Transcript Chemistry: McMurry and Fay, 5th Edition

John E. McMurry • Robert C. Fay
C H E M I S T R Y
Fifth Edition
Chapter 17
Electrochemistry
Lecture Notes
Alan D. Earhart
Southeast Community College • Lincoln, NE
Copyright © 2008 Pearson Prentice Hall, Inc.
Galvanic Cells
Electrochemistry: The area of chemistry concerned
with the interconversion of chemical and electrical
energy.
Galvanic (Voltaic) Cell: A spontaneous chemical
reaction which generates an electric current.
Electrolytic Cell: An electric current which drives a
nonspontaneous reaction.
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Chapter 17/2
Galvanic Cells
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Oxidation half-reaction:
Zn(s)
Reduction half-reaction:
Cu2+(aq) + 2e-
Zn2+(aq) + 2e-
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Cu(s)
Chapter 17/3
Galvanic Cells
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Galvanic Cells
•
Anode:
• The electrode where oxidation occurs.
• The electrode where electrons are produced.
• Is what anions migrate toward.
• Has a negative sign.
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Chapter 17/5
Galvanic Cells
•
Anode:
• The electrode where oxidation occurs.
• The electrode where electrons are produced.
• Is what anions migrate toward.
• Has a negative sign.
•
Cathode:
• The electrode where reduction occurs.
• The electrode where electrons are consumed.
• Is what cations migrate toward.
• Has a positive sign.
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Chapter 17/6
Galvanic Cells
Anode half-reaction:
Cathode half-reaction:
Overall cell reaction:
Zn(s)
Cu2+(aq) + 2eZn(s) + Cu2+(aq)
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Zn2+(aq) + 2eCu(s)
Zn2+(aq) + Cu(s)
Chapter 17/7
Shorthand Notation for
Galvanic Cells
Anode half-reaction:
Cathode half-reaction:
Overall cell reaction:
Zn2+(aq) + 2e-
Zn(s)
Cu2+(aq) + 2e-
Cu(s)
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Salt bridge
Anode half-cell
Cathode half-cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Electron flow
Phase boundary
Phase boundary
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Chapter 17/8
Cell Potentials and Free-Energy
Changes for Cell Reactions
Electromotive Force (emf): The force or electrical
potential that pushes the negatively charged electrons
away from the anode (- electrode) and pulls them toward
the cathode (+ electrode).
It is also called the cell potential (E) or the cell voltage.
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Chapter 17/9
Cell Potentials and Free-Energy
Changes for Cell Reactions
1J=1Cx1V
joule
SI unit of energy
volt
SI unit of electric potential
coulomb
Electric charge
1 coulomb is the amount of charge transferred
when a current of 1 ampere flows for 1 second.
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Chapter 17/10
Cell Potentials and Free-Energy
Changes for Cell Reactions
faraday or Faraday constant
the electric charge on 1 mol of electrons
96,5000 C/mol e-
DG = -nFE
free-energy change
or DG° = -nFE°
cell potential
number of moles of electrons
transferred in the reaction
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Chapter 17/11
Cell Potentials and Free-Energy
Changes for Cell Reactions
The standard cell potential at 25 °C is 0.10 V for the
reaction:
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Calculate the standard free-energy change for this reaction
at 25 °C.
DG° = -nFE°
= -(2 mol
e-)
96,500 C
mol e-
(1.10 V)
1J
1 kJ
1CV
1000 J
DG° = -212 kJ
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Chapter 17/12
Standard Reduction Potentials
H2(g) + Cu2+(aq)
2H1+(aq) + Cu(s)
Standard Reduction Potentials
Anode half-reaction:
H2(g)
Cathode half-reaction:
Overall cell reaction:
Cu2+(aq) + 2eH2(g) + Cu2+(aq)
2H1+(aq) + 2eCu(s)
2H1+(aq) + Cu(s)
The standard potential of a cell is the sum of the standard
half-cell potentials for oxidation at the anode and reduction
at the cathode:
E°cell = E°ox + E°red
The measured potential for this cell: E°cell = 0.34 V
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Chapter 17/14
Standard Reduction Potentials
The standard hydrogen electrode (S.H.E.) has been
chosen to be the reference electrode.
Standard Reduction Potentials
The standard hydrogen electrode (S.H.E.) has been
chosen to be the reference electrode.
