1. 21
Electrochemistry

2. Chapter Goals Electrical Conduction Electrodes Electrolytic Cells
The Electrolysis of Molten Sodium Chloride (the Downs Cell)
The Electrolysis of Aqueous Sodium Chloride
The Electrolysis of Aqueous Sodium Sulfate
Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
Commercial Applications of Electrolytic Cells

3. Voltaic or Galvanic Cells Standard Electrode Potentials
Chapter Goals
Voltaic or Galvanic Cells
The Construction of Simple Voltaic Cells
The Zinc-Copper Cell
The Copper-Silver Cell
Standard Electrode Potentials
The Standard Hydrogen Electrode
The Zinc-SHE Cell
The Copper-SHE Cell
Uses of Standard Electrode Potentials

4. Chapter Goals Standard Electrode Potentials for Other Half-Reactions
Corrosion
Corrosion Protection
Effect of Concentrations (or Partial Pressures) on Electrode Potentials
The Nernst Equation
Using Electrohemical Cells to Determine Concentrations
The Relationship of Eocell to Go and K

5. Secondary Voltaic Cells
Chapter Goals
Primary Voltaic Cells
Dry Cells
Secondary Voltaic Cells
The Lead Storage Battery
The Nickel-Cadmium (Nicad) Cell
The Hydrogen-Oxygen Fuel Cell

6. Electrochemistry
Electrochemical reactions are oxidation-reduction reactions.
The two parts of the reaction are physically separated.
The oxidation reaction occurs in one cell.
The reduction reaction occurs in the other cell.

7. Electrochemistry There are two kinds electrochemical cells.
Electrochemical cells containing in nonspontaneous chemical reactions are called electrolytic cells.
Electrochemical cells containing spontaneous chemical reactions are called voltaic or galvanic cells.

8. Electrical Conduction
Metals conduct electric currents well in a process called metallic conduction.
In metallic conduction there is electron flow with no atomic motion.
In ionic or electrolytic conduction ionic motion transports the electrons.
Positively charged ions, cations, move toward the negative electrode.
Negatively charged ions, anions, move toward the positive electrode.

9. Electrodes
The following convention for electrodes is correct for either electrolytic or voltaic cells:
The cathode is the electrode at which reduction occurs.
The cathode is negative in electrolytic cells and positive in voltaic cells.
The anode is the electrode at which oxidation occurs.
The anode is positive in electrolytic cells and negative in voltaic cells.

10. Electrodes
Inert electrodes do not react with the liquids or products of the electrochemical reaction.
Two examples of common inert electrodes are graphite and platinum.

11. Electrolytic Cells
Electrical energy is used to force nonspontaneous chemical reactions to occur.
The process is called electrolysis.
Two examples of commercial electrolytic reactions are:
The electroplating of jewelry and auto parts.
The electrolysis of chemical compounds.

12. Electrolytic Cells Electrolytic cells consist of:
A container for the reaction mixture.
Two electrodes immersed in the reaction mixture.
A source of direct current.

13. The Electrolysis of Molten Sodium Chloride
Liquid Sodium is produced at one electrode.
Indicates that the reaction Na+() + e-  Na(s) occurs at this electrode.
Is this electrode the anode or cathode?
Gaseous chlorine is produced at the other electrode.
Indicates that the reaction 2 Cl-  Cl2(g) + 2 e- occurs at this electrode.

14. The Electrolysis of Molten Sodium Chloride
Diagram of this electrolytic cell.

15. The Electrolysis of Molten Sodium Chloride
The nonspontaneous redox reaction that occurs is:

16. The Electrolysis of Molten Sodium Chloride
In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the negative electrode (cathode).

17. The Electrolysis of Aqueous Sodium Chloride
In this electrolytic cell, hydrogen gas is produced at one electrode.
The aqueous solution becomes basic near this electrode.
What reaction is occurring at this electrode? You do it!
Gaseous chlorine is produced at the other electrode.
These experimental facts lead us to the following nonspontaneous electrode reactions:

18. The Electrolysis of Aqueous Sodium Chloride

19. The Electrolysis of Aqueous Sodium Chloride
Cell diagram

20. The Electrolysis of Aqueous Sodium Sulfate
In this electrolysis, hydrogen gas is produced at one electrode.
The solution becomes basic near this electrode.
What reaction is occurring at this electrode?
You do it!
Gaseous oxygen is produced at the other electrode
The solution becomes acidic near this electrode.
These experimental facts lead us to the following electrode reactions:

21. The Electrolysis of Aqueous Sodium Sulfate

22. The Electrolysis of Aqueous Sodium Sulfate

23. Electrolytic Cells
In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized.

24. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
Faraday’s Law - The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the electrolytic cell.
A faraday is the amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode.

25. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
A coulomb is the amount of charge that passes a given point when a current of one ampere (A) flows for one second.
1 amp = 1 coulomb/second

26. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
Faraday’s Law states that during electrolysis, one faraday of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent.
This corresponds to the passage of one mole of electrons through the electrolytic cell.

27. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
Example 21-1: Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes.

28. Counting Electrons: Coulometry and Faraday’s Law of Electrolysis
Example 21-2: Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in example 21-1.

29. Commercial Applications of Electrolytic Cells
Electrolytic Refining and Electroplating of Metals
Impure metallic copper can be purified electrolytically to  100% pure Cu.
The impurities commonly include some active metals plus less active metals such as: Ag, Au, and Pt.
The cathode is a thin sheet of copper metal connected to the negative terminal of a direct current source.
The anode is large impure bars of copper.

30. Commercial Applications of Electrolytic Cells
The electrolytic solution is CuSO4 and H2SO4
The impure Cu dissolves to form Cu2+.
The Cu2+ ions are reduced to Cu at the cathode.

31. Commercial Applications of Electrolytic Cells
Any active metal impurities are oxidized to cations that are more difficult to reduce than Cu2+.
This effectively removes them from the Cu metal.

32. Commercial Applications of Electrolytic Cells
The less active metals are not oxidized and precipitate to the bottom of the cell.
These metal impurities can be isolated and separated after the cell is disconnected.
Some common metals that precipitate include:

33. Voltaic or Galvanic Cells
Electrochemical cells in which a spontaneous chemical reaction produces electrical energy.
Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference.
Examples of voltaic cells include:

34. The Construction of Simple Voltaic Cells
Voltaic cells consist of two half-cells which contain the oxidized and reduced forms of an element (or other chemical species) in contact with each other.
A simple half-cell consists of:
A piece of metal immersed in a solution of its ions.
A wire to connect the two half-cells.
And a salt bridge to complete the circuit, maintain neutrality, and prevent solution mixing.

35. The Construction of Simple Voltaic Cells

36. The Zinc-Copper Cell Cell components for the Zn-Cu cell are:
A metallic Cu strip immersed in 1.0 M copper (II) sulfate.
A metallic Zn strip immersed in 1.0 M zinc (II) sulfate.
A wire and a salt bridge to complete circuit
The cell’s initial voltage is 1.10 volts

37. The Zinc-Copper Cell

38. The Zinc-Copper Cell
In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

39. The Zinc-Copper Cell
There is a commonly used short hand notation for voltaic cells.
The Zn-Cu cell provides a good example.

40. The Copper - Silver Cell
Cell components:
A Cu strip immersed in 1.0 M copper (II) sulfate.
A Ag strip immersed in 1.0 M silver (I) nitrate.
A wire and a salt bridge to complete the circuit.
The initial cell voltage is 0.46 volts.

41. The Copper - Silver Cell

42. The Copper - Silver Cell
Compare the Zn-Cu cell to the Cu-Ag cell
The Cu electrode is the cathode in the Zn-Cu cell.
The Cu electrode is the anode in the Cu-Ag cell.
Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.

43. The Copper - Silver Cell
These experimental facts demonstrate that Cu2+ is a stronger oxidizing agent than Zn2+.
In other words Cu2+ oxidizes metallic Zn to Zn2+.
Similarly, Ag+ is is a stronger oxidizing agent than Cu2+.
Because Ag+ oxidizes metallic Cu to Cu 2+.
If we arrange these species in order of increasing strengths, we see that:

44. Standard Electrode Potential
To measure relative electrode potentials, we must establish an arbitrary standard.
That standard is the Standard Hydrogen Electrode (SHE).
The SHE is assigned an arbitrary voltage of … V

45. Standard Electrode Potential

46. The Zinc-SHE Cell For this cell the components are:
A Zn strip immersed in 1.0 M zinc (II) sulfate.
The other electrode is the Standard Hydrogen Electrode.
A wire and a salt bridge to complete the circuit.
The initial cell voltage is volts.

