium ion, compact
Fuel cells, hydrogen

30. Batteries
Lead-Acid

31. Batteries
NiMH

32. Electrochemical Cells at Nonstandard Conditions: The Nernst Equation
To determine E when a cell is not at standard conditions, the Nernst Equation is used.

33. Electrochemical Cells at Nonstandard Conditions: The Nernst Equation
n is the number of electrons transferred.
Qc is the reaction quotient.
Notice that if all concentrations are 1 M, E=E°.
F is the Faraday constant.

34. Electrochemical Cells at Nonstandard Conditions: The Nernst Equation
Zn(s) + Cu+2(aq) → Zn+2(aq) + Cu(s)
Figure 12.8

35. Electrochemical Cells at Nonstandard Conditions: The Nernst Equation
At equilibrium Qc = K and E = 0.
This provides an alternate equation for expressing equilibrium.

36. Electrolysis and Faraday’s Law
Voltaic cells operate spontaneously.
Electrolytic cells require an external power supply.
e. g. electroplating

37. Electrolysis and Faraday’s Law
Figure 12.9

38. Electrolysis and Faraday’s Law
The amount of a substance consumed or produced at one of the electrodes in an electrolytic cell is directly proportional to the amount of electricity that passes through the cell.

39. Electrolysis and Faraday’s Law
Amps × time (in secs) = Coulombs, C
F = 96,485 C/mol of e-
C/F = mol of e- passed
Grams of silver plated out can be determined from
[Ag(CN)2]-(aq) + e- ⇄ Ag(s) + 2CN-(aq)

40. Electrolysis of Molten NaCl
Figure 12.11

41. Electrolysis of Molten NaCl
CaCl2 added to the NaCl to lower the melting point. No effect on half reactions.
Na(l) less dense than NaCl(l).
Cl2(g) and Na(l) kept apart. Why?

42. Electrolysis of Aqueous NaCl
Figure 12.13

43. Electrolysis of Aqueous NaCl
Chloride is oxidized instead of water.
Water is reduced, not sodium ion.
Hydrogen gas and NaOH(aq) are produced and sold.

44. Electrolysis of Water
Figure 12.15

45. Electrolysis of Water
A salt which resists electrolysis is added to improve conductivity.
Similar (but not exactly) half reactions running in reverse describe a fuel cell.
If the gases were collected, what would their volume ratio be?

46. The Hydrogen Economy Using hydrogen gas as a common fuel.
Solar energy for electrolysis of water.
Fuel cells to generate electricity from hydrogen and reproduce the water.
No drain on fossil fuels.
No carbon emissions.
Water (seawater) already has the electrolyte for improved conductivity added!
Plenty of solar radiation.

47. The Hydrogen Economy Not a new idea. Big challenges.
Economical production of hydrogen.
Storage.
Distribution.
Better fuels cells.
Cheaper.
More robust.