1. Chapter 20 Electrochemistry
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Chapter 20 Electrochemistry
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.

2. Electrochemical Reactions
In electrochemical reactions, electrons are transferred from one species to another.

3. Oxidation Numbers
In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

4. Oxidation and Reduction
A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.

5. Oxidation and Reduction
A species is reduced when it gains electrons.
Here, each of the H+ gains an electron and they combine to form H2.

6. Given the following reaction, what chemical process is happening to Mg?
Mg + 2H+  Mg2+ + H2
oxidation
reduction
combustion
None of the above 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

7. Given the following reaction, what chemical process is happening to H+?
Mg + 2H+  Mg2+ + H2
oxidation
reduction
combustion
None of the above 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

8. Oxidation and Reduction
What is reduced is the oxidizing agent.
H+ oxidizes Zn by taking electrons from it.
What is oxidized is the reducing agent.
Zn reduces H+ by giving it electrons.

9. Given the following reaction, which species is acting as the reducing agent?
Mg + 2H+  Mg2+ + H2 H+ Mg
Mg2+ H2 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

10. Given the following reaction, which species is acting as the oxidizing agent?
Mg + 2H+  Mg2+ + H2 H2 Mg
Mg2+ H+ 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

11. Assigning Oxidation Numbers
Elements in their elemental form have an oxidation number of 0.
The oxidation number of a monatomic ion is the same as its charge.

12. What is the oxidation state of magnesium metal?
Mg + 2H+  Mg2+ + H2 -2 +1 +2 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

13. What is the oxidation state of hydrogen gas?
Mg + 2H+  Mg2+ + H2 -2 +1 +2 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

14. Assigning Oxidation Numbers
Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1.
Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.

15. What is the oxidation state of oxygen in the permanganate ion?
MnO4-1 -8 -1 -2 -4 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

16. What is the oxidation state of hydrogen in H2SO4?+1 -1 +2
None of the above 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

17. What is the oxidation state of hydrogen in LiH?+1 -1 +2
None of the above 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

18. Assigning Oxidation Numbers
Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
Fluorine always has an oxidation number of −1.
The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.

19. Assigning Oxidation Numbers
The sum of the oxidation numbers in a neutral compound is 0.
The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

20. What is the oxidation state of manganese in the permanganate ion?
MnO4-1 +8 -1 +2 +7 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

21. What is the oxidation state of sulfur in the H2SO4?+6 +4 +8 -6 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

22. What is the oxidation state of chlorine in HCl?+1 -1
Can’t tell 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

23. What is the oxidation state of chlorine in the HClO4?+4 +7 -7 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

24. What chemical process is happening to manganese in the following reaction?
MnO4− + C2O42-  Mn2+ + CO2
oxidation
reduction
Both 1 and 2
None of the above 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

25. What chemical process is happening to oxygen in the following reaction?
MnO4− + C2O42-  Mn2+ + CO2
oxidation
reduction
Both 1 and 2
None of the above 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

26. What chemical process is happening to carbon in the following reaction?
MnO4− + C2O42-  Mn2+ + CO2
oxidation
reduction
Both 1 and 2
None of the above 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

27. Balancing Oxidation-Reduction Equations
Perhaps the easiest way to balance the equation of an oxidation-reduction reaction is via the half-reaction method.

28. Balancing Oxidation-Reduction Equations
This involves treating (on paper only) the oxidation and reduction as two separate processes, balancing these half reactions, and then combining them to attain the balanced equation for the overall reaction.

29. Half-Reaction Method
Assign oxidation numbers to determine what is oxidized and what is reduced.
Write the oxidation and reduction half-reactions.

30. Half-Reaction Method Balance each half-reaction.
Balance elements other than H and O.
Balance O by adding H2O.
Balance H by adding H+.
Balance charge by adding electrons.
Multiply the half-reactions by integers so that the electrons gained and lost are the same.

31. Half-Reaction Method
Add the half-reactions, subtracting things that appear on both sides.
Make sure the equation is balanced according to mass.
Make sure the equation is balanced according to charge.

32. MnO4−(aq) + C2O42−(aq)  Mn2+(aq) + CO2(aq)
Half-Reaction Method
Consider the reaction between MnO4− and C2O42− :
MnO4−(aq) + C2O42−(aq)  Mn2+(aq) + CO2(aq)

33. Half-Reaction Method MnO4− + C2O42-  Mn2+ + CO2
First, we assign oxidation numbers.
MnO4− + C2O42-  Mn2+ + CO2 +7 +3 +4 +2
Since the manganese goes from +7 to +2, it is reduced.
Since the carbon goes from +3 to +4, it is oxidized.