47. The Zinc-SHE Cell

48. The Zinc-SHE Cell The cathode is the Standard Hydrogen Electrode.
In other words Zn reduces H+ to H2.
The anode is Zn metal.
Zn metal is oxidized to Zn2+ ions.

49. The Copper-SHE Cell The cell components are:
A Cu strip immersed in 1.0 M copper (II) sulfate.
The other electrode is a Standard Hydrogen Electrode.
A wire and a salt bridge to complete the circuit.
The initial cell voltage is volts.

50. The Copper-SHE Cell

51. The Copper-SHE Cell In this cell the SHE is the anode
The Cu2+ ions oxidize H2 to H+.
The Cu is the cathode.
The Cu2+ ions are reduced to Cu metal.

52. Uses of Standard Electrode Potentials
Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials.
Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials.
Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.
For example, the half-reaction for the standard potassium electrode is:
The large negative value tells us that this reaction
will occur only under extreme conditions.

53. Uses of Standard Electrode Potentials
Compare the potassium half-reaction to fluorine’s half-reaction:
The large positive value denotes that this reaction occurs readily as written.
Positive E0 values denote that the reaction tends to occur to the right.
The larger the value, the greater the tendency to occur to the right.
It is the opposite for negative values of Eo.

54. Uses of Standard Electrode Potentials
Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously.
Example 21-3: Will silver ions, Ag+, oxidize metallic zinc to Zn2+ ions, or will Zn2+ ions oxidize metallic Ag to Ag+ ions?
Steps for obtaining the equation for the spontaneous reaction.

55. Uses of Standard Electrode Potentials
Choose the appropriate half-reactions from a table of standard reduction potentials.
Write the equation for the half-reaction with the more positive E0 value first, along with its E0 value.
Write the equation for the other half-reaction as an oxidation with its oxidation potential, i.e. reverse the tabulated reduction half-reaction and change the sign of the tabulated E0.
Balance the electron transfer.
Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E0cell is positive, which indicates that the forward reaction is spontaneous.

56. Uses of Standard Electrode Potentials

57. Electrode Potentials for Other Half-Reactions
Example 21-4: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution?
Follow the steps outlined in the previous slides.
Note that E0 values are not multiplied by any stoichiometric relationships in this procedure.

58. Electrode Potentials for Other Half-Reactions
Example 21-4: Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution?
Thus permanganate ions will oxidize iron (II) ions to iron (III) and are reduced to manganese (II) ions in acidic solution.

59. Electrode Potentials for Other Half-Reactions
Example 21-5: Will nitric acid, HNO3, oxidize arsenous acid, H3AsO3, in acidic solution? The reduction product of HNO3 is NO in this reaction.
You do it!

60. Corrosion
Metallic corrosion is the oxidation-reduction reactions of a metal with atmospheric components such as CO2, O2, and H2O.

61. Corrosion Protection Some examples of corrosion protection.
Plate a metal with a thin layer of a less active (less easily oxidized) metal.

62. Corrosion Protection
Connect the metal to a sacrificial anode, a piece of a more active metal.

63. Corrosion Protection

64. Corrosion Protection
Allow a protective film to form naturally.

65. Corrosion Protection
Galvanizing, the coating of steel with zinc, provides a more active metal on the exterior.

66. Corrosion Protection
Paint or coat with a polymeric material such as plastic or ceramic.

67. Effect of Concentrations (or Partial Pressures) on Electrode Potentials
The Nernst Equation
Standard electrode potentials, those compiled in appendices, are determined at thermodynamic standard conditions.
Reminder of standard conditions.
1.00 M solution concentrations
1.00 atm of pressure for gases
All liquids and solids in their standard thermodynamic states.
Temperature of 250 C.

68. The Nernst Equation
The value of the cell potentials change if conditions are nonstandard.
The Nernst equation describes the electrode potentials at nonstandard conditions.
The Nernst equation is:

69. The Nernst Equation

70. The Nernst Equation
Substitution of the values of the constants into the Nernst equation at 25o C gives:

71. The Nernst Equation For this half-reaction:
The corresponding Nernst equation is:

72. The Nernst Equation Substituting E0 into the above expression gives:
If [Cu2+] and [Cu+] are both 1.0 M, i.e. at standard conditions, then E = E0 because the concentration term equals zero.

