Electrochemistry

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 Faradays Law
  of Electrolysis
• Commercial Applications of Electrolytic Cells

3
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
• Standard Electrode Potentials
• 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
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.
• Is this electrode the anode or cathode?

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.
• What reaction is occurring at this electrode?
  You do it!
• 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.
• What reaction is occurring at this electrode?
• You do it!
• 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 Faradays Law
of Electrolysis

• Faradays 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 Faradays 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 Faradays Law
of Electrolysis

• Faradays 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 Faradays 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 Faradays 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 cells 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
  0.000000 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 0.763 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 0.337 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 fluorines
  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

1. Connect the metal to a sacrificial anode, a piece
   of a more active metal.

63
Corrosion Protection
64
Corrosion Protection

1. 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

1. 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 Fe31.0 x 10-2 M and Fe20.1 M
  connected to a Sn4/Sn2 electrode in which
  Sn41.0 M and Sn20.10 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

1. Calculate E0cell using the appropriate
   half-reactions.

92
Relationship of E0cell to ?G0 and K

1. 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
  21-10 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

1. Calculate the standard cell potential E0cell.

96
Relationship of E0cell to ?G0 and K

1. 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 1.540 V, compared to E0cell 1.562 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 cells 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