← Oxidation-reduction processes

5 Standard potentials & cells

Watch: Standard potentials & cells Opens a YouTube search: pick a video that clicks for you Each half-reaction has a standard reduction potential E° (vs the standard hydrogen electrode, 0.00 V). A more positive E° means a stronger oxidant (more readily reduced). Cell potential E^\circ_{\text{cell}}=E^\circ_{\text{cathode}}-E^\circ_{\text{anode}} Equivalently E°(oxidant) − E°(reductant). A positive E°cell ⟹ spontaneous.

Using an E° table

• The half-reaction with the more positive E° proceeds as a reduction (cathode, +).
• The other one reverses (oxidation, anode, −).
• E°_cell > 0 ⟹ reaction is spontaneous as written; the bigger the value, the more feasible.
• E° is an intensive property. Do not multiply it when you scale a half-equation.

Worked example Zn / Cu cell E°(Cu²⁺/Cu) = +0.34 V, E°(Zn²⁺/Zn) = −0.76 V. Find E°cell and identify the anode. More positive E° = Cu²⁺/Cu ⟹ reduction (cathode). Zn is oxidised (anode). E°cell = E°(cathode) − E°(anode) = (+0.34) − (−0.76) = +1.10 V. Positive ⟹ spontaneous. Zn electrode is the anode (negative terminal). ✓ E°cell = +1.10 V; Zn is oxidised at the anode, Cu²⁺ reduced at the cathode.

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