Redox equilibria

Electrode potentials and cells

A half-cell's standard electrode potential E⦵ is measured against the standard hydrogen electrode (SHE, 0.00 V) at 298 K, 1 mol dm⁻³ and 100 kPa. Half-equations are written as reductions; the more positive E⦵, the stronger the oxidising agent (left side). $E^\circ_\text{cell}=E^\circ_\text{(positive, reduction)}-E^\circ_\text{(negative, oxidation)}$. A reaction is feasible if $E^\circ_\text{cell}>0$ (but may be slow, or conditions may be non-standard). Changing concentration shifts E (Nernst: $E=E^\circ+\frac{0.059}{z}\log\frac{[\text{ox}]}{[\text{red}]}$). $\Delta G^\circ=-nFE^\circ_\text{cell}$. Fuel cells and batteries apply this.