Electrochemical Changes
Core ideas
Electrochemistry links electron transfer to cell potential. Assign oxidation numbers; oxidation is loss of electrons at the anode, and reduction is gain of electrons at the cathode. In both galvanic and electrolytic cells the anode is where oxidation occurs; only spontaneity and the external circuit differ. A galvanic cell uses a spontaneous reaction to produce voltage; an electrolytic cell uses an external supply to drive a nonspontaneous reaction.
Standard reduction potentials, E°, are listed as reductions. Keep the cathode half-reaction as written and reverse the anode half-reaction; then E°cell = E°cathode - E°anode. Do not multiply E° when balancing electrons. A positive E°cell means spontaneous under standard conditions. The stronger oxidizing agent is the species reduced (higher E°); the stronger reducing agent is oxidized (lower E°). Cell notation: anode | anode solution || cathode solution | cathode.
Traps: changing the sign of E° when multiplying a half-reaction; calling the cathode positive in every cell (in electrolytic cells it is negative); forgetting that electrode mass changes depend on electron flow; and using standard potentials for nonstandard concentrations without noting the limitation.
Worked example: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s). E°cell = E°Cu2+/Cu - E°Zn2+/Zn = +0.34 - (-0.76) = +1.10 V, so the reaction is spontaneous and Zn is oxidized at the anode.