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CHEM 160 General Chemistry II
Lecture Presentation
Electrochemistry
December 1, 2004
Chapter 20
Electrochemistry
 Electrochemistry
 deals with interconversion between chemical and
electrical energy
Electrochemistry
 Electrochemistry
 deals with the interconversion between chemical and
electrical energy
 involves redox reactions
Electrochemistry
 Electrochemistry
 deals with interconversion between chemical and
electrical energy
 involves redox reactions
• electron transfer reactions
•Oh No! They’re back!
Redox reactions (quick review)
 Oxidation
 Reduction
 Reducing agent
 Oxidizing agent
Redox reactions (quick review)
 Oxidation
 loss of electrons
 Reduction
 Reducing agent
 Oxidizing agent
Redox reactions (quick review)
 Oxidation
 loss of electrons
 Reduction
 gain of electrons
 Reducing agent
 Oxidizing agent
Redox reactions (quick review)
 Oxidation
 loss of electrons
 Reduction
 gain of electrons
 Reducing agent
 donates the electrons and is oxidized
 Oxidizing agent
Redox reactions (quick review)
 Oxidation
 loss of electrons
 Reduction
 gain of electrons
 Reducing agent
 donates the electrons and is oxidized
 Oxidizing agent
 accepts electrons and is reduced
Redox Reactions
 Direct redox reaction
Redox Reactions
 Direct redox reaction
 Oxidizing and reducing agents are mixed together
Direct Redox Reaction
Zn rod
CuSO4(aq)
(Cu2+)
Direct Redox Reaction
Zn rod
CuSO4(aq)
(Cu2+)
Deposit of
Cu metal
forms
Redox Reactions
 Direct redox reaction
 Oxidizing and reducing agents are mixed together
 Indirect redox reaction
 Oxidizing and reducing agents are separated but
connected electrically
• Example
– Zn and Cu2+ can be reacted indirectly
 Basis for electrochemistry
– Electrochemical cell
Electrochemical Cells
Electrochemical Cells
 Voltaic Cell
 cell in which a spontaneous redox reaction generates
electricity
 chemical energy  electrical energy
Electrochemical Cells
Electrochemical Cells
Voltaic Cell
Electrochemical Cells
 Electrolytic Cell
 electrochemical cell in which an electric current
drives a nonspontaneous redox reaction
 electrical energy  chemical energy
Cell Potential
Cell Potential
 Cell Potential (electromotive force), Ecell (V)
 electrical potential difference between the two
electrodes or half-cells
• Depends on specific half-reactions, concentrations, and
temperature
• Under standard state conditions ([solutes] = 1 M, Psolutes =
1 atm), emf = standard cell potential, Ecell
• 1 V = 1 J/C
 driving force of the redox reaction
Cell Potential
high electrical
potential
low electrical
potential
Cell Potential
Ecell = Ecathode - Eanode = Eredn - Eox
E°cell = E°cathode - E°anode = E°redn - E°ox
(Ecathode and Eanode are reduction potentials by definition.)
Cell Potential
 E°cell = E°cathode - E°anode = E°redn - E°ox
 Ecell can be measured
• Absolute Ecathode and Eanode values cannot
 Reference electrode
 has arbitrarily assigned E
 used to measure relative Ecathode and Eanode for halfcell reactions
 Standard hydrogen electrode (S.H.E.)
 conventional reference electrode
Standard Hydrogen Electrode
 E = 0 V (by
definition; arbitrarily
selected)
 2H+ + 2e-  H2
Example 1
A voltaic cell is made by connecting a standard
Cu/Cu2+ electrode to a S.H.E. The cell potential
is 0.34 V. The Cu electrode is the cathode.
What is the standard reduction potential of the
Cu/Cu2+ electrode?
Example 2
A voltaic cell is made by connecting a standard
Zn/Zn2+ electrode to a S.H.E. The cell potential
is 0.76 V. The Zn electrode is the anode of the
cell. What is the standard reduction potential of
the Zn/Zn2+ electrode?
