1. 1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 20 Principles of Reactivity: Electron Transfer Reactions © 2006 Brooks/Cole Thomson Lectures written by John Kotz

2. 2 © 2006 Brooks/Cole - Thomson ELECTROCHEMISTRY Chapter 19

3. 3 © 2006 Brooks/Cole - Thomson TRANSFER REACTIONS Atom/Group transfer HCl + H 2 O ---> Cl - + H 3 O + Electron transfer Cu(s) + 2 Ag + (aq) ---> Cu 2+ (aq) + 2 Ag(s)

4. 4 © 2006 Brooks/Cole - Thomson Electron Transfer Reactions Electron transfer reactions are oxidation- reduction or redox reactions.Electron transfer reactions are oxidation- reduction or redox reactions. Redox reactions can result in the generation of an electric current or be caused by imposing an electric current.Redox reactions can result in the generation of an electric current or be caused by imposing an electric current. Therefore, this field of chemistry is often called ELECTROCHEMISTRY.Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

5. 5 © 2006 Brooks/Cole - Thomson Review of Terminology for Redox Reactions OXIDATION—loss of electron(s) by a species; increase in oxidation number.OXIDATION—loss of electron(s) by a species; increase in oxidation number. REDUCTION—gain of electron(s); decrease in oxidation number.REDUCTION—gain of electron(s); decrease in oxidation number. OXIDIZING AGENT—electron acceptor; species is reduced.OXIDIZING AGENT—electron acceptor; species is reduced. REDUCING AGENT—electron donor; species is oxidized.REDUCING AGENT—electron donor; species is oxidized. OXIDATION—loss of electron(s) by a species; increase in oxidation number.OXIDATION—loss of electron(s) by a species; increase in oxidation number. REDUCTION—gain of electron(s); decrease in oxidation number.REDUCTION—gain of electron(s); decrease in oxidation number. OXIDIZING AGENT—electron acceptor; species is reduced.OXIDIZING AGENT—electron acceptor; species is reduced. REDUCING AGENT—electron donor; species is oxidized.REDUCING AGENT—electron donor; species is oxidized.

6. 6 © 2006 Brooks/Cole - Thomson OXIDATION-REDUCTION REACTIONS Direct Redox Reaction Oxidizing and reducing agents in direct contact. Cu(s) + 2 Ag + (aq) ---> Cu 2+ (aq) + 2 Ag(s)

7. 7 © 2006 Brooks/Cole - Thomson OXIDATION-REDUCTION REACTIONS Indirect Redox Reaction A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent.

8. 8 © 2006 Brooks/Cole - Thomson Why Study Electrochemistry? BatteriesBatteries CorrosionCorrosion Industrial production of chemicals such as Cl 2, NaOH, F 2 and AlIndustrial production of chemicals such as Cl 2, NaOH, F 2 and Al Biological redox reactionsBiological redox reactions The heme group

9. 9 © 2006 Brooks/Cole - Thomson Electrochemical Cells An apparatus that allows a redox reaction to occur by transferring electrons through an external connector.An apparatus that allows a redox reaction to occur by transferring electrons through an external connector. Product favored reaction - voltaic or galvanic cell generates electric currentProduct favored reaction - voltaic or galvanic cell generates electric current Reactant favored reaction - electrolytic cell in which an electric current is used to cause chemical change.Reactant favored reaction - electrolytic cell in which an electric current is used to cause chemical change. Batteries are voltaic cells

10. 10 © 2006 Brooks/Cole - Thomson Electrochemistry Alessandro Volta, 1745-1827, Italian scientist and inventor. Luigi Galvani, 1737-1798, Italian scientist and inventor.

