1. Electrochemistry Applications of Redox AP Chemistry Chapter 20 Notes

2. Review  Oxidation reduction reactions involve a transfer of electrons.  OIL- RIG  Oxidation Involves Loss  Reduction Involves Gain  LEO-GER  Lose Electrons Oxidation  Gain Electrons Reduction

3. Applications  Moving electrons is electric current.  8H + +MnO 4 - + 5Fe +2 +5e -  Mn +2 + 5Fe +3 +4H 2 O  Helps to break the reactions into half rxns.  8H + +MnO 4 - +5e -  Mn +2 +4H 2 O  5Fe +2  5Fe +3 + 5e - )  In the same mixture it happens without doing useful work, but if separate

4. H + MnO 4 - Fe +2  Connected this way the reaction starts  Stops immediately because charge builds up.

5. H + MnO 4 - Fe +2 Galvanic Cell Salt Bridge allows current to flow

6. H + MnO 4 - Fe +2 e-e-  Electricity travels in a complete circuit  Instead of a salt bridge

7. H + MnO 4 - Fe +2 Porous Disk

8. Reducing Agent Oxidizing Agent e-e- e-e- e-e- e-e- e-e- e-e- AnodeCathode

9. Cell Potential  Oxidizing agent pushes the electron.  Reducing agent pulls the electron.  The push or pull (“driving force”) is called the cell potential E cell  Also called the electromotive force (emf)  Unit is the volt(V)  = 1 joule of work/coulomb of charge  Measured with a voltmeter

10. Zn +2 SO 4 -2 1 M HCl Anode 0.76 1 M ZnSO 4 H + Cl - H 2 in Cathode

11. 1 M HCl H + Cl - H 2 in Standard Hydrogen Electrode  This is the reference all other oxidations are compared to  E º = 0  º indicates standard states of 25ºC, 1 atm, 1 M solutions. 1 M solutions.

12. Cell Potential  Zn(s) + Cu +2 (aq)  Zn +2 (aq) + Cu(s)  The total cell potential is the sum of the potential at each electrode.  E º cell = E º Zn  Zn +2 + E º Cu +2  Cu  We can look up reduction potentials in a table.  One of the reactions must be reversed, so change it sign.

13. Cell Potential  Determine the cell potential for a galvanic cell based on the redox reaction.  Cu(s) + Fe +3 (aq)  Cu +2 (aq) + Fe +2 (aq)  Fe +3 (aq) + e -  Fe +2 (aq) E º = 0.77 V  Cu +2 (aq)+2e -  Cu(s) E º = 0.34 V  Cu(s)  Cu +2 (aq)+2e - E º = -0.34 V  2Fe +3 (aq) + 2e -  2Fe +2 (aq) E º = 0.77 V

14. Line Notation  solid  Aqueous  Aqueous  solid  Anode on the left  Cathode on the right  Single line different phases.  Double line porous disk or salt bridge.  If all the substances on one side are aqueous, a platinum electrode is indicated.  For the last reaction  Cu(s)  Cu +2 (aq)  Fe +2 (aq),Fe +3 (aq)  Pt(s)

15. Galvanic Cell  The reaction always runs spontaneously in the direction that produced a positive cell potential.  Four things for a complete description. 1)Cell Potential 2)Direction of flow 3)Designation of anode and cathode 4)Nature of all the components- electrodes and ions

16. Practice  Completely describe the galvanic cell based on the following half-reactions under standard conditions.  MnO 4 - + 8 H + +5e -  Mn +2 + 4H 2 O Eº=1.51  Fe +3 +3e -  Fe(s) Eº=0.036V

17. Potential, Work and  G  emf = potential (V) = work (J) / Charge(C)  E = work done by system / charge  E = -w/q  Charge is measured in coulombs.  -w = qE  Faraday = 96,485 C/mol e -  q = nF = moles of e - x charge/mole e -  w = -qE = -nFE =  G

18. Potential, Work and  G   Gº = -nFE º FIXED  if E º > 0, then  Gº 0, then  Gº < 0 spontaneous  if E º 0 nonspontaneous  In fact, reverse is spontaneous.  Calculate  Gº for the following reaction:  Cu +2 (aq)+ Fe(s)  Cu(s)+ Fe +2 (aq)  Fe +2 (aq) + e -  Fe(s) E º = 0.44 V  Cu +2 (aq)+2e -  Cu(s) E º = 0.34 V

19. Cell Potential and Concentration  Qualitatively - Can predict direction of change in E from LeChâtelier.  2Al(s) + 3Mn +2 (aq)  2Al +3 (aq) + 3Mn(s)  Predict if E cell will be greater or less than E º cell if [Al +3 ] = 1.5 M and [Mn +2 ] = 1.0 M  if [Al +3 ] = 1.0 M and [Mn +2 ] = 1.5M  if [Al +3 ] = 1.5 M and [Mn +2 ] = 1.5 M

20. The Nernst Equation   G =  Gº +RTln(Q)  -nFE = -nFEº + RTln(Q)  E = Eº - RT ln(Q) nF  2Al(s) + 3Mn +2 (aq)  2Al +3 (aq) + 3Mn(s) E º = 0.48 V  Always have to figure out n by balancing.  If concentration can gives voltage, then from voltage we can tell concentration.