Standard Electrode Potentials
 Standard Reduction Potentials, E°
 E°cell measured relative to S.H.E. (0 V)
• electrode of interest = cathode
 If E° < 0 V:
• Oxidizing agent is harder to reduce than H+
 If E° > 0 V:
• Oxidizing agent is easier to reduce than H+
Standard Reduction Potentials
Reduction Half-Reaction
E(V)
F2(g) + 2e-  2F-(aq)
2.87
Au3+(aq) + 3e-  Au(s)
1.50
Cl2(g) + 2 e-  2Cl-(aq)
1.36
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O
1.33
O2(g) + 4H+ + 4e-  2H2O(l)
1.23
Ag+(aq) + e-  Ag(s)
0.80
Fe3+(aq) + e-  Fe2+(aq)
0.77
Cu2+(aq) + 2e-  Cu(s)
0.34
Sn4+(aq) + 2e-  Sn2+(aq)
0.15
2H+(aq) + 2e-  H2(g)
0.00
Sn2+(aq) + 2e-  Sn(s)
-0.14
Ni2+(aq) + 2e-  Ni(s)
-0.23
Fe2+(aq) + 2e-  Fe(s)
-0.44
Zn2+(aq) + 2e-  Zn(s)
-0.76
Al3+(aq) + 3e-  Al(s)
-1.66
Mg2+(aq) + 2e-  Mg(s)
-2.37
Li+(aq) + e-  Li(s)
-3.04
Uses of Standard Reduction
Potentials
 Compare strengths of reducing/oxidizing agents.
 the more - E°, stronger the red. agent
 the more + E°, stronger the ox. agent
Reduction Half-Reaction
E(V)
F2(g) + 2e-  2F-(aq)
2.87
Au3+(aq) + 3e-  Au(s)
1.50
Cl2(g) + 2 e-  2Cl-(aq)
1.36
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O
1.33
O2(g) + 4H+ + 4e-  2H2O(l)
1.23
Ag+(aq) + e-  Ag(s)
0.80
Fe3+(aq) + e-  Fe2+(aq)
0.77
Cu2+(aq) + 2e-  Cu(s)
0.34
Sn4+(aq) + 2e-  Sn2+(aq)
0.15
2H+(aq) + 2e-  H2(g)
0.00
Sn2+(aq) + 2e-  Sn(s)
-0.14
Ni2+(aq) + 2e-  Ni(s)
-0.23
Fe2+(aq) + 2e-  Fe(s)
-0.44
Zn2+(aq) + 2e-  Zn(s)
-0.76
Al3+(aq) + 3e-  Al(s)
-1.66
Mg2+(aq) + 2e-  Mg(s)
-2.37
Li+(aq) + e-  Li(s)
-3.04
Red. agent strength increases
Ox. agent strength increases
Standard Reduction Potentials
Uses of Standard Reduction
Potentials
 Determine if oxidizing and reducing agent react
spontaneously
 diagonal rule
ox. agent
red. agent
Uses of Standard Reduction
Potentials
 Determine if oxidizing and reducing agent react
spontaneously
more +
Cathode
(reduction)
Anode
(oxidation)
more -
Standard Reduction Potentials
Reduction Half-Reaction
E(V)
F2(g) + 2e-  2F-(aq)
2.87
Au3+(aq) + 3e-  Au(s)
1.50
Cl2(g) + 2 e-  2Cl-(aq)
1.36
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O
1.33
O2(g) + 4H+ + 4e-  2H2O(l)
1.23
Ag+(aq) + e-  Ag(s)
0.80
Fe3+(aq) + e-  Fe2+(aq)
0.77
Cu2+(aq) + 2e-  Cu(s)
0.34
Sn4+(aq) + 2e-  Sn2+(aq)
0.15
2H+(aq) + 2e-  H2(g)
0.00
Sn2+(aq) + 2e-  Sn(s)
-0.14
Ni2+(aq) + 2e-  Ni(s)
-0.23
Fe2+(aq) + 2e-  Fe(s)
-0.44
Zn2+(aq) + 2e-  Zn(s)
-0.76
Al3+(aq) + 3e-  Al(s)
-1.66
Mg2+(aq) + 2e-  Mg(s)
-2.37
Li+(aq) + e-  Li(s)
-3.04
Uses of Standard Reduction
Potentials
 Calculate E°cell
 E°cell = E°cathode - E°anode
• Greater E°cell, greater the driving force
 E°cell > 0 : spontaneous redox reactions
 E°cell < 0 : nonspontaeous redox reactions
Example 3
A voltaic cell consists of a Ag electrode in 1.0 M
AgNO3 and a Cu electrode in 1 M Cu(NO3)2.