73. The Nernst Equation

74. The Nernst Equation
Example 21-6: Calculate the potential for the Cu2+/ Cu+ electrode at 250C when the concentration of Cu+ ions is three times that of Cu2+ ions.

75. The Nernst Equation
Example 21-6: Calculate the potential for the Cu2+/ Cu+ electrode at 250C when the concentration of Cu+ ions is three times that of Cu2+ ions.

76. The Nernst Equation
Example 21-6: Calculate the potential for the Cu2+/ Cu+ electrode at 250C when the concentration of Cu+ ions is three times that of Cu2+ ions.

77. The Nernst Equation

78. The Nernst Equation
Example 21-7: Calculate the potential for the Cu2+/Cu+ electrode at 250C when the Cu+ ion concentration is 1/3 of the Cu2+ ion concentration.
You do it!

79. The Nernst Equation
Example 21-7: Calculate the potential for the Cu2+/Cu+ electrode at 250C when the concentration of Cu+ ions is 1/3 that of Cu2+ ions.

80. The Nernst Equation

81. The Nernst Equation
Example 21-8: Calculate the electrode potential for a hydrogen electrode in which the [H+] is 1.0 x 10-3 M and the H2 pressure is 0.50 atmosphere.

82. The Nernst Equation
Example 21-8: Calculate the electrode potential for a hydrogen electrode in which the [H+] is 1.0 x 10-3 M and the H2 pressure is 0.50 atmosphere.

83. The Nernst Equation
The Nernst equation can also be used to calculate the potential for a cell that consists of two nonstandard electrodes.
Example 21-9: Calculate the initial potential of a cell that consists of an Fe3+/Fe2+ electrode in which [Fe3+]=1.0 x 10-2 M and [Fe2+]=0.1 M connected to a Sn4+/Sn2+ electrode in which [Sn4+]=1.0 M and [Sn2+]=0.10 M . A wire and salt bridge complete the circuit.

84. The Nernst Equation
Calculate the E0 cell by the usual procedure.

85. The Nernst Equation
Substitute the ion concentrations into Q to calculate Ecell.

86. The Nernst Equation

87. Relationship of E0cell to G0 and K
From previous chapters we know the relationship of G0 and K for a reaction.

88. Relationship of E0cell to G0 and K
The relationship between G0 and E0cell is also a simple one.

89. Relationship of E0cell to G0 and K
Combine these two relationships into a single relationship to relate E0cell to K.

90. Relationship of E0cell to G0 and K
Example 21-10: Calculate the standard Gibbs free energy change, G0 , at 250C for the following reaction.

91. Relationship of E0cell to G0 and K
Calculate E0cell using the appropriate half-reactions.

92. Relationship of E0cell to G0 and K
Now that we know E0cell , we can calculate G0 .

93. Relationship of E0cell to G0 and K
Example 21-11: Calculate the thermodynamic equilibrium constant for the reaction in example at 250C.

94. Relationship of E0cell to G0 and K
Example 21-12: Calculate the Gibbs Free Energy change, G and the equilibrium constant at 250C for the following reaction with the indicated concentrations.

95. Relationship of E0cell to G0 and K
Calculate the standard cell potential E0cell.

96. Relationship of E0cell to G0 and K
Use the Nernst equation to calculate Ecell for the given concentrations.

97. Relationship of E0cell to G0 and K

98. Relationship of E0cell to G0 and K

99. Relationship of E0cell to G0 and K
Ecell = V, compared to E0cell = V.
We can use this information to calculate G.
The negative G tells us that the reaction is spontaneous.

100. Relationship of E0cell to G0 and K
Equilibrium constants do not change with reactant concentration.
We can use the value of E0cell at 250C to get K.

101. Primary Voltaic Cells
As a voltaic cell discharges, its chemicals are consumed.
Once the chemicals are consumed, further chemical action is impossible.
The electrodes and electrolytes cannot be regenerated by reversing current flow through cell.
These cells are not rechargable.