11. 11 © 2006 Brooks/Cole - Thomson Balancing Equations for Redox Reactions Some redox reactions have equations that must be balanced by special techniques. MnO 4 - + 5 Fe 2+ + 8 H + ---> Mn 2+ + 5 Fe 3+ + 4 H 2 O Mn = +7 Fe = +2 Fe = +3 Mn = +2

12. 12 © 2006 Brooks/Cole - Thomson Balancing Equations Cu + Ag + --give--> Cu 2+ + Ag

13. 13 © 2006 Brooks/Cole - Thomson Balancing Equations Step 1:Divide the reaction into half- reactions, one for oxidation and the other for reduction. OxCu ---> Cu 2+ Red Ag + ---> Ag Step 2:Balance each for mass. Already done in this case. Step 3:Balance each half-reaction for charge by adding electrons. Ox Cu ---> Cu 2+ + 2e- Red Ag + + e- ---> Ag

14. 14 © 2006 Brooks/Cole - Thomson Balancing Equations Step 4:Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires. Reducing agent Cu ---> Cu 2+ + 2e- Oxidizing agent 2 Ag + + 2 e- ---> 2 Ag Step 5:Add half-reactions to give the overall equation. Cu + 2 Ag + ---> Cu 2+ + 2Ag The equation is now balanced for both charge and mass.

15. 15 © 2006 Brooks/Cole - Thomson Reduction of VO 2 + with Zn

16. 16 © 2006 Brooks/Cole - Thomson Balancing Equations Balance the following in acid solution— VO 2 + + Zn ---> VO 2+ + Zn 2+ VO 2 + + Zn ---> VO 2+ + Zn 2+ Step 1:Write the half-reactions OxZn ---> Zn 2+ RedVO 2 + ---> VO 2+ Step 2:Balance each half-reaction for mass. OxZn ---> Zn 2+ Red VO 2 + ---> VO 2+ + H 2 O 2 H + + Add H 2 O on O-deficient side and add H + on other side for H-balance.

17. 17 © 2006 Brooks/Cole - Thomson Balancing Equations Step 3:Balance half-reactions for charge. OxZn ---> Zn 2+ + 2e- Rede- + 2 H + + VO 2 + ---> VO 2+ + H 2 O Step 4:Multiply by an appropriate factor. OxZn ---> Zn 2+ + 2e- Red 2e- + 4 H + + 2 VO 2 + ---> 2 VO 2+ + 2 H 2 O Step 5:Add balanced half-reactions Zn + 4 H + + 2 VO 2 + ---> Zn 2+ + 2 VO 2+ + 2 H 2 O

18. 18 © 2006 Brooks/Cole - Thomson Tips on Balancing Equations Never add O 2, O atoms, or O 2- to balance oxygen ONLY add H 2 O or OH -.Never add O 2, O atoms, or O 2- to balance oxygen ONLY add H 2 O or OH -. Never add H 2 or H atoms to balance hydrogen ONLY add H + or H 2 O.Never add H 2 or H atoms to balance hydrogen ONLY add H + or H 2 O. Be sure to write the correct charges on all the ions.Be sure to write the correct charges on all the ions. Check your work at the end to make sure mass and charge are balanced.Check your work at the end to make sure mass and charge are balanced. PRACTICE!PRACTICE!

19. 19 © 2006 Brooks/Cole - Thomson Oxidation: Zn(s) ---> Zn 2+ (aq) + 2e- Reduction: Cu 2+ (aq) + 2e- ---> Cu(s) -------------------------------------------------------- Cu 2+ (aq) + Zn(s) ---> Zn 2+ (aq) + Cu(s) Electrons are transferred from Zn to Cu 2+, but there is no useful electric current. CHEMICAL CHANGE ---> ELECTRIC CURRENT With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.”

20. 20 © 2006 Brooks/Cole - Thomson To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs through an external wire.To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs through an external wire. CHEMICAL CHANGE ---> ELECTRIC CURRENT This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery.