Calculate E°cell for the spontaneous cell reaction
at 25°C.
Standard Reduction Potentials
Reduction Half-Reaction
E(V)
F2(g) + 2e-  2F-(aq)
2.87
Au3+(aq) + 3e-  Au(s)
1.50
Cl2(g) + 2 e-  2Cl-(aq)
1.36
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O
1.33
O2(g) + 4H+ + 4e-  2H2O(l)
1.23
Ag+(aq) + e-  Ag(s)
0.80
Fe3+(aq) + e-  Fe2+(aq)
0.77
Cu2+(aq) + 2e-  Cu(s)
0.34
Sn4+(aq) + 2e-  Sn2+(aq)
0.15
2H+(aq) + 2e-  H2(g)
0.00
Sn2+(aq) + 2e-  Sn(s)
-0.14
Ni2+(aq) + 2e-  Ni(s)
-0.23
Fe2+(aq) + 2e-  Fe(s)
-0.44
Zn2+(aq) + 2e-  Zn(s)
-0.76
Al3+(aq) + 3e-  Al(s)
-1.66
Mg2+(aq) + 2e-  Mg(s)
-2.37
Li+(aq) + e-  Li(s)
-3.04
Example 4
A voltaic cell consists of a Ni electrode in 1.0 M
Ni(NO3)2 and an Fe electrode in 1 M Fe(NO3)2.
Calculate E°cell for the spontaneous cell reaction
at 25°C.
Standard Reduction Potentials
Reduction Half-Reaction
E(V)
F2(g) + 2e-  2F-(aq)
2.87
Au3+(aq) + 3e-  Au(s)
1.50
Cl2(g) + 2 e-  2Cl-(aq)
1.36
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O
1.33
O2(g) + 4H+ + 4e-  2H2O(l)
1.23
Ag+(aq) + e-  Ag(s)
0.80
Fe3+(aq) + e-  Fe2+(aq)
0.77
Cu2+(aq) + 2e-  Cu(s)
0.34
Sn4+(aq) + 2e-  Sn2+(aq)
0.15
2H+(aq) + 2e-  H2(g)
0.00
Sn2+(aq) + 2e-  Sn(s)
-0.14
Ni2+(aq) + 2e-  Ni(s)
-0.23
Fe2+(aq) + 2e-  Fe(s)
-0.44
Zn2+(aq) + 2e-  Zn(s)
-0.76
Al3+(aq) + 3e-  Al(s)
-1.66
Mg2+(aq) + 2e-  Mg(s)
-2.37
Li+(aq) + e-  Li(s)
-3.04
Cell Potential
 Is there a relationship between Ecell and DG for a
redox reaction?
Cell Potential
 Relationship between Ecell and DG:
 DG = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn.
Cell Potential
 Relationship between Ecell and DG:
 DG = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn.
