Electrochemistry and Redox
What you must understand
Identify oxidation and reduction by changes in oxidation number: oxidation is loss of electrons and an increase in oxidation number. Balance redox equations in acidic and basic media by the half-reaction method. Know how to build galvanic cells: anode oxidation, cathode reduction, salt bridge, and cell notation. Use standard reduction potentials to compute E°cell as cathode minus anode and to predict spontaneity. Relate E° to Gibbs energy with ΔG° = -nFE° and to the equilibrium constant. Understand electrolysis, Faraday's laws, and the Nernst equation for nonstandard conditions.
Traps to watch for
- Reversing the sign of the anode potential incorrectly; E°cell = E°cathode - E°anode uses tabulated reduction potentials as written.
- Forgetting that multiplying a half-reaction to balance electrons does not change E°.
- Confusing anode and cathode in a galvanic cell versus an electrolytic cell; oxidation is always at the anode.
- Mishandling oxygen and hydrogen in basic-media balancing, where OH- and water replace H+ and water.
- Using moles of electrons instead of charge in coulometry; remember Q = It and n = Q/F.
Worked example
For a cell with a copper cathode (E° = +0.34 V) and a zinc anode (E° = -0.76 V), E°cell is 0.34 minus (-0.76), or +1.10 V. With n = 2, ΔG° is -2 times 96485 times 1.10, about -212 kJ per mole, confirming a spontaneous reaction.
Rule of thumb
A positive E°cell means a spontaneous galvanic reaction; a negative value means electrolysis is required.