Topic 16: Electrochemistry : Topic 16: Electrochemistry Zumdahl 7e, Chapter 17
II. Standard Reduction Potentials : II. Standard Reduction Potentials Cell Potential
E°cell is the combination of two half-reactions’ potentials (voltage generated/required)
Standard of comparison is the hydrogen electrode: 2 H+ + 2 e- ? H2 E°cell = 0.00 V
SRP table lists half reactions at 298K and 1 atm, all in REDUCTION form
If you need the reverse/oxidation reaction, change the sign
BUT multiplying by coefficients does not change E°cell
Calculating the Cell Potential of a Galvanic Cell
E°cell > 0.00 V in order for the cell to work; that’s the amount of electricity generated by the reaction
E°cell < 0.00 V means that the reaction doesn’t happen spontaneously; you need to put in that much electricity to make the reaction happen
Sample Problem : Sample Problem Calculate the E°cell for the following reaction:
Mg (s) + 2 H+ (aq) ? Mg2+ (aq) + H2 (g)
re: 2 H+ (aq) + 2 e- ? H2 (g) E°cell = 0.00 V
ox: Mg (s) ? Mg2+ (aq) + 2 e- E°cell = -(-2.37) V
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E°cell = 2.37 V
This reaction happens spontaneously.
Sample Problem : Sample Problem Calculate the E°cell for the following reaction:
2 Ag (s) + Cu2+ (aq) ? 2 Ag+ (aq) + Cu (s)
re: Cu2+ (aq) + 2 e- ? Cu (s) E°cell = +0.34 V
ox: 2 Ag (s) ? 2 Ag+ (aq) + 2 e- E°cell = -0.80 V
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E°cell = -0.46 V
This reaction does not happen spontaneously.
Do Example 2.
c. Complete Description of a Galvanic Cell : c. Complete Description of a Galvanic Cell Cell potential: E°cell
Balanced cell reaction
Direction of electron flow (anode to cathode, more active metal to less active metal)
Designate the anode and the cathode
State/phase of each electrode, which ions present
Do Example 3.
III. Cell Potential and Free Energy : III. Cell Potential and Free Energy Spontaneous Reactions and Gibbs Free Energy
?G° = -nFE°cell
(J) = -(mol e- transferred) (96,485 C/mol e-) (J/C)
If ?G < 0 and E°cell > 0 then the reaction is spontaneous
?G° = -RT ln k = -nFE°cell
E°cell = RT/nF ln k @ T = 298 K
E°cell = 0.0257/n ln k
Do Examples 4 & 5
IV. Dependence of Cell Potential on Concentration : IV. Dependence of Cell Potential on Concentration The Effects of Concentration on E°cell
Standard = 1.0 M
Use Le Chatelier’s Principle to predict shifts
The spontaneous reaction is the forward reaction
?fwd reaction = ?E°cell ; ?reverse = ?E°cell
Concentration Cells
Both compartments have the same components, but at different concentrations; very low E°cell
E.g. 0.1 M anode; 1.0 M cathode
c. The Nernst Equation : c. The Nernst Equation ?G = ?G° + RT ln Q, used to calculate the free energy of a cell not under standard conditions
Q is reaction quotient (initial concentrations)
Ecell = E°cell - RT/nF ln Q
@ 298K, Ecell = E°cell – 0.0591/n log Q
When Q = K, Ecell = 0; ?G = 0; dead battery
Do Examples 6, 7 & 8
#7) Fe| 0.01 M Fe2+|| Fe| 0.1 M Fe2+
V. Electrolysis : V. Electrolysis Determining the Strength of Oxidizing and Reducing Agents
Use E°cell to determine if a species is better at oxidizing or reducing (which reaction gives a more positive voltage)
Anything that can still accept electrons is a potential oxidizing agent (e.g. Cu2+, Cu+, Cl2)
Anything that can still donate electrons is a potential reducing agent (e.g. Cu, Cu+, Cl-)
Do Example 9
b. Electrolytic Cells : b. Electrolytic Cells Non-spontaneous reaction forced to happen by hooking a cell up to a power source
E.g. Zn2+ + Cu ? Zn + Cu2+ E°cell = -1.10 V
The anode is the less active metal, the cathode is the more active metal (opposite of galvanic cells!)
Often used to electroplate metals (jewelry, chrome bumpers, nails)
c. Stoichiometry of Electrolytic Processes : c. Stoichiometry of Electrolytic Processes Ampere (A) = 1 coulomb per second
Faraday (F) = 96,485 coulombs per mol e-
Use dimensional analysis to solve problems
Do Examples 10 & 11.
Slide 12 :