1. AP Chemistry
Chapter 20 Notes
Electrochemistry
Applications of Redox

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++MnO4-+ 5Fe+2 +5e- ® Mn+2 + 5Fe+3 +4H2O
Helps to break the reactions into half rxns.
8H++MnO4-+5e- ® Mn+2 +4H2O
5Fe+2 ® 5Fe+3 + 5e- )
In the same mixture it happens without doing useful work, but if separate

4. Connected this way the reaction starts
Stops immediately because charge builds up. H+
MnO4-
Fe+2

5. Galvanic Cell
Salt Bridge allows current to flow H+
MnO4-
Fe+2

6. Electricity travels in a complete circuit
Instead of a salt bridge e- H+
MnO4-
Fe+2

7. Porous Disk H+
MnO4-
Fe+2

8. e- e- e- e-
Anode
Cathode e- e-
Reducing Agent
Oxidizing Agent

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

10. 0.76
H2 in
Cathode
Anode
H+ Cl-
Zn+2 SO4-2
1 M ZnSO4
1 M HCl

11. Standard Hydrogen Electrode
This is the reference all other oxidations are compared to
Eº = 0
º indicates standard states of 25ºC, 1 atm,
1 M solutions.
H2 in
H+ Cl-
1 M HCl

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.
Eº cell = EºZn® Zn+2 + Eº 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) Eº = 0.77 V
Cu+2(aq)+2e- ® Cu(s) Eº = 0.34 V
Cu(s) ® Cu+2(aq)+2e Eº = V
2Fe+3(aq) + 2e-® 2Fe+2(aq) Eº = 0.77 V

14. Line Notation solid½Aqueous½½Aqueous½solid
Anode on the left½½Cathode 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)½Cu+2(aq)½½Fe+2(aq),Fe+3(aq)½Pt(s)

15. Galvanic Cell
The reaction always runs spontaneously in the direction that produced a positive cell potential.
Four things for a complete description.
Cell Potential
Direction of flow
Designation of anode and cathode
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 H+ +5e- ® Mn+2 + 4H2O Eº=1.51
Fe+3 +3e- ® Fe(s) Eº=0.036V

17. Potential, Work and DG 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 = DG

18. Potential, Work and DG DGº = -nFE º
if E º < 0, then DGº > 0 spontaneous
if E º > 0, then DGº < 0 nonspontaneous
In fact, reverse is spontaneous.
Calculate DGº for the following reaction:
Cu+2(aq)+ Fe(s) ® Cu(s)+ Fe+2(aq)
Fe+2(aq) + e-® Fe(s) Eº = 0.44 V
Cu+2(aq)+2e- ® Cu(s) Eº = 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 Ecell will be greater or less than Eºcell 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 E = Eº - RTln(Q) nF DG = DGº +RTln(Q)
-nFE = -nFEº + RTln(Q)
E = Eº - RTln(Q) nF
2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s) Eº = 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 0 = Eº - RTln(K) nF Eº = RTln(K) nF
As reactions proceed concentrations of products increase and reactants decrease.
Reach equilibrium where Q = K and Ecell = 0
0 = Eº - RTln(K) nF
Eº = RTln(K) nF
nFEº = ln(K) RT

22. Batteries are Galvanic Cells
Car batteries are lead storage batteries.
Pb +PbO2 +H2SO4 ®PbSO4(s) +H2O
Dry Cell Zn + NH4+ +MnO2 ® Zn+2 + NH3 + H2O
Alkaline Zn +MnO2 ® ZnO+ Mn2O3 (in base)
NiCad
NiO2 + Cd + 2H2O ® Cd(OH)2 +Ni(OH)2

23. Corrosion Rusting - spontaneous oxidation.
Most structural metals have reduction potentials that are less positive than O2 .
Fe ® Fe+2 +2e- Eº= 0.44 V
O2 + 2H2O + 4e- ® 4OH- Eº= 0.40 V
Fe+2 + O2 + H2O ® Fe2 O3 + H+
Reaction happens in two places.

