Electrochemistry and Oxidation-Reduction
Redox fundamentals
Oxidation is loss of electrons (increase in oxidation number); reduction is gain. Assign oxidation numbers by rules: elements are 0; monatomic ions equal charge; O is usually −2; H is +1 with nonmetals. Balance redox by half-reactions, balancing atoms, charge, and electrons, then combining. In acidic solution use H2O and H+; in basic solution use OH−.
Cells and potentials
A galvanic cell converts chemical energy to electrical energy. Oxidation occurs at the anode; reduction at the cathode. Cell potential E°cell = E°cathode − E°anode using reduction potentials. A positive E°cell means spontaneous. The Nernst equation adjusts for nonstandard concentrations. Electrolytic cells use external voltage to drive nonspontaneous reactions. Current × time gives charge; Faraday's constant (96,485 C/mol e−) converts charge to moles of electrons.
Worked example
For Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s), E°cell = E°Cu − E°Zn = 0.34 V − (−0.76 V) = 1.10 V. Zn is oxidized at the anode; Cu2+ is reduced at the cathode. Because E°cell is positive, the reaction is spontaneous under standard conditions.
Traps
Reversing the sign of the anode potential incorrectly; calling the cathode negative in a galvanic cell; forgetting to balance electrons before adding half-reactions; and using E° for nonstandard concentrations without the Nernst correction.