Redox & electrochemistry

Oxidation states, half-equations, electrode potentials

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

• Measured against the standard hydrogen electrode, E° = 0.00 V by definition.
• 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). The more negative half-cell is oxidised (is the anode). Daniell cell check: 0.34 − (−0.76) = 1.10 V.
• A reaction is feasible if EMF > 0 — but feasible ≠ fast (kinetics can block it).
• 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

1. Balance the main atom, then O with H₂O, then H with H⁺, then charge with e⁻.
2. Example: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O (purple → colourless: the classic self-indicating titration).

3.1.7Combining half-equations into full redox equations

3.1.11.1Cells — conventional representation

• Written most-negative half on the left: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). Single line = phase boundary, double line = salt bridge; EMF = E°(right) − E°(left).
• Salt bridge (KNO₃-soaked): completes the circuit by ion flow without mixing the solutions.
• Non-standard conditions shift E: increase [Cu²⁺] → its half-equation shifts right → E(cell) rises (Le Chatelier applied to electrodes).

3.1.11.2Commercial cells

• Non-rechargeable (e.g. zinc–carbon): reaction irreversible, EMF falls as reagents deplete.
• Rechargeable (lithium-ion): reactions reversed on charging by an applied potential.
• Hydrogen–oxygen fuel cell (alkaline): H₂ + 2OH⁻ → 2H₂O + 2e⁻ and O₂ + 2H₂O + 4e⁻ → 4OH⁻; overall 2H₂ + O₂ → 2H₂O. Constant EMF while fuel is supplied; water is the only product — but hydrogen is bulky to store and mostly made from fossil fuels.