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

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

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

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

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. hydrogen cation & water
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 H2O

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

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

39. CHECK
final equation
for BALANCE of
ATOMS and CHARGE

40. The Redox Blues

41. 1. in acid
AuCl4-(aq) + AsH3(g) --->
H3AsO3(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 H2O
4. Balance charge by adding e-

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

45. OR
Balance by acidic method & then “neutralize” the H+ by adding OH- and adjusting the H2O

46. 2. in base
Am3+(aq) + S2O82-(aq) ---->
AmO2+(aq) + SO42-(aq)

47. 3. MnO4-(aq) + H2C2O4(aq) 
Mn2+(aq) + CO2(g)

48. 4. Bi(OH)3 + SnO22- 
Bi(s) + SnO32-

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+(aq) M M
X-(aq)

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

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

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

56. a constant representing the charge on one mole of electrons
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
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 104 amps if the cell is 75% efficient?

59. 5: At what current must a cell be run in order to produce 5
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: wmax = - q . Emax

62. but: wmax = G
and q = nF
thus if: wmax = - q . Emax
then G = - nFE

63. G = G0 + RT ln Q
G = - nFE
- nFE = - nFE0 + RT ln Q

64. NERNST EQUATION

65. if: aA + bB  cC + dD

66. IF T = 250C = K
ln Q = log Q
R = J/mol.K
F = 96,485 C/mol

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

69. 6: Calculate the equilibrium constant at 400C for the cell:
Cd(s) Cd2+ (1M) Pb2+ (1M) Pb(s)

70. 7a: Calculate the standard free energy for the cell:
Cr(s) Cr3+ (1M) Fe2+ (1M) Fe(s)

71. 7a: Calculate the standard free energy for the cell:
Cr(s) Cr3+ (1M) Fe2+ (1M) Fe(s)
7b: What will be the voltage if [Fe2+] = 0.50M and [Cr3+] = 0.30M at 200C?

72. 8: Through electrochemical calculations, determine the Ksp for silver bromide.
AgBr + e-  Ag + Br-
E0 = 0.10 V

73. 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. The ease with which a chemical species can be reduced
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. M2+(aq)
M1+(aq) M1
X-(aq)
X-(aq) M2

85. Anode M1  M1+ + e- Cathode M2+ + e-  M2 M2+(aq) M1+(aq) M1 X-(aq)

86. Cathode Anode M2+ + e-  M2 M1  M1+ + e- M2+(aq) M1+(aq) M1 X-(aq)
K+(aq) NO3-(aq)
M2+(aq)
M1+(aq) M1
X-(aq)
X-(aq) M2
Cathode
M e-  M2
Anode
M1  M e-

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

88. e- e- Cathode Anode M2+ + e-  M2 M1  M1+ + e- M2+(aq) M1+(aq) M1
K+(aq) NO3-(aq)
M2+(aq)
M1+(aq) M1
X-(aq)
X-(aq) M2
Cathode
M e-  M2
Anode
M1  M e-

89. shorthand notation
oxidation reduction
M1 | M1+ || M2+ | M2
anode  cathode
 e- flow 

90. this e- flow can accomplish work

91. Electrochemical Standard State Conditions
[ions] = 1 M
T = 250C
Pgas = 1 atm

92. An electrochemical cell is spontaneous if: Oxidation-reduction occurs
Ered + Eox > 0

93. 2b. Which of the oxidizing agents listed is (are) capable of oxidizing Br- to BrO3- ?

95. Line Notation: ANODE CATHODE Ni(s)|Ni2+ (aq, 1 M)||Au3+(aq, 1 M)|Au(s)
oxidation reduction

96. Line Notation: ANODE CATHODE
Al(s) | Al3+(aq, 1 M) || Ni2+(aq, 1 M) | Ni (s)
oxidation reduction