21. 21 © 2006 Brooks/Cole - Thomson Electrons travel through external wire. Salt bridge allows anions and cations to move between electrode compartments.Salt bridge allows anions and cations to move between electrode compartments. Electrons travel through external wire. Salt bridge allows anions and cations to move between electrode compartments.Salt bridge allows anions and cations to move between electrode compartments. Zn --> Zn 2+ + 2e- Cu 2+ + 2e- --> Cu <--AnionsCations--> OxidationAnodeNegativeOxidationAnodeNegativeReductionCathodePositiveReductionCathodePositive

22. 22 © 2006 Brooks/Cole - Thomson The Cu|Cu 2+ and Ag|Ag + Cell

23. 23 © 2006 Brooks/Cole - Thomson Electrochemical Cell Electrons move from anode to cathode in the wire. Anions & cations move thru the salt bridge.

24. 24 © 2006 Brooks/Cole - Thomson Terms Used for Voltaic Cells Figure 20.6

25. 25 © 2006 Brooks/Cole - Thomson The Voltaic Pile Drawing done by Volta to show the arrangement of silver and zinc disks to generate an electric current. What voltage does a cell generate?

26. 26 © 2006 Brooks/Cole - Thomson CELL POTENTIAL, E Electrons are “driven” from anode to cathode by an electromotive force or emf.Electrons are “driven” from anode to cathode by an electromotive force or emf. For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M.For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M. Zn and Zn 2+, anode Cu and Cu 2+, cathode 1.10 V 1.0 M

27. 27 © 2006 Brooks/Cole - Thomson CELL POTENTIAL, E For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M.For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M. This is the STANDARD CELL POTENTIAL, E oThis is the STANDARD CELL POTENTIAL, E o —a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.—a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.

28. 28 © 2006 Brooks/Cole - Thomson Calculating Cell Voltage Balanced half-reactions can be added together to get overall, balanced equation.Balanced half-reactions can be added together to get overall, balanced equation. Zn(s) ---> Zn 2+ (aq) + 2e- Cu 2+ (aq) + 2e- ---> Cu(s) -------------------------------------------- Cu 2+ (aq) + Zn(s) ---> Zn 2+ (aq) + Cu(s) Zn(s) ---> Zn 2+ (aq) + 2e- Cu 2+ (aq) + 2e- ---> Cu(s) -------------------------------------------- Cu 2+ (aq) + Zn(s) ---> Zn 2+ (aq) + Cu(s) If we know E o for each half-reaction, we could get E o for net reaction.

29. 29 © 2006 Brooks/Cole - Thomson CELL POTENTIALS, E o STANDARD HYDROGEN CELL Can’t measure 1/2 reaction E o directly. Therefore, measure it relative to a STANDARD HYDROGEN CELL 2 H + (aq, 1 M) + 2e- H 2 (g, 1 atm) E o = 0.0 V

30. 30 © 2006 Brooks/Cole - Thomson Zn/Zn 2+ half-cell hooked to a SHE. E o for the cell = +0.76 V Zn/Zn 2+ half-cell hooked to a SHE. E o for the cell = +0.76 V Negative electrode Supplier of electrons Acceptor of electrons Positive electrode 2 H + + 2e- --> H 2 ReductionCathode Zn --> Zn 2+ + 2e- OxidationAnode

31. 31 © 2006 Brooks/Cole - Thomson Reduction of H + by Zn Active Figure 20.13

32. 32 © 2006 Brooks/Cole - Thomson Overall reaction is reduction of H + by Zn metal. Zn(s) + 2 H + (aq) --> Zn 2+ + H 2 (g) E o = +0.76 V Therefore, E o for Zn ---> Zn 2+ (aq) + 2e- is +0.76 V Zn is a better reducing agent than H 2.

33. 33 © 2006 Brooks/Cole - Thomson Zn/Cu Electrochemical Cell Zn(s) ---> Zn 2+ (aq) + 2e-E o = +0.76 V Cu 2+ (aq) + 2e- ---> Cu(s)E o = +0.34 V --------------------------------------------------------------- Cu 2+ (aq) + Zn(s) ---> Zn 2+ (aq) + Cu(s) E o (calc’d) = +1.10 V Cathode, positive, sink for electrons Anode, negative, source of electrons +

34. 34 © 2006 Brooks/Cole - Thomson Uses of E o Values Organize half-reactions by relative ability to act as oxidizing agents Use this to predict direction of redox reactions and cell potentials.Use this to predict direction of redox reactions and cell potentials. Cu 2+ (aq) + 2e- ---> Cu(s)E o = +0.34 V Zn 2+ (aq) + 2e- ---> Zn(s) E o = –0.76 V Note that when a reaction is reversed the sign of E˚ is reversed!

