Redox and electrochemistry
Redox and electrochemistry links electron transfer to oxidation numbers, cell potentials and electrolysis. The exam tests balancing redox equations, predicting spontaneous cells, and quantitative electrolysis.
Core rules
- Oxidation is loss of electrons; reduction is gain. An oxidising agent is reduced, and a reducing agent is oxidised.
- Oxidation numbers: elements are 0; monatomic ions equal their charge; oxygen is usually −2 (except peroxides, −1); hydrogen is +1 with non-metals and −1 with metals; the sum equals the overall charge.
- Balance half-equations by atoms and charge, adding H2O and H+ in acidic solution or OH− in basic solution. Combine so electrons cancel.
- In an electrochemical cell, oxidation occurs at the anode and reduction at the cathode. Electrons flow through the external circuit from anode to cathode; a salt bridge or porous barrier completes the circuit.
- Standard electrode potentials E° are measured against the standard hydrogen electrode (SHE, 0.00 V). E°cell = E°cathode − E°anode. A positive E°cell indicates a spontaneous reaction; ΔG° = −nFE°cell.
- Electrolysis uses electrical energy to drive a non-spontaneous reaction. Charge Q = It, and moles of electrons = Q/F, with F = 96 500 C mol−1.
Worked example
Calculate E°cell for Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s), given E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = −0.76 V. The cathode is Cu2+/Cu (reduction); the anode is Zn2+/Zn (oxidation). E°cell = 0.34 − (−0.76) = +1.10 V. Positive, so the reaction is spontaneous under standard conditions; zinc dissolves and copper is deposited.
Traps that cost marks
- Multiplying E° values when multiplying half-equations; electrode potentials are intensive.
- Reversing the sign incorrectly when treating a reduction potential as an oxidation potential.
- Confusing anode and cathode in a cell versus an electrolytic cell; oxidation is always at the anode.
- Using standard potentials for non-standard concentrations without the Nernst equation.
- Forgetting that electrolysis may produce hydrogen or oxygen from water rather than the metal, depending on potentials.