A level Chemistry · Edexcel 9CH0 · Topic 14A

Electrode potentials

Key points

A half-cell contains an element in two oxidation states; standard electrode potential $E^\circ$ is measured against the standard hydrogen electrode (H₂ at 100 kPa, 1.00 mol dm⁻³ H⁺, Pt electrode) at 298 K.
Standard conditions: 1.00 mol dm⁻³ solutions, 100 kPa, 298 K. Use a platinum electrode when no metal is present, and a salt bridge (e.g. KNO₃ on filter paper) to complete the circuit.
More positive $E^\circ$ = stronger oxidising agent (more easily reduced); more negative = stronger reducing agent.
$E^\circ_{cell} = E^\circ_{right} - E^\circ_{left}$ using cell diagrams like Zn(s)|Zn²⁺(aq)||Cu²⁺(aq)|Cu(s).
Positive $E^\circ_{cell}$ means feasible, but the reaction may be slow or conditions non-standard; $E_{cell}$ is proportional to $\Delta S_{total}$ and $\ln K$ .
Changing concentration shifts the electrode equilibrium and so changes the electrode potential.

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An electrode potential measures how readily a half-cell gains electrons. Comparing two potentials tells you which way electrons will flow and whether a redox reaction is feasible.

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