Chapter 12 Oxidation-Reduction Reactions

1. 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. Chapter 12Oxidation-Reduction Reactions

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° =+0.7628 2H+ + 2e-⇄ H2 E° =+0.0000 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.