Chemistry
Third Edition
Julia Burdge
Lecture PowerPoints
Chapter 19
Electrochemistry
Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display.
CHAPTER
19.1
19.2
19.3
19.4
19.5
19.6
19.7
19.8
19
Electrochemistry
Balancing Redox Reactions
Galvanic Cells
Standard Reduction Potentials
Spontaneity of Redox Reactions Under StandardState Conditions
Spontaneity of Redox Reactions Under Conditions
Other Than Standard State
Batteries
Electrolysis
Corrosion
2
19.1
Balancing Redox Reactions
Topics
Balancing Redox Reactions
3
19.1
Balancing Redox Reactions
Balancing Redox Reactions
A redox reaction is one in which there are changes in
oxidation states.
4
19.1
Balancing Redox Reactions
Balancing Redox Reactions
Consider the aqueous reaction of the iron(II) ion with the
dichromate ion:
• The reaction takes place in aqueous solution, so we can
add H2O as needed to balance the equation.
• This particular reaction takes place in acidic solution, so
we can add H+ as needed to balance the equation.
5
19.1
Balancing Redox Reactions
Balancing Redox Reactions
1. Separate the unbalanced reaction into half-reactions. A
half-reaction is an oxidation or a reduction that occurs as
part of the overall redox reaction.
6
19.1
Balancing Redox Reactions
Balancing Redox Reactions
2. Balance each of the half-reactions with regard to atoms
other than O and H.
In this case, no change is required for the oxidation halfreaction. We adjust the coefficient of the chromium(III)
ion to balance the reduction half-reaction.
7
19.1
Balancing Redox Reactions
Balancing Redox Reactions
3. Balance both half-reactions for O by adding H2O. Again,
the oxidation in this case requires no change, but we must
add seven water molecules to the product side of the
reduction.
8
19.1
Balancing Redox Reactions
Balancing Redox Reactions
4. Balance both half-reactions for H by adding H+.
Once again, the oxidation in this case requires no change,
but we must add 14 hydrogen ions to the reactant side of
the reduction.
9
19.1
Balancing Redox Reactions
Balancing Redox Reactions
5. Balance both half-reactions for charge by adding
electrons.
10
19.1
Balancing Redox Reactions
Balancing Redox Reactions
6. Multiply one or both of the half-reactions by the
number(s) required to make the number of electrons the
same in both.
11
19.1
Balancing Redox Reactions
Balancing Redox Reactions
7. Add the balanced half-reactions back together and cancel
the electrons, in addition to any other identical terms that
appear on both sides.
12
19.1
Balancing Redox Reactions
Balancing Redox Reactions
8. For each H+ ion in the final equation, add one OH– ion to
each side of the equation, combining the H+ and OH– ions
to produce H2O.
9. Make any additional cancellations made necessary by the
new H2O molecules.
13
SAMPLE PROBLEM
19.1
Permanganate ion and iodide ion react in basic solution to
produce manganese(IV) oxide and molecular iodine. Use the
half-reaction method to balance the equation:
Setup
14
SAMPLE PROBLEM
19.1
Solution
15
SAMPLE PROBLEM
19.1
Solution
16
SAMPLE PROBLEM
19.1
Solution
17
SAMPLE PROBLEM
19.1
Solution
18
SAMPLE PROBLEM
19.1
Solution
19
SAMPLE PROBLEM
19.1
Solution
20
19.2
Galvanic Cells
Topics
Galvanic Cells
21
19.2
Galvanic Cells
Galvanic Cells
The experimental apparatus for generating electricity through
the use of a spontaneous reaction is called a galvanic cell.
22
19.2
Galvanic Cells
Galvanic Cells
The zinc and copper bars are called electrodes.
By definition, the anode in a galvanic cell is the electrode at
which oxidation occurs and the cathode is the electrode at
which reduction occurs.
Each combination of container, electrode, and solution is
called a half-cell.
23
19.2
Galvanic Cells
Galvanic Cells
To complete the electric circuit, and allow electrons to flow
through the external wire, the solutions must be connected
by a conducting medium through which the cations and
anions can move from one half-cell to the other.
This requirement is satisfied by a salt bridge .
24
19.2
Galvanic Cells
Galvanic Cells
Experimentally the difference in electrical potential between
the anode and the cathode is measured by a voltmeter and
the reading (in volts) is called the cell potential (Ecell).
25
19.3
Standard Reduction Potentials
Topics
Standard Reduction Potentials
26
19.3
Standard Reduction Potentials
Standard Reduction Potentials
The standard hydrogen
electrode (SHE):
27
19.3
Standard Reduction Potentials
Standard Reduction Potentials
28
19.3
Standard Reduction Potentials
Standard Reduction Potentials
29
19.3
Standard Reduction Potentials
Standard Reduction Potentials
30
SAMPLE PROBLEM
19.2
A galvanic cell consists of an Mg electrode in a 1.0 M
Mg(NO3)2 solution and a Cd electrode in a 1.0 M Cd(NO3)2
solution.