21. The Nernst Equation  As reactions proceed concentrations of products increase and reactants decrease.  Reach equilibrium where Q = K and E cell = 0  0 = Eº - RT ln(K) nF  Eº = RT ln(K) nF  nFEº = ln(K) RT

22. Batteries are Galvanic Cells  Car batteries are lead storage batteries.  Pb +PbO 2 +H 2 SO 4  PbSO 4 (s) +H 2 O  Dry Cell Zn + NH 4 + +MnO 2  Zn +2 + NH 3 + H 2 O  Alkaline Zn +MnO 2  ZnO+ Mn 2 O 3 (in base)  NiCad  NiO 2 + Cd + 2H 2 O  Cd(OH) 2 +Ni(OH) 2

23. Corrosion  Rusting - spontaneous oxidation.  Most structural metals have reduction potentials that are less positive than O 2.  Fe  Fe +2 +2e - Eº= 0.44 V  O 2 + 2H 2 O + 4e -  4OH - Eº= 0.40 V  Fe +2 + O 2 + H 2 O  Fe 2 O 3 + H +  Reaction happens in two places.

24. Water Rust Iron Dissolves- Fe  Fe +2 e-e- Salt speeds up process by increasing conductivity

25. Preventing Corrosion  Coating to keep out air and water.  Galvanizing - Putting on a zinc coat  Has a lower reduction potential, so it is more. easily oxidized.  Alloying with metals that form oxide coats.  Cathodic Protection - Attaching large pieces of an active metal like magnesium that get oxidized instead.

26.  Running a galvanic cell backwards.  Put a voltage bigger than the potential and reverse the direction of the redox reaction.  Used for electroplating. Electrolysis

27. 1.0 M Zn +2 e-e- e-e- Anode Cathode 1.10 Zn Cu 1.0 M Cu +2

28. 1.0 M Zn +2 e-e- e-e- Anode Cathode A battery >1.10V Zn Cu 1.0 M Cu +2

29. Calculating plating  Have to count charge.  Measure current I (in amperes)  1 amp = 1 coulomb of charge per second  q = I x t  q/nF = moles of metal  Mass of plated metal  How long must 5.00 amp current be applied to produce 15.5 g of Ag from Ag +

30. Other uses  Electroysis of water.  Seperating mixtures of ions.  More positive reduction potential means the reaction proceeds forward.  We want the reverse.  Most negative reduction potential is easiest to plate out of solution.

31. Balancing Redox Equations

32. Redox Reactions Environment [acidic or basic] is very important Particles available but sometimes are not given in the reaction

33. acid hydrogen cation & water base hydroxide ion & water

34. Identify species undergoing 1. oxidation 2. reduction

35. Split overall reaction into 1. Oxidation half-reaction 2. Reduction half-reaction

36. Half-reaction in ACID Sol’n 1. Balance species changing oxidation # 2. Balance oxygen by adding H 2 O

37. 3. Balance hydrogen by adding H + 4. Balance charge by adding e - Half-reaction in ACID Sol’n

38. 5. Add the two half- reactions eliminating all electrons (least common multiple concept) Half-reaction in ACID Sol’n

39. CHECK final equation for BALANCE of ATOMS and CHARGE

40. The Redox Blues

41. 1. in acid AuCl 4 - (aq) + AsH 3 (g) ---> H 3 AsO 3 (aq) + Au(s) + Cl - (aq)

42. Half-reaction in BASIC Sol’n 1. Balance species changing oxidation # 2. Balance oxygen by adding twice as many OH-

43. Half-reaction in BASIC Sol’n 3. Balance hydrogen by adding H 2 O 4. Balance charge by adding e -

44. 5. Add the two half- reactions eliminating all electrons (least common multiple concept) Half-reaction in BASIC Sol’n

45. OR Balance by acidic method & then “neutralize” the H + by adding OH - and adjusting the H 2 O

46. 2. in base Am 3+ (aq) + S 2 O 8 2- (aq) ----> AmO 2 + (aq) + SO 4 2- (aq)

47. 3. MnO 4 - (aq) + H 2 C 2 O 4 (aq)  Mn 2+ (aq) + CO 2 (g)

48. 4. Bi(OH) 3 + SnO 2 2-  Bi(s) + SnO 3 2-

49. ELECTROLYTIC CELLS

50. Electrolytic Cell a cell that uses electrical energy to produce a chemical change that would otherwise NOT occur spontaneously

51. Process referred to as electrolysis

52. (+)(-) M M M + (aq) X - (aq)

53. (+)(-) e-e- e-e- M + (aq) X - (aq) MM Anode M M + + e - oxidation Cathode M + + e - M reduction

54. Ampere a unit of electrical current equal to one coulomb of charge per second 1 amp = 1 coul sec

55. Coulomb a unit of electric charge equal to the quantity of charge in about 6 x 10 19 electrons

56. Faraday a constant representing the charge on one mole of electrons 1 F = 96,485 C  96,500 C

57. 3: It is necessary to replate a silver teapot with 15.0 g of silver. If the electrolytic cell runs at 2.00 amps, how long will it take to plate the teapot?