102. The Dry Cell
One example of a dry cell is flashlight, and radio, batteries.
The cell’s container is made of zinc which acts as an electrode.
A graphite rod is in the center of the cell which acts as the other electrode.
The space between the electrodes is filled with a mixture of:
ammonium chloride, NH4Cl
manganese (IV) oxide, MnO2
zinc chloride, ZnCl2
and a porous inactive solid.

103. The Dry Cell
As electric current is produced, Zn dissolves and goes into solution as Zn2+ ions.
The Zn electrode is negative and acts as the anode.

104. The Dry Cell The anode reaction is:
The graphite rod is the positive electrode (cathode).
Ammonium ions from the NH4Cl are reduced at the cathode.

105. The Dry Cell
The cell reaction is:

106. The Dry Cell
The other components in the cell are included to remove the byproducts of the reaction.
MnO2 prevents H2 from collecting on graphite rod.
At the anode, NH3 combines with Zn2+ to form a soluble complex and removing the Zn2+ ions from the reaction.

107. The Dry Cell

108. The Dry Cell
Alkaline dry cells are similar to ordinary dry cells except that KOH, an alkaline substance, is added to the mixture.
Half reactions for an alkaline cell are:

109. The Dry Cell
Alkaline dry cells are similar to ordinary dry cells except that KOH, an alkaline substance, is added to the mixture.
Half reactions for an alkaline cell are:

110. Secondary Voltaic Cells
Secondary cells are reversible, rechargeable.
The electrodes in a secondary cell can be regenerated by the addition of electricity.
These cells can be switched from voltaic to electrolytic cells.
One example of a secondary voltaic cell is the lead storage or car battery.

111. The Lead Storage Battery
In the lead storage battery the electrodes are two sets of lead alloy grids (plates).
Holes in one of the grids are filled with lead (IV) oxide, PbO2.
The other holes are filled with spongy lead.
The electrolyte is dilute sulfuric acid.

112. The Lead Storage Battery
Diagram of the lead storage battery.

113. The Lead Storage Battery
As the battery discharges, spongy lead is oxidized to lead ions and the plate becomes negatively charged.
The Pb2+ ions that are formed combine with SO42- from sulfuric acid to form solid lead sulfate on the Pb electrode.

114. The Lead Storage Battery
The net reaction at the anode during discharge is:
Electrons are produced at the Pb electrode.
These electrons flow through an external circuit (the wire and starter) to the PbO2 electrode.
PbO2 is reduced to Pb2+ ions, in the acidic solution.
The Pb2+ ions combine with SO42- to form PbSO4 and coat the PbO2 electrode.
PbO2 electrode is the positive electrode (cathode).

115. The Lead Storage Battery
As the cell discharges, the cathode reaction is:
The cell reaction for a discharging lead storage battery is:

116. The Lead Storage Battery
As the cell discharges, the cathode reaction is:
The cell reaction for a discharging lead storage battery is:

117. The Lead Storage Battery
What happens at each electrode during recharging?
At the lead (IV) oxide, PbO2, electrode, lead ions are oxidized to lead (IV) oxide.
The concentration of the H2SO4 decreases as the cell discharges.
Recharging the cell regenerates the H2SO4.

118. The Lead Storage Battery
What happens at each electrode during recharging?
At the lead (IV) oxide, PbO2, electrode, lead ions are oxidized to lead (IV) oxide.
The concentration of the H2SO4 decreases as the cell discharges.
Recharging the cell regenerates the H2SO4.

119. The Nickel-Cadmium (Nicad) Cell
Nicad batteries are the rechargeable cells used in calculators, cameras, watches, etc.
As the battery discharges, the half-reactions are:

120. The Hydrogen-Oxygen Fuel Cell
Fuel cells are batteries that must have their reactants continuously supplied in the presence of appropriate catalysts.
A hydrogen-oxygen fuel cell is used in the space shuttle
The fuel cell is what exploded in Apollo 13.
Hydrogen is oxidized at the anode.
Oxygen is reduced at the cathode.

121. The Hydrogen-Oxygen Fuel Cell

122. The Hydrogen-Oxygen Fuel Cell
Notice that the overall reaction is the combination of hydrogen and oxygen to form water.
The cell provides a drinking water supply for the astronauts as well as the electricity for the lights, computers, etc. on board.
Fuel cells are very efficient.
Energy conversion rates of 60-70% are common!

123. 21
Electrochemistry