2H1+(aq, 1 M) + 2eH2(g, 1 atm)
H2(g, 1 atm)
E°ox = 0 V
2H1+(aq, 1 M) + 2e-
E°red = 0 V
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Chapter 17/16
Standard Reduction Potentials
Anode half-reaction:
H2(g)
Cathode half-reaction:
Cu2+(aq) + 2eH2(g) + Cu2+(aq)
Overall cell reaction:
2H1+(aq) + 2eCu(s)
2H1+(aq) + Cu(s)
E°cell = E°ox + E°red
0.34 V = 0 V + E°red
A standard reduction potential can be defined:
Cu2+(aq) + 2e-
Cu(s)
E° = 0.34 V
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Chapter 17/17
Standard Reduction Potentials
2H1+(aq) + Zn(s)
H2(g) + Zn2+(aq)
Standard Reduction Potentials
Anode half-reaction:
2H1+(aq) + 2e-
Cathode half-reaction:
Overall cell reaction:
Zn(s)
2H1+(aq) + Zn(s)
H2(g)
Zn2+(aq) + 2eH2(g) + Zn2+(aq)
E°cell = E°ox + E°red
0.76 V = E°ox + 0 V
Zn(s)
Zn2+(aq) + 2e-
E° = 0.76 V
As a standard reduction potential:
Zn2+(aq) + 2e-
Zn(s)
E° = -0.76 V
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Chapter 17/19
Standard Reduction Potentials
Using Standard Reduction
Potentials
Zn(s)
Cu2+(aq) + 2eZn(s) + Cu2+(aq)
2 x [Ag1+(aq) + eCu(s)
2Ag(s) + Cu2+(aq)
Zn2+(aq) + 2e-
E° = -(-0.76 V)
Cu(s)
E° = 0.34 V
Zn2+(g) + Cu(s)
E° = 1.10 V
Ag(s)]
E° = 0.80 V
Cu2+(aq) + 2e-
E° = -0.34 V
2Ag2+(g) + Cu(s)
E° = 0.46 V
Half-cell potentials are intensive properties.
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Chapter 17/21
The Nernst Equation
DG = DG° + RT ln Q
Using:
Nernst Equation:
DG = -nFE and DG° = -nFE°
RT
E = E° ln Q
nF
or
2.303RT
E = E° log Q
nF
or
0.0592 V
E = E° log Q in volts, at 25°C
n
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Chapter 17/22
The Nernst Equation
Consider a galvanic cell that uses the reaction:
Cu(s) + 2Fe3+(aq)
Cu2+(aq) + 2Fe2+(aq)
What is the potential of a cell at 25 °C that has the following
ion concentrations?
[Fe3+] = 1.0 x 10-4 M
[Cu2+] = 0.25 M
[Fe2+] = 0.20 M
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Chapter 17/23
The Nernst Equation
E = E° -
0.0592 V
log Q
n
Calculate E°:
Cu(s)
Fe3+(aq) + e-
Cu2+(aq) + 2e-
E° = -0.34 V
Fe2+(aq)
E° = 0.77 V
E°cell = -0.34 V + 0.77 V = 0.43 V
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Chapter 17/24
The Nernst Equation
E = E° -
0.0592 V
log Q
n
Calculate E:
0.0592 V
[Cu2+][Fe2+]2
E = E° log
n
[Fe3+]2
0.0592 V
(0.25)(0.20)2
= 0.43 V log
2
(1.0 x 10-4)2
E = 0.25 V
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Chapter 17/25
Standard Cell Potentials and
Equilibrium Constants
Using DG° = -nFE°
DG° = -RT ln K
and
-nFE° = -RT ln K
E° =
RT
ln K =
2.303 RT
nF
E° =
log K
nF
0.0592 V
log K
in volts, at 25°C
n
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Chapter 17/26
Standard Cell Potentials and
Equilibrium Constants
Standard Cell Potentials and
Equilibrium Constants
Three methods to determine equilibrium constants:
1. K from concentration data:
K=
[C]c[D]d
[A]a[B]b
-DG°
2. K from thermochemical data: ln K =
RT
3. K from electrochemical data: E° = RT ln K
nF
or
nFE°
ln K =
RT
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Chapter 17/28
Batteries
Lead Storage Battery
Anode:
Cathode:
Overall:
Pb(s) + HSO41-(aq)
PbO2(s) + 3H1+(aq) + HSO41-(aq) + 2ePb(s) + PbO2(s) + 2H1+(aq) + 2HSO41-(aq)
PbSO4(s) + H1+(aq) + 2ePbSO4(s) + 2H2O(l)
2PbSO4(s) + 2H2O(l)
Batteries
Dry-Cell Batteries
Leclanché cell
Anode:
Zn(s)
Cathode: 2MnO2(s) + 2NH41+(aq) + 2e-
Zn2+(aq) + 2e-
Mn2O3(s) + 2NH3(aq)+ H2O(l)
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Chapter 17/30
Batteries
Dry-Cell Batteries
Alkaline cell
Anode:
Zn(s) + 2OH1-(aq)
Cathode: 2MnO2(s) + H2O(l) + 2e-
ZnO(s) + H2O(l) + 2eMn2O3(s) + 2OH1-(aq)
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Chapter 17/31
Batteries
Nickel-Cadmium (“ni-cad”) Batteries
Anode:
Cd(s) + 2OH1-(aq)
Cathode: NiO(OH)(s) + H2O(l) + e-
