A-level chemistry · Chapter 8

Redox & electrochemistry

Oxidation states, half-equations, electrode potentials

Oxidation is loss of electrons (oxidation state increases); reduction is gain. The oxidising agent is itself reduced.

Standard electrode potentials (298 K, 100 kPa, 1 mol dm⁻³)

Every potential is measured against the standard hydrogen electrode, whose E° is 0.00 V by definition. A run of values worth recognising on sight:

Zn²⁺/Zn −0.76 V · Fe²⁺/Fe −0.44 V · H⁺/H₂ 0.00 V · Cu²⁺/Cu +0.34 V · Fe³⁺/Fe²⁺ +0.77 V · Ag⁺/Ag +0.80 V · Cl₂/Cl⁻ +1.36 V · MnO₄⁻/Mn²⁺ +1.51 V

Cell EMF = E°(positive electrode) − E°(negative electrode), and the more negative half-cell is the one oxidised — it is the anode. The Daniell cell checks out: 0.34 − (−0.76) = 1.10 V.

A reaction is feasible if EMF > 0 — but feasible ≠ fast, because kinetics can still block it. And note that AQA's data booklet contains no E° table: values are always supplied in the question, so learn the method, not the numbers.

Balancing half-equations in acid

Balance the main atom first, then O with H₂O, then H with H⁺, and finally the charge with e⁻. Applied to manganate(VII) this gives MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O — purple → colourless, the classic self-indicating titration.

3.1.7Combining half-equations into full redox equations

Worked example — manganate(VII) titrating iron(II).
  1. Write both halves — oxidation: Fe²⁺ → Fe³⁺ + e⁻  ·  reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.
  2. Scale the electrons equal by multiplying the iron half-equation by 5.
  3. Add the halves and cancel the electrons: 5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O
  4. Check the charge: +17 on both sides ✓

3.1.11.1Cells — conventional representation

A cell is written with the most negative half on the left: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). A single line marks a phase boundary, the double line is the salt bridge, and EMF = E°(right) − E°(left). The salt bridge (KNO₃-soaked) completes the circuit by ion flow without letting the solutions mix.

Non-standard conditions shift E. Increase [Cu²⁺] and its half-equation shifts right, so E(cell) rises — Le Chatelier applied to electrodes.

3.1.11.2Commercial cells

Non-rechargeable cells (e.g. zinc–carbon) run an irreversible reaction, so the EMF falls as the reagents deplete. Rechargeable cells (lithium-ion) use reactions that are reversed on charging by an applied potential.

The alkaline hydrogen–oxygen fuel cell runs H₂ + 2OH⁻ → 2H₂O + 2e⁻ and O₂ + 2H₂O + 4e⁻ → 4OH⁻, giving overall 2H₂ + O₂ → 2H₂O. It holds a constant EMF while fuel is supplied and water is the only product — but hydrogen is bulky to store and mostly made from fossil fuels.

Exam tip. "Feasible" from E° values means EMF > 0 for the direction written — but always add the kinetic caveat: a feasible reaction can be immeasurably slow (high activation energy).
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