Determine the overall cell reaction, and calculate the
standard cell potential at 25°C.
Setup
31
SAMPLE PROBLEM
19.2
Solution
32
SAMPLE PROBLEM
19.3
Predict what redox reaction will take place, if any, when
molecular bromine (Br2) is added to
(a) a 1-M solution of NaI and
(b) a 1-M solution of NaCl. (Assume a temperature of 25°C.)
Setup
33
SAMPLE PROBLEM
19.3
Solution
(a)
Because the reduction potential of Br2 is greater than that of
I2, Br2 will be reduced to Br– and I– will be oxidized to I2.
Thus, the preceding reaction will occur.
34
SAMPLE PROBLEM
19.3
Solution
(b)
Because the reduction potential of Br2 is smaller than that of
Cl2, this reaction will not occur.
Cl2 is more readily reduced than Br2, so Br2 is not reduced by
Cl–.
35
19.4
Spontaneity of Redox Reactions Under
Standard-State Conditions
Topics
Spontaneity of Redox Reactions Under Standard-State
Conditions
36
19.4
Spontaneity of Redox Reactions Under
Standard-State Conditions
Spontaneity of Redox Reactions Under Standard-State
Conditions
37
19.4
Spontaneity of Redox Reactions Under
Standard-State Conditions
Spontaneity of Redox Reactions Under Standard-State
Conditions
38
19.4
Spontaneity of Redox Reactions Under
Standard-State Conditions
Spontaneity of Redox Reactions Under Standard-State
Conditions
39
SAMPLE PROBLEM
19.4
Calculate the standard free-energy change for the following
reaction at 25°C:
Setup
40
SAMPLE PROBLEM
19.4
Solution
41
SAMPLE PROBLEM
19.5
Calculate the equilibrium constant for the following reaction
at 25°C:
Setup
42
SAMPLE PROBLEM
19.5
Solution
43
19.5
Spontaneity of Redox Reactions Under
Conditions Other Than Standard State
Topics
The Nernst Equation
Concentration Cells
44
19.5
Spontaneity of Redox Reactions Under
Conditions Other Than Standard State
The Nernst Equation
45
19.5
Spontaneity of Redox Reactions Under
Conditions Other Than Standard State
The Nernst Equation
The Nernst equation
46
SAMPLE PROBLEM
19.6
Predict whether the following reaction will occur
spontaneously as written at 298 K:
assuming [Co2+] = 0.15 M and [Fe2+] = 0.68 M.
47
SAMPLE PROBLEM
19.6
Setup
48
SAMPLE PROBLEM
19.6
Solution
49
19.5
Spontaneity of Redox Reactions Under
Conditions Other Than Standard State
Concentration Cells
Because electrode potential depends on ion concentrations, it
is possible to construct a galvanic cell from two half-cells
composed of the same material but differing in ion
concentrations.
Such a cell is called a concentration cell.
50
19.5
Spontaneity of Redox Reactions Under
Conditions Other Than Standard State
51
19.5
Spontaneity of Redox Reactions Under
Conditions Other Than Standard State
Concentration Cells
52
SAMPLE PROBLEM
19.7
An electrochemical cell is constructed for the purpose of
determining the Ksp of silver cyanide (AgCN) at 25°C.
One half-cell consists of a silver electrode in a 1.00-M solution
of silver nitrate.
The other half-cell consists of a silver electrode in a saturated
solution of silver cyanide.
The cell potential is measured and found to be 0.470 V.
Determine the concentration of silver ion in the saturated
silver cyanide solution and the value of Ksp for AgCN.
53
SAMPLE PROBLEM
19.7
Setup
Because this is a concentration cell, E°cell = 0 V.
The reaction quotient, Q, is (x M)/(1.00 M); and the value of n
is 1.
54
SAMPLE PROBLEM
19.7
Solution
55
19.6
Batteries
Topics
Dry Cells and Alkaline Batteries
Lead Storage Batteries
Lithium-Ion Batteries
Fuel Cells
56
19.6
Batteries
Dry Cells and Alkaline Batteries
A battery is a galvanic cell, or a series of connected galvanic
cells, that can be used as a portable, self-contained source of
direct electric current.
A dry cell, so named because it has
no fluid component .
57
19.6
Batteries
Dry Cells and Alkaline Batteries
58
19.6
Batteries
Lead Storage Batteries
59
19.6
Batteries
Lead Storage Batteries
Unlike dry cells and alkaline batteries, the lead storage
battery is rechargeable.
60
19.6
Batteries
Lithium-Ion Batteries
The overall cell potential is 3.4 V, which is relatively large.
61
19.6
Batteries
Fuel Cells
To generate electricity, heat produced by the reaction is first
used to convert water to steam, which then drives a turbine,
which then drives a generator.
A significant fraction of the energy released in the form of
heat is lost to the surroundings at each step .
62
19.6
Batteries
Fuel Cells
Greater efficiency can be accomplished by a device known as
a fuel cell, a galvanic cell that requires a continuous supply of
reactants to keep functioning.
