Redox Equilibria

Revise Redox Equilibria for Chemistry 5 (IAL) WCH15 (A2 Level) — revision notes and instant AI marking.

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Edexcel IAL Chemistry — Unit 5

Redox Equilibria

Big idea: Every half-cell "wants" to gain or lose electrons to a different degree — we measure that "want" as a standard electrode potential (E°), and comparing two E° values tells us which way electrons will actually flow, whether a reaction is feasible, and how to calculate unknown concentrations by titration.

Chapter Summary

Oxidation and reduction can be defined three ways: gain/loss of oxygen, gain/loss of hydrogen, and gain/loss of electrons (the last one is the one that matters here) — remember with OIL RIG.
Standard electrode potential (E°) is measured against the Standard Hydrogen Electrode (SHE), which is fixed at 0.00 V, under standard conditions (1.00 mol dm⁻³, 298 K, 100 kPa).
Three types of half-cell exist: metal/metal-ion, non-metal/non-metal-ion, and ion/ion (different oxidation states) — the last two need an inert platinum electrode.
The salt bridge (usually KNO₃) completes the circuit by letting ions flow, balancing charge, without letting electrons pass through it.
E°_cell = E°_right − E°_left = E°_reduction − E°_oxidation. The more positive half-cell is reduced; the less positive one is oxidised.
Conventional cell diagrams use | for phase boundaries and , for species in the same phase, with the oxidation half on the left and reduction half on the right.
A reaction is thermodynamically feasible (spontaneous) when E°_cell is positive — but this says nothing about the rate of reaction.
E° is directly proportional to total entropy change (ΔS) and to ln K (the equilibrium constant).
Fuel cells (like hydrogen-oxygen) generate electricity directly from a chemical reaction without combustion — cleaner and more efficient, but hydrogen storage brings its own problems.
Redox titrations (manganate(VII) and iodine-thiosulfate) let us calculate unknown concentrations using stoichiometric ratios from balanced half-equations.

1. Reduction & Oxidation Recap

The Three Definitions

Chemists use three interchangeable definitions of oxidation and reduction, depending on what's easiest to spot in a given reaction. They all describe the same underlying process — electron transfer — just viewed from different angles.

Process	Oxidation	Reduction
Oxygen	Addition of oxygen 2Mg + O₂ → 2MgO	Loss of oxygen 2CuO + C → 2Cu + CO₂
Hydrogen	Loss of hydrogen	Addition of hydrogen
Electrons	Loss of electrons Al → Al³⁺ + 3e⁻	Gain of electrons F₂ + 2e⁻ → 2F⁻

🧠 Memory trick — OIL RIG Oxidation Is Loss (of electrons) | Reduction Is Gain (of electrons). This is the definition that matters for everything in this chapter — electrode potentials are entirely about electron transfer.

How different blocks of the periodic table behave

Because this chapter deals with ions constantly, it helps to know instinctively what charge an element's ion will carry:

s-block elements are oxidised, losing electrons to form 1+ or 2+ ions (e.g. Na → Na⁺ + e⁻, Ca → Ca²⁺ + 2e⁻).
p-block metals are also oxidised to positive ions — usually with a charge matching their group number (Al → Al³⁺ + 3e⁻), though some (like Sn → Sn²⁺ + 2e⁻) don't follow this rule exactly.
p-block non-metals
d-block elements