35. 35 © 2006 Brooks/Cole - Thomson

36. 36 Potential Ladder for Reduction Half-Reactions Best oxidizing agents Best reducing agents Figure 20.14

37. 37 © 2006 Brooks/Cole - Thomson TABLE OF STANDARD REDUCTION POTENTIALS 2 E o (V) Cu 2+ + 2e- Cu+0.34 2 H + + 2e- H0.00 Zn 2+ + 2e- Zn-0.76 oxidizing ability of ion reducing ability of element

38. 38 © 2006 Brooks/Cole - Thomson Using Standard Potentials, E o Table 20.1 Which is the best oxidizing agent: O 2, H 2 O 2, or Cl 2 ? _________________ Which is the best reducing agent: Hg, Al, or Sn? ____________________

39. 39 © 2006 Brooks/Cole - Thomson Standard Redox Potentials, E o Any substance on the right will reduce any substance higher than it on the left. Zn can reduce H + and Cu 2+.Zn can reduce H + and Cu 2+. H 2 can reduce Cu 2+ but not Zn 2+H 2 can reduce Cu 2+ but not Zn 2+ Cu cannot reduce H + or Zn 2+.Cu cannot reduce H + or Zn 2+.

40. 40 © 2006 Brooks/Cole - Thomson Standard Redox Potentials, E o Cu 2+ + 2e- --> Cu +0.34 + 2 H + 2e- --> H 2 0.00 Zn 2+ + 2e- --> Zn -0.76 Northwest-southeast rule: product-favored reactions occur between reducing agent at southeast corner reducing agent at southeast corner oxidizing agent at northwest corner oxidizing agent at northwest corner Any substance on the right will reduce any substance higher than it on the left. Ox. agent Red. agent

41. 41 © 2006 Brooks/Cole - Thomson Cu(s) | Cu 2+ (aq) || H + (aq) | H 2 (g) Cu 2+ + 2e- --> Cu Or Cu --> Cu 2+ + 2 e- H 2 --> 2 H + + 2 e- or 2 H + + 2e- --> H 2 CathodePositiveAnodeNegative Electrons<----------

42. 42 © 2006 Brooks/Cole - Thomson Cu(s) | Cu 2+ (aq) || H + (aq) | H 2 (g) Cu 2+ + 2e- --> Cu H 2 --> 2 H + + 2 e- CathodePositiveAnodeNegative Electrons<---------- The sign of the electrode in Table 20.1 is the polarity when hooked to the H + /H 2 half-cell.

43. 43 © 2006 Brooks/Cole - Thomson Using Standard Potentials, E o Table 20.1 In which direction do the following reactions go?In which direction do the following reactions go? Cu(s) + 2 Ag + (aq) ---> Cu 2+ (aq) + 2 Ag(s)Cu(s) + 2 Ag + (aq) ---> Cu 2+ (aq) + 2 Ag(s) –Goes right as written 2 Fe 2+ (aq) +2 Fe 2+ (aq) + Sn 2+ (aq) ---> 2 Fe 3+ (aq) + Sn(s) –Goes LEFT opposite to direction written What is E o net for the overall reaction?

44. 44 © 2006 Brooks/Cole - Thomson Cd --> Cd 2+ + 2e- or Cd 2+ + 2e- --> Cd Fe --> Fe 2+ + 2e- or Fe 2+ + 2e- --> Fe E o for a Voltaic Cell All ingredients are present. Which way does reaction proceed? Calculate E o for this cell.

