1. Electrochemistry Chapter 20

2. oxidation: lose e- -increase oxidation number reduction: gain e- -reduces oxidation number LEO goes GER Oxidation-Reduction Review

3. Oxidation Numbers Review Review rules p. 132 1.HNO 3 6. Ag 10. CO 2 1.CuCl 2 7. PbSO 4 11. (NH 4 ) 2 Ce(SO 4 ) 3 1.O 2 8. PbO 2 12. Cr 2 O 3 1.H 2 O 2 9. Na 2 C 2 O 4 1.MgSO 4

4. Oxidation-Reduction Review 1.CH 4 (g) + H 2 O(g) → CO(g) + 3H 2 (g) 1.2AgNO 3 (aq) + Cu(s) → Cu(NO 3 ) 2 (aq) + 2Ag(s) 1.Zn(s) + 2HCl(aq) → ZnCl 2 (aq) + H 2 (g) 1.2H + (aq) + 2CrO 4 2- (aq) → Cr 2 O 7 2- (aq) + H 2 O(l)

5. Balancing Redox Reactions Half Reactions: show only the oxidation or only the reduction part of the reaction Sn 2+ (aq) + 2Fe 3+ (aq) → Sn 4+ (aq) + 2Fe 2+ (aq) ox.: Sn 2+ (aq) → Sn 4+ (aq) + 2e - red.: 2e - + 2Fe 3+ (aq) → 2Fe 2+ (aq)

6. Balancing Redox Reactions 1.split into the half-reactions 2.balance each half-reactions a.first balance elements other than H and O b.balance O atoms by adding H 2 O as needed c.balance H atoms by adding H + as needed d.balance charge by adding e - as needed 3.multiply half-reactions by as needed to make the # of e - lost in ox. equal to # of e - gained in red. 4.add half reactions, simplify by canceling 5.double check everything

7. Balancing Redox Reactions What is reduced? What is oxidized? Reducing agent? Oxidizing agent?

8. Balancing Redox Reactions MnO 4 - (aq) + C 2 O 4 2- (aq) → Mn 2+ (aq) + CO 2 (aq)

9. Balancing Redox Reactions Do free electrons appear anywhere in the balanced equation for a redox reaction? Explain

10. Balancing Redox Reactions Cr 2 O 7 2- (aq) + Cl - (aq) → Cr 3+ (aq) + Cl 2 (g) (acidic solution)

11. Balancing Redox Reactions Cu(s) + NO 3 - (aq) → Cu 2+ (aq) + NO 2 (aq) (acidic)

12. Balancing Redox Reactions Mn 2+ (aq) + NaBiO 3 (s) → Bi 3+ (aq) + MnO 4 - (aq) (acidic)

13. Balancing Redox Reactions CN - (aq) + MnO 4 - (aq) → CNO - (aq) + MnO 2 (s) (basic)

14. Balancing Redox Reactions NO 2 - (aq) + Al(s) → NH 3 (aq) + Al(OH) 4 - (aq) (basic)

15. Balancing Redox Reactions Cr(OH) 3 (s) + ClO - (aq) → CrO 4 2- (aq) + Cl 2 (g)

16. Electrochemistry study of the interchange of chemical and electrical energy

17. Galvanic or Voltaic Cells -Device in which chemical energy is changed to electrical energy -Uses a spontaneous redox reaction to produce a current that can be used to do work -oxidizing agent is separated from the reducing agent and the electrons are forced to transfer through a wire -the current produced through the wire can then be used to do work

18. Galvanic or Voltaic Cells

19. Salt Bridge: allows ions to flow so the net charge is zero -without the salt bridge, the current would stop flowing due to a charge buildup -negative at red. side and positive at ox. side

20. Galvanic or Voltaic Cells Anode: where oxidation occurs Cathode: where reduction occurs AN OX, Red Cat

21. Why does Na + migrate into the cathode half-cell as the cell reaction proceeds?

22. Cell Potential Cell Potential: (E° cell ) -the oxidizing agent pulls electrons through the wire from the reducing agent. aka. electromotive force -unit is the volt: 1V = 1J of work per coulomb of charge (J/coulomb) -measured with a voltmeter: -drawing of current through a known resistance -or a potentiometer: measures opposition to current (compares to a known emf)

23. Standard Reduction Potentials All galvanic cell reactions are redox reactions that can be broken down into 2 half reactions -can assign a potential to each half reaction and calculate the cell potential by adding the potentials of the two half reactions The standard: -platinum electrode in contact with 1M H+ and bathed by H2(g) at 1atm has a reduction potential of 0V → called a standard hydrogen electrode 2H + + 2e - → H 2 E° = 0v