• 1 J = CV
• DG < 0, Ecell > 0 = spontaneous
Equilibrium Constants from Ecell
 Relationship between Ecell and DG:
 DG = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn
• 1 J = CV
• DG < 0, Ecell > 0 = spontaneous
 Under standard state conditions:
 DG° = -nFE°cell
Equilibrium Constants from Ecell
 Relationship between Ecell and DG:
 DG = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s
transferred redox rxn
• 1 J = CV
• DG < 0, Ecell > 0 = spontaneous
 Under standard state conditions:
 DG° = -nFE°cell
Equilibrium Constants from Ecell
 Relationship between Ecell and DG:
 DG = -nFEcell
• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox
rxn
• 1 J = CV
• DG < 0, Ecell > 0 = spontaneous
 Under standard state conditions:
 DG° = -nFE°cell
and
 DG° = -RTlnK
so
 -nFE°cell = -RTlnK
Calorimetric Data
DH° DS°
Composition
Data
DG°
Electrochemical
Data
E°cell
Equilibrium
constants
K
Example 5
Calculate E°cell, DG°, and K for the voltaic cell
that uses the reaction between Ag and Cl2 under
standard state conditions at 25°C.
The Nernst Equation
 DG depends on concentrations
 DG = DG° + RTlnQ
and
 DG = -nFEcell and DG° = -nFE°cell
thus
 -nFEcell = -nFE°cell + RTlnQ
or
 Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)
The Nernst Equation
 Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)
 At 298 K (25°C), RT/F = 0.0257 V
so
 Ecell = E°cell - (0.0257/n)lnQ
or
 Ecell = E°cell - (0.0592/n)logQ
Example 7
 Calculate the voltage produced by the galvanic
cell which uses the reaction below if [Ag+] =
0.001 M and [Cu2+] = 1.3 M.
2Ag+(aq) + Cu(s)  2Ag(s) + Cu2+(aq)
Reduction Half-Reaction
E(V)
F2(g) + 2e-  2F-(aq)
2.87
Au3+(aq) + 3e-  Au(s)
1.50
Cl2(g) + 2 e-  2Cl-(aq)
1.36
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O
1.33
O2(g) + 4H+ + 4e-  2H2O(l)
1.23
Ag+(aq) + e-  Ag(s)
0.80
Fe3+(aq) + e-  Fe2+(aq)
0.77
Cu2+(aq) + 2e-  Cu(s)
0.34
Sn4+(aq) + 2e-  Sn2+(aq)
0.15
2H+(aq) + 2e-  H2(g)
0.00
Sn2+(aq) + 2e-  Sn(s)
-0.14
Ni2+(aq) + 2e-  Ni(s)
-0.23
Fe2+(aq) + 2e-  Fe(s)
-0.44
Zn2+(aq) + 2e-  Zn(s)
-0.76
Al3+(aq) + 3e-  Al(s)
-1.66
Mg2+(aq) + 2e-  Mg(s)
-2.37
Li+(aq) + e-  Li(s)
-3.04
Red. agent strength increases
Ox. agent strength increases
Standard Reduction Potentials
Commercial Voltaic Cells
 Battery
 commercial voltaic cell used as portable source of
electrical energy
 types
 primary cell
• Nonrechargeable
• Example: Alkaline battery
 secondary cell
• Rechargeable
• Example: Lead storage battery
How Does a Battery Work
Assume a generalized battery
Seal/cap
cathode (+)
Electrolyte
Paste
anode (-)
Battery
Placing the battery into a flashlight,
etc., and turning the power on
completes the circuit and allows
electron flow to occur
Electrolyte paste:
ion migration occurs
here
cathode (+):
Reduction occurs
here
anode (-):
oxidation
occurs here
e- flow
How Does a Battery Work
 Battery reaction when producing electricity
(spontaneous):
Cathode: O1 + e-  R1
Anode: R2  O2 + eOverall: O1 + R2  R1 + O2
 Recharging a secondary cell
 Redox reaction must be reversed, i.e., current is
reversed (nonspontaneous)
Recharge: O2 + R1  R2 + O1
 Performed using electrical energy from an external
power source
Batteries
 Read the textbook to fill in the details on
specific batteries.