45. 45 © 2006 Brooks/Cole - Thomson E at Nonstandard Conditions The NERNST EQUATIONThe NERNST EQUATION E = potential under nonstandard conditionsE = potential under nonstandard conditions n = no. of electrons exchangedn = no. of electrons exchanged F = Faraday’s constantF = Faraday’s constant R = gas constantR = gas constant T = temp in KelvinsT = temp in Kelvins ln = “natural log”ln = “natural log” Q = reaction quotientQ = reaction quotient

46. 46 © 2006 Brooks/Cole - Thomson BATTERIES Primary, Secondary, and Fuel Cells

47. 47 © 2006 Brooks/Cole - Thomson Dry Cell Battery Anode (-) Zn ---> Zn 2+ + 2e- Cathode (+) 2 NH 4 + + 2e- ---> 2 NH 3 + H 2 Primary battery — uses redox reactions that cannot be restored by recharge.

48. 48 © 2006 Brooks/Cole - Thomson Nearly same reactions as in common dry cell, but under basic conditions. Alkaline Battery Anode (-): Zn + 2 OH - ---> ZnO + H 2 O + 2e- Cathode (+): 2 MnO 2 + H 2 O + 2e- ---> Mn 2 O 3 + 2 OH -

49. 49 © 2006 Brooks/Cole - Thomson Lead Storage Battery Secondary batterySecondary battery Uses redox reactions that can be reversed.Uses redox reactions that can be reversed. Can be restored by rechargingCan be restored by recharging

50. 50 © 2006 Brooks/Cole - Thomson Ni-Cad Battery Anode (-) Cd + 2 OH - ---> Cd(OH) 2 + 2e- Cathode (+) NiO(OH) + H 2 O + e- ---> Ni(OH) 2 + OH -

51. 51 © 2006 Brooks/Cole - Thomson Fuel Cells: H 2 as a Fuel Fuel cell - reactants are supplied continuously from an external source.Fuel cell - reactants are supplied continuously from an external source. Cars can use electricity generated by H 2 /O 2 fuel cells.Cars can use electricity generated by H 2 /O 2 fuel cells. H 2 carried in tanks or generated from hydrocarbons.H 2 carried in tanks or generated from hydrocarbons.

52. 52 © 2006 Brooks/Cole - Thomson Hydrogen—Air Fuel Cell Figure 20.12

53. 53 © 2006 Brooks/Cole - Thomson H 2 as a Fuel Comparison of the volumes of substances required to store 4 kg of hydrogen relative to car size. (Energy, p. 290)

54. 54 © 2006 Brooks/Cole - Thomson Storing H 2 as a Fuel One way to store H 2 is to adsorb the gas onto a metal or metal alloy. (Energy, p. 290)

55. 55 © 2006 Brooks/Cole - Thomson Electrolysis Using electrical energy to produce chemical change. Sn 2+ (aq) + 2 Cl - (aq) ---> Sn(s) + Cl 2 (g)

56. 56 © 2006 Brooks/Cole - Thomson Electrolysis Electric Energy ---> Chemical Change Electrolysis of molten NaCl. Electrolysis of molten NaCl. Here a battery “pumps” electrons from Cl - to Na +. Here a battery “pumps” electrons from Cl - to Na +. NOTE: Polarity of electrodes is reversed from batteries. NOTE: Polarity of electrodes is reversed from batteries.

57. 57 © 2006 Brooks/Cole - Thomson Electrolysis of Molten NaCl Figure 20.18

58. 58 © 2006 Brooks/Cole - Thomson Electrolysis of Molten NaCl Anode (+) 2 Cl - ---> Cl 2 (g) + 2e- Cathode (-) Na + + e- ---> Na E o for cell (in water) = - 4.07 V (in water) External energy needed because E o is (-).