24. Standard Hydrogen Electrode

25. Standard Reduction Potentials -Compare all half reactions to the standard hydrogen electrode to obtain the reduction potential of the half-reaction E° means standard state: 1M, 1atm -Potentials of half reactions are given as reduction potentials - pink sheet -oxidation potentials are the reverse E° is an intensive property: does not change when multiplied E° cell = E° cathode - E° anode

26. SRP Calculation

29. Calculating Cell Potential Voltaic cell based on the reaction: Fe 3+ (aq) + Cu(s) → Cu 2+ + Fe 2+

30. Calculating Cell Potential Mg anode and Sn cathode

31. Calculating Cell Potential Cu cathode and Zn anode

32. Calculating Cell Potential Cr anode and Ag cathode

33. Strength of Oxidizing & Reducing Agents more negative E = most likely to be reversed and run as an oxidation reaction -Li + : E = -3.05: poor oxidizing agent -difficult to reduce (gain e-) - it is a very good reducing agent: will easily lose e- Would the halogens be good reducing agents or good oxidizing agents? Why

34. SRP Table In the reduction half reaction: the reactants are oxidizing agents the products are reducing agents

35. SRP Table

36. Relative Strengths of Oxidizing Agents Rank the following ions in order of increasing strength as oxidizing agents: NO 3 -, Ag +, Cr 2 O 7 2- Rank the following ions in order of decreasing strength as reducing agents: I - (aq), Fe(s), Al(s)

37. Determining Spontaneity +E = spontaneous -E = nonspontaneous Are the following reactions spontaneous? 1.Cu(s) + 2H + (aq) → Cu 2+ (aq) + H 2 (g) 2. Cl 2 (g) + 2I - (aq) → 2Cl - (aq) + I 2 (s)

38. Free Energy and Redox Gibbs Free Energy: measure of the spontaneity of a process at standard conditions -EMF (electromotive force) also indicates spontaneity ΔG = -nFE n = moles of electrons transferred in balanced equation F = Faraday’s constant = 96485C/mol or 96485 J/Vmol Units of ΔG: J/mol -means per mole of reaction

39. Free Energy and Redox does it make sense? ΔG = -nFE What sign of ΔG indicates a spontaneous reaction? What about E?

40. Free Energy and Redox Calculate ΔG and K at 298K for: 4Ag(s) + O 2 (g) + 4H + (aq) → 4Ag + (aq) + 2H 2 O

41. Free Energy and Redox Would the answer be different if the reaction was written: 2Ag(s) + ½ O 2 (g) + 2H + (aq) → 2Ag + (aq) + H 2 O(l)

42. Nonstandard Conditions Use LeChatelier’s Principle to predict the direction the reaction will shift (there is something called the Nernst equation, but that has been removed from the AP curriculum - you will probably come across it at some point)

43. Cell Potential and Concentration Standard conditions: 1M concentration Given: 2Al(s) + 3Mn 2+ (aq) 2Al 3+ (aq) + 3Mn(s); E° cell = 0.48V Change concentrations: 1.[Al 3+ ] = 2.0M and [Mn 2+ ] = 1.0M 2. [Al 3+ ] = 1.0M and [Mn 2+ ] = 3.0M

44. Electrolysis non spontaneous redox reaction -requires an outside source of electrical energy electrolytic cell - 2 electrodes in molten salt or a solution -the electrodes are inert (don’t react, just allow the reaction to occur

45. Electrolysis Electroplating -the cathode is active -metal deposits on the cathode Ni 2+ + 2e - → Ni(s) -2 moles of e- needed to plate 1 mole of Ni from Ni 2+

46. Electrolysis 1 mole of electrons has a charge of 96,485C (C for coulombs) coulomb is the amount of charge Current is the charge per second = ampere I = q/t I = current (ampere) q = charge (coulombs) t = time (seconds)

47. Electrolysis Calculate the number of grams of aluminum produced in 1.00h by the electrolysis of molten AlCl 3 if the electrical current is 10.0A

48. Electrolysis The half-reaction for formation of magnesium metal upon electrolysis of molten MgCl 2 is Mg 2+ + 2e - → Mg. 1.Calculate the mass of magnesium formed upon passage of a current of 60.0A for a period of 4.00x10 3 s. b. How many seconds would be required to produce 50.0g of Mg from MgCl 2 if the current is 100.0A?