Cd(OH)2(s) + 2eNi(OH)2(s) + OH1-(aq)
Nickel-Metal Hydride (“NiMH”) Batteries
Anode:
Cathode:
Overall:
MHab(s) + OH1-(aq)
NiO(OH)(s) + H2O(l) + eMHab(s) + NiO(OH)(s)
M(s) + H2O(l) + eNi(OH)2(s) + OH1-(aq)
M(s) + Ni(OH)2(s)
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Chapter 17/32
Batteries
Lithium and Lithium Ion Batteries
Lithium
Anode:
Cathode:
xLi(s)
MnO2(s) + xLi1+(soln) + xe-
xLi1+(soln) + xe-
LixMnO2(s)
Lithium Ion
Anode:
LixC6(s)
Cathode: Li1-xCoO2(s) + xLi1+(soln) + xe-
xLi1+(soln) + 6C(s) + xeLiCoO2(s)
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Chapter 17/33
Fuel Cells
Hydrogen-Oxygen Fuel Cell
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Chapter 17/34
Corrosion
Corrosion: The oxidative deterioration of a metal.
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Chapter 17/35
Corrosion
Prevention of Corrosion
For some metals, oxidation protects the metal (aluminum,
chromium, magnesium, titanium, zinc, and others). For
other metals, there are two main techniques.
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Chapter 17/36
Corrosion
Prevention of Corrosion
1. Galvanization: The coating of iron with zinc.
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Chapter 17/37
Corrosion
Prevention of Corrosion
1. Galvanization: The coating of iron with zinc.
When some of the iron is oxidized (rust), the process
is reversed since zinc will reduce Fe2+ to Fe:
Fe2+(aq) + 2e-
Fe(s)
E° = -0.45 V
Zn2+(aq) + 2e-
Zn(s)
E° = -0.76 V
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Chapter 17/38
Corrosion
Prevention of Corrosion
2. Cathodic Protection: Instead of coating the entire
surface of the first metal with a second metal, the
second metal is placed in electrical contact with the
first metal:
Anode:
Mg(s)
Cathode: O2(g) + 4H1+(aq) + 4e-
Mg2+(aq) + 2e- E° = 2.37 V
2H2O(l)
E° = 1.23 V
Attaching a magnesium stake to iron will corrode the
magnesium instead of the iron. Magnesium acts as a
sacrificial anode.
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Chapter 17/39
Electrolysis and Electrolytic
Cells
Electrolysis: The process of using an electric current
to bring about chemical change.
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Chapter 17/40
Electrolysis and Electrolytic
Cells
Electrolysis of Molten Sodium Chloride
Anode:
Cathode:
Overall:
2Cl1-(l)
2Na1+(l) + 2e-
2Na1+(l) + 2Cl1-(l)
Cl2(g) + 2e2Na(l)
2Na(l) + Cl2(g)
Electrolysis and Electrolytic
Cells
Electrolysis of Aqueous Sodium Chloride
Anode:
2Cl1-(aq)
Cathode:
Overall:
2H2O(l) + 2e-
2Cl1-(l) + 2H2O(l)
Cl2(g) + 2eH2(g) + 2OH1-(aq)
Cl2(g) + H2(g) + 2OH1-(aq)
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Chapter 17/42
Electrolysis and Electrolytic
Cells
Electrolysis of Water
Anode:
2H2O(l)
O2(g) + 4H1+(aq) + 4e-
Cathode: 4H2O(l) + 4e-
2H2(g) + 4OH1-(aq)
Overall:
2H2(g) + O2(g) + 4H1+ + 4OH1-(aq)
6H2O(l)
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Chapter 17/43
Commercial Applications of
Electrolysis
Down’s Cell for the Production of Sodium Metal
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Chapter 17/44
Commercial Applications of
Electrolysis
A Membrane Cell for Electrolytic Production of Cl2 and NaOH
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Chapter 17/45
Commercial Applications of
Electrolysis
Hall-Heroult Process for the Production of Aluminum
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Chapter 17/46
Commercial Applications of
Electrolysis
Electrorefining of copper metal
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Chapter 17/47
Quantitative Aspects of
Electrolysis
Charge(C) = Current(A) x Time(s)
Moles of
e-
1
mol
e
= Charge(A) x
96,500 C
Faraday constant
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Chapter 17/48