59. 59 © 2006 Brooks/Cole - Thomson Electrolysis of Aqueous NaI Anode (+): 2 I - ---> I 2 (g) + 2e- Cathode (-): 2 H 2 O + 2e- ---> H 2 + 2 OH - E o for cell = -1.36 V

60. 60 © 2006 Brooks/Cole - Thomson Electrolysis of Aqueous CuCl 2 Anode (+) 2 Cl - ---> Cl 2 (g) + 2e- Cathode (-) Cu 2+ + 2e- ---> Cu E o for cell = -1.02 V Note that Cu 2+ is more easily reduced than either H 2 O or Na +.

61. 61 © 2006 Brooks/Cole - Thomson Michael Faraday 1791-1867 Originated the terms anode, cathode, anion, cation, electrode. Discoverer of electrolysiselectrolysis magnetic props. of mattermagnetic props. of matter electromagnetic inductionelectromagnetic induction benzene and other organic chemicalsbenzene and other organic chemicals Was a popular lecturer.

62. 62 © 2006 Brooks/Cole - Thomson E o and Thermodynamics E o is related to ∆G o, the free energy change for the reaction.E o is related to ∆G o, the free energy change for the reaction. ∆G˚ is proportional to –nE˚∆G˚ is proportional to –nE˚ ∆G o = -nFE o where F = Faraday constant = 9.6485 x 10 4 J/Vmol of e - (or 9.6485 x 10 4 coulombs/mol) and n is the number of moles of electrons transferred

63. 63 © 2006 Brooks/Cole - Thomson E o and ∆G o ∆G o = - n F E o For a product-favored reaction Reactants ----> Products Reactants ----> Products ∆G o 0 E o is positive For a reactant-favored reaction Reactants <---- Products Reactants <---- Products ∆G o > 0 and so E o 0 and so E o < 0 E o is negative

64. 64 © 2006 Brooks/Cole - Thomson E o and Equilibrium Constant  G o = -RT ln K  G o = -n F E o

65. 65 © 2006 Brooks/Cole - Thomson Quantitative Aspects of Electrochemistry Consider electrolysis of aqueous silver ion. Ag + (aq) + e- ---> Ag(s) 1 mol e----> 1 mol Ag If we could measure the moles of e-, we could know the quantity of Ag formed. But how to measure moles of e-?

66. 66 © 2006 Brooks/Cole - Thomson But how is charge related to moles of electrons? Quantitative Aspects of Electrochemistry = 96,500 C/mol e- = 1 Faraday Michael Faraday 1791-1867

67. 67 © 2006 Brooks/Cole - Thomson Quantitative Aspects of Electrochemistry 1.50 amps flow through a Ag + (aq) solution for 15.0 min. What mass of Ag metal is deposited? Solution (a)Calc. charge Charge (C) = current (A) x time (t) = (1.5 amps)(15.0 min)(60 s/min) = 1350 C

68. 68 © 2006 Brooks/Cole - Thomson Quantitative Aspects of Electrochemistry Solution (a)Charge = 1350 C (b)Calculate moles of e- used 1.50 amps flow through a Ag + (aq) solution for 15.0 min. What mass of Ag metal is deposited? (c)Calc. quantity of Ag

69. 69 © 2006 Brooks/Cole - Thomson Quantitative Aspects of Electrochemistry The anode reaction in a lead storage battery is Pb(s) + HSO 4 - (aq) ---> PbSO 4 (s) + H + (aq) + 2e- If a battery delivers 1.50 amp, and you have 454 g of Pb, how long will the battery last? Solution a)454 g Pb = 2.19 mol Pb b)Calculate moles of e- c)Calculate charge 4.38 mol e- 96,500 C/mol e- = 423,000 C 4.38 mol e- 96,500 C/mol e- = 423,000 C

70. 70 © 2006 Brooks/Cole - Thomson Quantitative Aspects of Electrochemistry The anode reaction in a lead storage battery is Pb(s) + HSO 4 - (aq) ---> PbSO 4 (s) + H + (aq) + 2e- If a battery delivers 1.50 amp, and you have 454 g of Pb, how long will the battery last? Solution a)454 g Pb = 2.19 mol Pb b)Mol of e- = 4.38 mol c)Charge = 423,000 C About 78 hours d)Calculate time