58. 4: Sodium metal and chlorine gas are prepared industrially in a Down’s Cell from the electrolysis of molten NaCl. What mass of metal and volume of gas can be made per day if the cell operates at 7.0 volts and 4.0 x 10 4 amps if the cell is 75% efficient?

59. 5: At what current must a cell be run in order to produce 5.0 kg of aluminum in 8.0 hours if the cell produces solid aluminum from molten aluminum chloride?

60. ELECTROCHEMISTRY, FREE ENERGY, & EQUILIBRIUM

61. thus: w max = - q. E max

62. but: w max =  G and q = nF thus if: w max = - q. E max then  G = - nFE

63.  G =  G 0 + RT ln Q  G = - nFE - nFE = - nFE 0 + RT ln Q

64. NERNST EQUATION

65. if: aA + bB  cC + dD

66. IF T = 25 0 C = 298.15 K ln Q = 2.303 log Q R = 8.314 J/mol. K F = 96,485 C/mol

67. what if : Q = K eq ? then: E = 0.0 V

69. 6: Calculate the equilibrium constant at 40 0 C for the cell: Cd (s) Cd 2+ (1M) Pb 2+ (1M) Pb (s)

70. 7a: Calculate the standard free energy for the cell: Cr (s) Cr 3+ (1M) Fe 2+ (1M) Fe (s)

71. 7a: Calculate the standard free energy for the cell: Cr (s) Cr 3+ (1M) Fe 2+ (1M) Fe (s) 7b: What will be the voltage if [Fe 2+ ] = 0.50M and [Cr 3+ ] = 0.30M at 20 0 C?

72. 8: Through electrochemical calculations, determine the K sp for silver bromide. AgBr + e -  Ag + Br - E 0 = 0.10 V

73. Review of Redox & Electrochemical Cells Review of Redox & Electrochemical Cells

74. Review Oxidation: loss of e - [increase in ox #] [reducing agent] Reduction: gain of e - [decrease in ox #] [oxidizing agent]

75. Reduction Potential The ease with which a chemical species can be reduced

76. Standard Reduction Potential Appendix M Table 20.1 in text

78. 1. Which of the following elements listed is the best reducing agent? Cu Zn Fe Ag Cr

79. 2a.Choosing from among the reactants in the given half reactions, identify the strongest and weakest oxidizing agents.

80. Anode and Cathode  OXIDATION occurs at the ANODE.  REDuction occurs at the CAThode.

81. Electrochemical Cell device in which chemical energy is spontaneously changed to electrical energy

82. battery voltaic cell galvanic cell

83. An electrochemical cell consists of ???

84. M1M1 M2M2 M 1 + (aq) X - (aq) M 2 + (aq)

85. M1M1 M2M2 M 1 + (aq) X - (aq) M 2 + (aq) Anode M 1  M 1 + + e - Cathode M 2 + + e -  M 2

86. M1M1 M2M2 M 1 + (aq) X - (aq) M 2 + (aq) Anode M 1  M 1 + + e - Cathode M 2 + + e -  M 2 K + (aq) NO 3 - (aq)

87. e - flow is from source of high “concentration” to source of low “concentration”

88. M1M1 M2M2 M 1 + (aq) X - (aq) M 2 + (aq) Anode M 1  M 1 + + e - Cathode M 2 + + e -  M 2 K + (aq) NO 3 - (aq) e-e- e-e-

89. shorthand notation oxidation reduction M 1 | M 1 + || M 2 + | M 2 anode  cathode  e - flow 

90. this e - flow can accomplish work

91. Electrochemical Standard State Conditions [ions] = 1 M T = 25 0 C P gas = 1 atm

92. An electrochemical cell is spontaneous if: Oxidation-reduction occurs E red + E ox > 0

93. 2b.Which of the oxidizing agents listed is (are) capable of oxidizing Br - to BrO 3 - ?

95. Line Notation: ANODE CATHODE Ni(s)|Ni 2+ (aq, 1 M)||Au 3+ (aq, 1 M)|Au(s) oxidation reduction

96. Line Notation: ANODE CATHODE Al (s) | Al 3+ (aq, 1 M) || Ni 2+ (aq, 1 M) | Ni (s) oxidation reduction