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Prentice-Hall © 2007General Chemistry: Chapter 20Slide 1 of 54
CHEMISTRYNinth
Edition GENERAL
Principles and Modern Applications
Petrucci • Harwood • Herring • Madura
Chapter 20: Electrochemistry
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 2 of 54
Contents
20-1 Electrode Potentials and Their Measurement
20-2 Standard Electrode Potentials
20-3 Ecell, ΔG, and Keq
20-4 Ecell as a Function of Concentration
20-5 Batteries: Producing Electricity Through Chemical Reactions
20-6 Corrosion: Unwanted Voltaic Cells
20-7 Electrolysis: Causing Non-spontaneous Reactions to Occur
20-8 Industrial Electrolysis Processes
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 3 of 54
20-1 Electrode Potentials and Their Measurement
Cu(s) + 2Ag+(aq)
Cu2+(aq) + 2 Ag(s)
Cu(s) + Zn2+(aq)
No reaction
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 4 of 54
An Electrochemical Half Cell
Anode
Cathode
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 5 of 54
An Electrochemical Cell
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 6 of 54
Terminology
Electromotive force, Ecell.
The cell voltage or cell potential.
Cell diagram. Shows the components of the cell in a symbolic way. Anode (where oxidation occurs) on the left. Cathode (where reduction occurs) on the right.
◦ Boundary between phases shown by |.
◦ Boundary between half cells (usually a salt bridge) shown by ||.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 7 of 54
Terminology
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Ecell = 1.103 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 8 of 54
Terminology
Galvanic cells. Produce electricity as a result of spontaneous reactions.
Electrolytic cells. Non-spontaneous chemical change driven by electricity.
Couple, M|Mn+
A pair of species related by a change in number of e-.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 9 of 54
20-2 Standard Electrode Potentials
Cell voltages, the potential differences between electrodes, are among the most precise scientific measurements.
The potential of an individual electrode is difficult to establish.
Arbitrary zero is chosen.
The Standard Hydrogen Electrode (SHE)
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 10 of 54
Standard Hydrogen Electrode2 H+(a = 1) + 2 e- H2(g, 1 bar) E° = 0 V
Pt|H2(g, 1 bar)|H+(a = 1)
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 11 of 54
Standard Electrode Potential, E°
E° defined by international agreement. The tendency for a reduction process to occur at
an electrode. All ionic species present at a=1 (approximately 1 M). All gases are at 1 bar (approximately 1 atm). Where no metallic substance is indicated, the potential
is established on an inert metallic electrode (ex. Pt).
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 12 of 54
Reduction Couples
Cu2+(1M) + 2 e- → Cu(s) E°Cu2+/Cu = ?
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = 0.340 V
Standard cell potential: the potential difference of a cell formed from two standard electrodes.
E°cell = E°cathode - E°anode
cathodeanode
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 13 of 54
Standard Cell Potential
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = 0.340 V
E°cell = E°cathode - E°anode
E°cell = E°Cu2+/Cu - E°H+/H2
0.340 V = E°Cu2+/Cu - 0 V
E°Cu2+/Cu = +0.340 V
H2(g, 1 atm) + Cu2+(1 M) → H+(1 M) + Cu(s) E°cell = 0.340 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 14 of 54
Measuring Standard Reduction Potential
anodeanode cathode cathode
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 15 of 54
Standard Reduction Potentials
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 16 of 54
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 17 of 54
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 18 of 54
20-3 Ecell, ΔG, and Keq
Cells do electrical work. Moving electric charge.
Faraday constant, F = 96,485 C mol-1
elec = -nFE
ΔG = -nFE
ΔG° = -nFE°
Michael Faraday 1791-1867
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 19 of 54
Combining Half Reactions
Fe3+(aq) + 3e- → Fe(s) E°Fe3+/Fe = ?
Fe2+(aq) + 2e- → Fe(s) E°Fe2+/Fe = -0.440 V
Fe3+(aq) + 3e- → Fe2+(aq) E°Fe3+/Fe2+ = 0.771 V
Fe3+(aq) + 3e- → Fe(s)
ΔG° = +0.880 J
ΔG° = -0.771 J
ΔG° = +0.109 VE°Fe3+/Fe = +0.331 V
ΔG° = +0.109 V = -nFE°
E°Fe3+/Fe = +0.109 V /(-3F) = -0.0363 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 20 of 54
Spontaneous Change
ΔG < 0 for spontaneous change. Therefore E°cell > 0 because ΔGcell = -nFE°cell
E°cell > 0 Reaction proceeds spontaneously as written.