34. Oxidation Half-Reaction
C2O42−  CO2
To balance the carbon, we add a coefficient of 2:
C2O42−  2 CO2

35. Oxidation Half-Reaction
C2O42−  2 CO2
The oxygen is now balanced as well. To balance the charge, we must add 2 electrons to the right side.
C2O42−  2 CO2 + 2 e−

36. Reduction Half-Reaction
MnO4−  Mn2+
The manganese is balanced; to balance the oxygen, we must add 4 waters to the right side.
MnO4−  Mn H2O

37. Reduction Half-Reaction
MnO4−  Mn H2O
To balance the hydrogen, we add 8 H+ to the left side.
8 H+ + MnO4−  Mn H2O

38. Reduction Half-Reaction
8 H+ + MnO4−  Mn H2O
To balance the charge, we add 5 e− to the left side.
5 e− + 8 H+ + MnO4−  Mn H2O

39. Combining the Half-Reactions
Now we evaluate the two half-reactions together:
C2O42−  2 CO2 + 2 e−
5 e− + 8 H+ + MnO4−  Mn H2O
To attain the same number of electrons on each side, we will multiply the first reaction by 5 and the second by 2.

40. Combining the Half-Reactions
5 C2O42−  10 CO e−
10 e− + 16 H+ + 2 MnO4−  2 Mn H2O
When we add these together, we get:
10 e− + 16 H+ + 2 MnO4− + 5 C2O42− 
2 Mn H2O + 10 CO2 +10 e−

41. Combining the Half-Reactions
10 e− + 16 H+ + 2 MnO4− + 5 C2O42− 
2 Mn H2O + 10 CO2 +10 e−
The only thing that appears on both sides are the electrons. Subtracting them, we are left with:
16 H+ + 2 MnO4− + 5 C2O42− 
2 Mn H2O + 10 CO2

42. Balancing in Basic Solution
If a reaction occurs in basic solution, one can balance it as if it occurred in acid.
Once the equation is balanced, add OH− to each side to “neutralize” the H+ in the equation and create water in its place.
If this produces water on both sides, you might have to subtract water from each side.

43. Activity series Some elements are more reactive than others.
What factors make elements more reactive?
Atomic radius (shielding/nuclear charge)
Stability
Activity series shows the order of reactivity

44. Do metals… Lose electrons? Gain electrons? Both of the above
Neither 1 nor 2 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

45. Which alkali metal is the most reactive?
Lithium
Sodium
Potassium
Rubidium 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

46. Activity Series
Activity series is usually written in terms of species most likely to be oxidized
Elements higher on the chart want to be oxidized more than elements lower on the chart.
Example: Potassium wants to be oxidized more than Sodium
What elements want to be oxidized more than Silver?
In Zumdahl, it is similar to the Standard Reduction Potential chart which is written in terms of reduction
Species more likely to be reduced are less likely to be oxidized
So in Zumdahl, the chart is in the opposite order

47. Activity Series
Activity series allow you to predict whether a redox reaction will happen or not.
Elements higher on the chart will displace a metal ion of an element lower on the chart.
Mg + Zn2+ will react to form Zn and Mg2+

48. Will K reduce Na+? Yes No Need more information 1 2 3 4 5 6 7 8 9 1011 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

49. Will Mg2+ reduce Na+? Yes No Need more information 1 2 3 4 5 6 7 8 910 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

50. Will Mg reduce Na+? Yes No Need more information 1 2 3 4 5 6 7 8 9 1011 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

51. Predicting Activity Series from Data
X + Y2+  No reaction
Y + Z2+  Reacts
Z + X2+  Reacts
Based on reaction 1, I know that Y is more active than X
Based on reaction 2, I know that Y is more reactive than Z
Based on reaction 3, I know that Z is more reactive than X
Activity series from most to least: Y > Z > X

52. Given the following data, put the elements in order from most to least reactive.
X + Y2+  X2+ + Y
M + Y2+  M+1 + Y
M + X2+  no reaction
M, X, Y
Y, X, M
Y, M, X
X, M, Y 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

53. Given the following data, put the elements in order from most to least reactive.
P + Q2+  Q2+ + P
M + Q2+  M2+ + Q
P, Q, M
M, Q, P
P, M, Q
Q, M, P 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

54. Standard Reduction Potentials
Reduction potentials for many electrodes have been measured and tabulated.

55. Voltaic Cells
In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.

56. Voltaic Cells
We can use that energy to do work if we make the electrons flow through an external device.
We call such a setup a voltaic cell.