 Alkaline battery
 Lead storage battery
 Nicad battery
 Fuel cell
Corrosion
 Corrosion
 deterioration of metals by a spontaneous redox
reaction
• Attacked by species in environment
– Metal becomes a “voltaic” cell
• Metal is often lost to a solution as an ion
 Rusting of Iron
Corrosion of Iron
Corrosion of Iron
Half-reactions
anode: Fe(s)  Fe2+(aq) + 2ecathode: O2(g) + 4H+(aq) + 4e-  2H2O(l)
overall: 2Fe(s) + O2(g) + 4H+(aq) 
2Fe2+(aq) + 2H2O(l)
Ecell > 0 (Ecell = 0.8 to 1.2 V), so process is
spontaneous!
Corrosion of Iron
Rust formation:
4Fe2+(aq) + O2(g) + 4H+(aq)  4Fe3+(aq) + 2H2O(l)
2Fe3+(aq) + 4H2O(l)  Fe2O3H2O(s) + 6H+(aq)
Prevention of Corrosion
 Cover the Fe surface with a protective coating
 Paint
 Passivation
• surface atoms made inactive via oxidation
2Fe(s) + 2Na2CrO4(aq) + 2H2O(l) -->
Fe2O3(s) + Cr2O3(s) + 4NaOH(aq)
 Other metal
• Tin
• Zn
– Galvanized iron
Prevention of Corrosion
 Cathodic Protection
 metal to be protected is brought into contact with a
more easily oxidized metal
 “sacrificial” metal becomes the anode
• “Corrodes” preferentially over the iron
• Iron serves only as the cathode
Standard Electrode Potentials
Half-reaction
F2(g) + 2e- -> 2F-(aq)
Ag+(aq) + e- -> Ag(s)
Cu2+(aq) + 2e- -> Cu(s)
2H+(aq) + 2e- -> H2(g)
Ni2+(aq) + 2e- -> Ni(s)
Fe2+(aq) + 2e- -> Fe(s)
Zn2+(aq) + 2e- -> Zn(s)
Al3+(aq) + 3e- -> Al(s)
Mg2+(aq) + 2e- ->Mg(s)
E°
+2.87 V
+0.80 V
+0.34 V
0V
-0.25 V
-0.44 V
-0.76 V
Metals more
-1.66 V easily oxidized
-2.38 V
than Fe have
more negative
E°’s
Cathodic Protection
galvanized steel (Fe)
Cathodic Protection
(anode)
(cathode)
(electrolyte)
Electrolysis
 Electrolysis
 process in which electrical energy drives a
nonspontaneous redox reaction
• electrical energy is converted into chemical energy
 Electrolytic cell
 electrochemical cell in which an electric current
drives a nonspontaneous redox reaction
Electrolysis
 Same principles apply to both electrolytic and
voltaic cells
 oxidation occurs at the anode
 reduction occurs at the cathode
 electrons flow from anode to cathode in the external
circuit
• In an electrolytic cell, an external power source pumps the
electrons through the external circuit
Electrolysis of Molten NaCl
Quantitative Aspects of Electrochemical Cells
 For any half-reaction, the amount of a substance
oxidized or reduced at an electrode is proportional to
the number of electrons passed through the cell
 Faraday’s law of electrolysis
 Examples
• Na+ + 1e-  Na
• Al3+ + 3e-  Al
 Number of electrons passing through cell is measured by
determining the quantity of charge (coulombs) that has
passed
• 1 C= 1Ax 1 s
• 1 F = 1 mole e- = 96500 C
Steps for Quantitative Electrolysis
Calculations
current (A) and time
(s), A x s
Number of
moles of e-
charge in
coulombs
(C)
moles of substance
oxidized or reduced
mass of substance
oxidized or reduced
Example 8
 What mass of copper metal can be produced by
a 3.00 A current flowing through a copper(II)
sulfate (CuSO4) solution for 5.00 hours?
Example 9
 An aqueous solution of an iron salt is
electrolyzed by passing a current of 2.50 A for
3.50 hours. As a result, 6.1 g of iron metal are
formed at the cathode. Calculate the charge on
the iron ions in the solution.