Chem 1B
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Chapter 18 - Electrochemistry
18.2 Galvanic Cells (aka Voltaic Cells)
A. Oxidation-Reduction Reactions
B. Galvanic Cell (Voltaic Cells) –
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a. Example of a galvanic cell with Zn as the anode and Cu as the cathode
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b. Example of a galvanic cell with the standard hydrogen electrode (SHE)
C. Cell Potential (Ecell)
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18.3 Standard Reduction Potentials
A. Standard Reduction Potentials (E°cell) –
Selected Standard Reduction Potentials (298K)
Half-Reaction
E° half cell (V) Half-Reaction
E° half cell (V)
3+
+
-1.66
0.00
Al + 3e ⇄ Al
2H + 2e ⇄ H2
2+
2+
-1.18
0.34
Mn + 2e ⇄ Mn
Cu + 2e ⇄ Cu
2+
-0.76
0.40
Zn + 2e ⇄ Zn
2H2O + O2 + 4e ⇄ 4OH
3+
+
-0.74
0.52
Cr + 3e ⇄ Cr
Cu + e ⇄ Cu
2+
3+
-0.44
0.77
Fe + 2e ⇄ Fe
Fe + 3e ⇄ Fe
2+
+
-0.25
0.80
Ni + 2e ⇄ Ni
Ag + e ⇄ Ag
2+
-0.28
1.09
Co + 2e ⇄ Co
Br2 + 2e ⇄ 2Br
2+
3+
2+
Cr
O
+
14H
6e
⇄
2Cr
+
7H
O
-0.14
1.33
2 7
2
Sn + 2e ⇄ Sn
2+
3+
-0.13
1.50
Pb + 2e ⇄ Pb
Au + 3e ⇄ Au
*The more negative the E°half cell, the more likely to be oxidized (stronger reducing agent)
*The more positive the E°half cell, the more likely to be reduced (stronger oxidizing agent)
Example: Rank the following from strongest to weakest oxidizing agent:
Cu+ (aq), Br2(l), Cr2O72- (aq), and Fe3+(aq)
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5 Example: Rank the following from strongest to weakest reducing agent: Sn,
Ni, Cr, and Pb.
Example: Of Al, Ni, and Ag, which would protect Fe from being oxidized to
Fe2+?
B. Standard Cell Potential (E°cell) –
Example – What is the cell potential for a galvanic cell
with Cu as cathode and Zn as anode?
Example: Calculate the cell potential for the reaction:
2 Fe3+(aq) + Cu(s) → 2 Fe2+(aq) + Cu2+(aq)
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18.4 Call Potential, Electrical Work, and Free Energy
A. Work and ΔG°
B. E°cell and K
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Example: Consider the following reaction:
2Ag+ (aq) + Zn (s) → 2Ag (s) + Zn2+(aq)
a. What is the standard cell potential?
b. What is the value of ΔG° at 25°C?
c. What is the value of K at 25°C?
Example: Calculate K at 25 °C for the following reaction:
2 Fe3+(aq) + Cu(s) → 2 Fe2+(aq) + Cu2+(aq) E°cell = +0.43 V
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8 18.5 Dependence of Cell Potential on Concentration
A. Nonstandard Conditions (starting concentrations are NOT 1M or 1atm)
1. Derivation of the Nernst Equation
2. Simplification at 25°C
3. Example: Calculate the cell potential for the reaction below at 25°C
when [H+] = 1.0 M, [Zn2+] = 0.0010 M and the pressure of hydrogen gas is 0.10
atm.
Zn (s) + 2H+ (aq) → Zn2+ (aq) + H2 (g)
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4. Nernst equation shows us that the cell potential depends on the
concentrations of solutions in the cell:
a. Q = 1
b. Q<1
c. Q>1
d. Q = K
B. Concentration Cells –
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10 Example: What is the cell potential for the cell above?
18.6 Batteries (some examples)
A. Alkaline Battery (non-rechargeable)
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11 B. Lead-Acid Battery (car battery, rechargeable)
C. Fuel Cells 18.7 Corrosion –
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A. The Corrosion of iron:
step 1: 2 Fe(s) → 2 Fe2+(aq) + 4e–
step 2: O2(g) + 4 H+(aq) + 4 e– → 2 H2O(l)
step 3: 2 Fe2+(aq) + (n + 2) H2O(l) + 1/2 O2(g) → Fe2O3·nH2O(s) (rust) + 4 H+
Overall: 2 Fe(s) + 3/2 O2(g) + nH2O(l) → Fe2O3·nH2O(s) (rust)
B. The Effect of Metal-Metal Contact on the Corrosion of Iron
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18.8 Electrolysis –
A. Electrolytic vs Galvanic Cells: What’s the difference?
C. Applications of Electrolysis
1. Electroplating
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2. Electrolysis of pure molten salts
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15 B. Stoichiometry in Electrolysis
Example: In an electrolytic cell that plates copper, the cell was operated
for 5.1 hours at 68.1 mA. What is the mass of copper plated?
Example: In an attempt to protect an iron automobile bumper from
corrosion, a technician wants to electroplate chromium metal onto the
bumper.
a. Is chromium metal an appropriate choice to protect the iron
bumper from corrosion? Explain.
b. Could chromium metal be electroplated from a molten mixture of
CrCl3 and ZnCl2?
c. Assuming the technician finds the best way to electroplate
chromium onto the bumper, if he uses 200. A of current and 58.0
minutes is allowed for the process, how much Cr can be plated?