Strictly speaking, a fuel cell is not a battery because it is not
self-contained.
63
19.6
Batteries
Fuel Cells
64
19.6
Batteries
Fuel Cells
65
19.6
Batteries
Fuel Cells
66
19.7
Electrolysis
Topics
Electrolysis of Molten Sodium Chloride
Electrolysis of Water
Electrolysis of an Aqueous Sodium Chloride Solution
Quantitative Applications of Electrolysis
67
19.7
Electrolysis
Electrolysis of Molten Sodium Chloride
The process of using electric energy to drive a nonspontaneous chemical reaction is called electrolysis.
An electrolytic cell is one used to carry out electrolysis.
68
19.7
Electrolysis
Electrolysis of Molten Sodium Chloride
69
19.7
Electrolysis
Electrolysis of Molten Sodium Chloride
70
19.7
Electrolysis
Electrolysis of Water
71
19.7
Electrolysis
Electrolysis of Water
72
19.7
Electrolysis
Electrolysis of an Aqueous Sodium Chloride Solution
The reductions that might occur at the cathode are
73
19.7
Electrolysis
Electrolysis of an Aqueous Sodium Chloride Solution
The oxidations that might occur at the anode are
74
19.7
Electrolysis
Electrolysis of an Aqueous Sodium Chloride Solution
The standard reduction potentials of the two reactions are
not very different, but the values do suggest that the
oxidation of H2O should occur more readily.
75
19.7
Electrolysis
Electrolysis of an Aqueous Sodium Chloride Solution
However, by experiment we find that the gas produced at the
anode is Cl2, not O2.
In the study of electrolytic processes, we sometimes find that the
voltage required for a reaction is considerably higher than the
electrode potentials would indicate.
The overvoltage is the difference between the calculated voltage
and the actual voltage required to cause electrolysis.
The overvoltage for O2 formation is quite high. Under normal
operating conditions, therefore, Cl2 gas forms at the anode instead
of O2.
76
19.7
Electrolysis
Electrolysis of an Aqueous Sodium Chloride Solution
Thus, the half-cell reactions in the electrolysis of aqueous
sodium chloride are
77
19.7
Electrolysis
Quantitative Applications of Electrolysis
If we know the current and how long it is
applied, we can calculate the charge.
Knowing the charge enables us to determine
the number of moles of electrons.
And knowing the number of moles of
electrons allows us to use stoichiometry to
determine the number of moles of product.
78
19.7
Electrolysis
Quantitative Applications of Electrolysis
Consider an electrolytic cell in which molten CaCl2 is
separated into its constituent elements, Ca and Cl2.
Suppose a current of 0.452 A is passed through the cell for
1.50 h.
How much product will be formed at each electrode?
79
19.7
Electrolysis
Quantitative Applications of Electrolysis
80
SAMPLE PROBLEM
19.8
A current of 1.26 A is passed through an electrolytic cell
containing a dilute sulfuric acid solution for 7.44 h.
Write the half-cell reactions, and calculate the volume of
gases generated at STP.
Setup
81
SAMPLE PROBLEM
19.8
Solution
82
SAMPLE PROBLEM
19.8
Solution
83
19.8
Corrosion
Topics
Corrosion
84
19.8
Corrosion
Corrosion
The term corrosion generally refers to the deterioration of a
metal by an electrochemical process.
85
19.8
Corrosion
Corrosion
The Fe2+ ions formed at the anode are further oxidized by
oxygen as follows:
86
19.8
Corrosion
Corrosion
87
19.8
Corrosion
Corrosion
Coinage metals such as copper and silver also corrode, but
much more slowly than either iron or aluminum:
Copper forms a layer of copper carbonate (CuCO3), a green
substance referred to as patina, that protects the metal
underneath from further corrosion.
Likewise, silverware that comes into contact with foodstuffs
develops a layer of silver sulfide (Ag2S).
88
19.8
Corrosion
Corrosion
An iron container can be covered with a layer of another
metal such as tin or zinc.
A “tin” can is made by applying a thin layer of tin over iron.
Rust formation is prevented as long as the tin layer remains
intact.
However, once the surface has been breached by a scratch or
a dent, rusting occurs rapidly.
89
19.8
Corrosion
Corrosion
90
19.8
Corrosion
Corrosion
Zinc-plating, or galvanization, protects iron from corrosion by
a different mechanism.
Zinc is more easily oxidized than iron. Like aluminum, zinc
oxidizes to form a protective coating.
Even when the zinc layer is compromised, though, and the
underlying iron is exposed, zinc is still the more easily oxidized
of the two metals and will act as the anode. Iron will be the
cathode, thereby remaining reduced.
91
19.8
Corrosion
Corrosion
Galvanization is one example of cathodic protection, the
process by which a metal is protected by being made the
cathode in what amounts to a galvanic cell.
Another example is the use of zinc or magnesium bars to
protect underground storage tanks and ships.
When a steel tank or hull is connected to a more easily
oxidized metal, corrosion of the steel is prevented.
92