E°cell = 0 Reaction is at equilibrium.
E°cell < 0 Reaction proceeds in the reverse direction spontaneously.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 21 of 54
The Behavior or Metals Toward Acids
M(s) → M2+(aq) + 2 e- E° = -E°M2+/M
2 H+(aq) + 2 e- → H2(g) E°H+/H2 = 0 V
2 H+(aq) + M(s) → H2(g) + M2+(aq)
E°cell = E°H+/H2 - E°M2+/M = -E°M2+/M
When E°M2+/M < 0, E°cell > 0. Therefore ΔG° < 0.
Metals with negative reduction potentials react with acids.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 22 of 54
Relationship Between E°cell and Keq
ΔG° = -RT ln Keq = -nFE°cell
E°cell = nF
RTln Keq
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 23 of 54
Summary of Thermodynamic, Equilibrium and Electrochemical Relationships.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 24 of 54
20-4 Ecell as a Function of Concentration
ΔG = ΔG° -RT ln Q
-nFEcell = -nFEcell° -RT ln Q
Ecell = Ecell° - ln QnF
RT
Convert to log10 and calculate constants.
Ecell = Ecell° - log Qn
0.0592 VThe Nernst Equation:
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 25 of 54
Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)
Applying the Nernst Equation for Determining Ecell. What is the value of Ecell for the voltaic cell pictured below and diagrammed as follows?
EXAMPLE 20-8
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 26 of 54
Ecell = Ecell° - log Qn
0.0592 V
Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s)
Ecell = Ecell° - logn
0.0592 V [Fe3+][Fe2+] [Ag+]
Fe2+(aq) + Ag+(aq) → Fe3+(aq) + Ag (s)
Ecell = 0.029 V – 0.018 V = 0.011 V
EXAMPLE 20-8
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 27 of 54
Concentration Cells
Two half cells with identical electrodes but different ion concentrations.
2 H+(1 M) → 2 H+(x M)
Pt|H2 (1 atm)|H+(x M)||H+(1.0 M)|H2(1 atm)|Pt(s)
2 H+(1 M) + 2 e- → H2(g, 1 atm)
H2(g, 1 atm) → 2 H+(x M) + 2 e-
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 28 of 54
Concentration Cells
Ecell = Ecell° - logn
0.0592 V x2
12
Ecell = 0 - log2
0.0592 V x2
1
Ecell = - 0.0592 V log x
Ecell = (0.0592 V) pH
2 H+(1 M) → 2 H+(x M)Ecell = Ecell° - log Qn
0.0592 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 29 of 54
Measurement of Ksp
Ag+(0.100 M) → Ag+(sat’d M)
Ag|Ag+(sat’d AgI)||Ag+(0.10 M)|Ag(s)
Ag+(0.100 M) + e- → Ag(s)
Ag(s) → Ag+(sat’d) + e-
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 30 of 54
Using a Voltaic Cell to Determine Ksp of a Slightly Soluble Solute. With the date given for the reaction on the previous slide, calculate Ksp for AgI.
AgI(s) → Ag+(aq) + I-(aq)
Let [Ag+] in a saturated Ag+ solution be x:
Ag+(0.100 M) → Ag+(sat’d M)
Ecell = Ecell° - log Q = n
0.0592 VEcell° - log
n
0.0592 V
[Ag+]0.10 M soln
[Ag+]sat’d AgI
EXAMPLE 20-10
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 31 of 54
Ecell = Ecell° - log n
0.0592 V
[Ag+]0.10 M soln
[Ag+]sat’d AgI
Ecell = Ecell° - log n
0.0592 V
0.100
x
0.417 =0 - (log x – log 0.100) 1
0.0592 V
0.417log 0.100 -
0.0592log x = = -1 – 7.04 = -8.04
x = 10-8.04 = 9.110-9 Ksp = x2 = 8.310-17
EXAMPLE 20-10
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 32 of 54
20-5 Batteries: Producing Electricity Through Chemical Reactions
Primary Cells (or batteries). Cell reaction is not reversible.
Secondary Cells. Cell reaction can be reversed by passing electricity
through the cell (charging).