57. Voltaic Cells A typical cell looks like this.
The oxidation occurs at the anode.
The reduction occurs at the cathode.

58. Which electrode is made of the most active metal?
Anode
Cathode
They have the same activity 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

59. Voltaic Cells
In the cell, then, electrons leave the anode and flow through the wire to the cathode.
As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

60. Voltaic Cells
As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode.
The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

61. In a voltaic cell, electrons flow from
Positive to negative electrode
Cathode to negative electrode
Anode to negative electrode
Negative to positive electrode 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

62. Which electrode loses mass?
anode
cathode
Both gain mass
Both lose mass 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

63. Voltaic Cells: Purpose of Salt Bridge
Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop.

64. Voltaic Cells: Ions, not electrons, move through salt bridge
Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced.
Cations move toward the cathode.
Anions move toward the anode.

65. Without the salt bridge, which of the following is true
Electrons build up in the anode cell
Excess positive charge builds up in the anode cell.
Excess negative charge builds up in the anode cell
Excess positive charge builds up in the cathode cell 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

66. Which is true about the salt bridge?
Electrons move from anode cell to cathode cell
Electrons move from cathode cell to anode cell
Cations move to the anode cell
Anions move to the anode cell 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

67. If you made a voltaic cell with iron and magnesium, which would be at the cathode?
Iron (III) Nitrate
Magnesium Nitrate 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

68. How can you remember? a RED CAT and AN OX REDuction = CAThode
ANode = OXidation

69. Applications of Oxidation-Reduction Reactions

70. Batteries

71. Alkaline Batteries

72. Hydrogen Fuel Cells Write the ½ reactions for this process.
Check your answers with data book.

73. Electrolytic Cell What if we made the current go backwards? How?
With a power supply…
Do you have any batteries that you plug in to recharge?

74. - + Voltaic Cell Anode: Oxidation occurs here
Positive
Electrode
Anions are attracted to cathode
Cathode: Reduction occurs here
Negative
Electrode
Cations
attracted to anode
Electrolyte: allows for charges to move from one electrode to the other.

75. - + Electrolytic Cell Power Source Cathode: Reduction occurs here
Negative
Electrode
Cations
attracted to anode
Anode: Oxidation occurs here
Positive
Electrode
Anions are attracted to cathode
Electrolyte: allows for charges to move from one electrode to the other.

76. Major Differences Direction of electron flow changes
Cathode and anode switch so that oxidation still occurs at the anode and reduction still occurs at the cathode.
Cathode is now the negative electrode, it attracts cations
Anode is now the positive electrode, it attracts anions
Power supply needed because reaction is nonspontaneous
Salt bridge isn’t necessary

77. Electrolysis Uses
Electroplating (putting a coating of a metal on some cheaper metal) (Cu)
Producing pure elements (PbBr2)
Requires molten salt
Example: NaCl(l)  Na(l) + Cl2(g)
Producing Sodium hydroxide (chlor-alkali industry)
Dilute aqueous solution

78. Electrons travel from negative to positive electrode in electolytic cells.
True
False 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24

79. What species are produced at the positive and negative electrodes during the electrolysis of molten sodium chloride?
+ electrode - electrode
A. Na+(l) Cl2(g)
B. Cl–(l) Na+(l)
C. Na(l) Cl2(g)
D. Cl2(g) Na(l)
True
False
Sdf
xcv 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

80. Demo Electrolysis of Sodium Iodide Solution NaI  NaOH + I2
See evidence of base b/c phenopthalein
See evidence of I2 b/c of starch
With dilute aqueous solutions you get a metal hydroxide instead of a pure metal
Iodine is oxidized
Hydrogen from the water is reduced to form H2(g)

81. What are the products of the electrolysis of dilute aqueous sodium chloride?
NaOH and Cl2
H2 and Cl2
Na and Cl2
OH- and Cl2 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30

82. What are the products of the electrolysis of molten sodium chloride?
NaOH and Cl2
H2 and CL2
Na and Cl2
OH- and Cl2 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30