Oxidation–reduction chemistry
This standard tests oxidation numbers, half-equations, electron transfer and electrochemical cells. You must balance half-equations in acidic solution, combine them, and use E° values to predict and explain spontaneity.
Core rules
- Oxidation number increases on oxidation (loss of electrons) and decreases on reduction (gain). Oxygen is usually −2, hydrogen +1, and the sum equals the charge on the ion, or zero for a neutral species.
- The oxidant is itself reduced; the reductant is itself oxidised.
- Balance half-equations by mass first (add H₂O for oxygen, H⁺ for hydrogen in acid), then by charge using electrons.
Cells
- Galvanic cell: the anode is where oxidation occurs (negative electrode) and the cathode where reduction occurs (positive). Electrons flow anode → cathode through the wire; ions migrate through the salt bridge.
- Electrolytic cell: an external supply drives the non-spontaneous reaction; the anode is still the oxidation electrode.
- E°cell = E°(cathode) − E°(anode). A positive E°cell means the reaction is spontaneous under standard conditions.
Worked example
Acidified permanganate oxidises Fe²⁺ to Fe³⁺. Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (E° = +1.51 V) Oxidation: Fe²⁺ → Fe³⁺ + e⁻ (E° = +0.77 V) Multiply the iron half-equation by 5 to equalise electrons: MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺ E°cell = 1.51 − 0.77 = +0.74 V, so the reaction is spontaneous and the purple colour fades.
Traps
- Multiplying E° by the balancing coefficient - electrode potential is intensive.
- Mishandling the sign of the oxidation potential instead of using E°(cathode) − E°(anode).
- Omitting H⁺ and H₂O when balancing oxygen in acidic solution.
- Swapping anode and cathode labels between galvanic and electrolytic cells.