Flow Batteries and Fuel Cells. Materials pass through the battery which converts
chemical energy to electric energy.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 33 of 54
The Leclanché (Dry) Cell
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 34 of 54
Dry Cell
Zn(s) → Zn2+(aq) + 2 e-Oxidation:
2 MnO2(s) + H2O(l) + 2 e- → Mn2O3(s) + 2 OH-Reduction:
NH4+ + OH- → NH3(g) + H2O(l) Acid-base reaction:
NH3 + Zn2+(aq) + Cl- → [Zn(NH3)2]Cl2(s)Precipitation reaction:
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 35 of 54
Alkaline Dry Cell
Zn2+(aq) + 2 OH- → Zn (OH)2(s)
Zn(s) → Zn2+(aq) + 2 e-
Oxidation reaction can be thought of in two steps:
2 MnO2(s) + H2O(l) + 2 e- → Mn2O3(s) + 2 OH-Reduction:
Zn (s) + 2 OH- → Zn (OH)2(s) + 2 e-
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 36 of 54
Lead-Acid (Storage) Battery
The most common secondary battery.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 37 of 54
Lead-Acid Battery
PbO2(s) + 3 H+(aq) + HSO4-(aq) + 2 e- → PbSO4(s) + 2 H2O(l)
Oxidation:
Reduction:
Pb (s) + HSO4-(aq) → PbSO4(s) + H+(aq) + 2 e-
PbO2(s) + Pb(s) + 2 H+(aq) + HSO4-(aq) → 2 PbSO4(s) + 2 H2O(l)
E°cell = E°PbO2/PbSO4 - E°PbSO4/Pb = 1.74 V – (-0.28 V) = 2.02 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 38 of 54
The Silver-Zinc Cell: A Button Battery
Zn(s),ZnO(s)|KOH(sat’d)|Ag2O(s),Ag(s)
Zn(s) + Ag2O(s) → ZnO(s) + 2 Ag(s) Ecell = 1.8 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 39 of 54
The Nickel-Cadmium Cell
Cd(s) + 2 NiO(OH)(s) + 2 H2O(L) → 2 Ni(OH)2(s) + Cd(OH)2(s)
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 40 of 54
Fuel Cells
O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq)
2{H2(g) + 2 OH-(aq) → 2 H2O(l) + 2 e-}
2H2(g) + O2(g) → 2 H2O(l)
E°cell = E°O2/OH- - E°H2O/H2
= 0.401 V – (-0.828 V) = 1.229 V
= ΔG°/ ΔH° = 0.83
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 41 of 54
Air Batteries
4 Al(s) + 3 O2(g) + 6 H2O(l) + 4 OH- → 4 [Al(OH)4](aq)
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 42 of 54
20-6 Corrosion: Unwanted Voltaic Cells
O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq)
2 Fe(s) → 2 Fe2+(aq) + 4 e-
2 Fe(s) + O2(g) + 2 H2O(l) → 2 Fe2+(aq) + 4 OH-(aq)
Ecell = 0.841 V
EO2/OH- = 0.401 V
EFe/Fe2+ = -0.440 V
In neutral solution:
In acidic solution:
O2(g) + 4 H+(aq) + 4 e- → 4 H2O (aq) EO2/OH- = 1.229 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 43 of 54
Corrosion
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 44 of 54
Corrosion Protection
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 45 of 54
Corrosion Protection
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 46 of 54
20-7 Electrolysis: Causing Non-spontaneous Reactions to Occur
Galvanic Cell:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) EO2/OH- = 1.103 V
Electolytic Cell:
Zn2+(aq) + Cu(s) → Zn(s) + Cu2+(aq) EO2/OH- = -1.103 V
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 47 of 54
Predicting Electrolysis Reaction
An Electrolytic Cell e- is the reverse of the
voltaic cell. Battery must have a
voltage in excess of 1.103 V in order to force the non-spontaneous reaction.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 48 of 54
Complications in Electrolytic Cells
Overpotential. Competing reactions. Non-standard states. Nature of electrodes.
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 49 of 54
Quantitative Aspects of Electrolysis
1 mol e- = 96485 C
Charge (C) = current (C/s) time (s)
ne- = I t
F
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 50 of 54
20-8 Industrial Electrolysis Processes
Prentice-Hall © 2007General Chemistry: Chapter 20Slide 51 of 54
Chlor-Alkali Process
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