Standard electrode potentials E°, standard cell potentials Ecell and the Nernst equation — common mistakes

A simple way to remember is E°cell = E°(more positive) - E°(more negative). This always gives a positive E°cell for a spontaneous reaction. The half-reaction with the more positive E° is the one that proceeds as a reduction (as written), and the one with the more negative E° is forced to reverse and proceed as an oxidation.

For A-Level, you are more likely to be asked to predict the effect of changing concentration on E than to perform a full Nernst calculation. For example, increasing [Cu²⁺] in the Cu²⁺/Cu half-cell makes its potential E more positive than its standard E° value.

It is not a measured value but a value defined by international agreement. By setting a universal reference point to zero, we can create a consistent scale to measure and compare the potentials of all other half-cells.

No, you do not need to memorise them. A table of standard electrode potentials is provided in the Data Booklet for your examination.

An E°cell of zero means the system is at equilibrium under standard conditions, and ΔG° is also zero. There is no net tendency for the reaction to proceed in either the forward or reverse direction.

No. Electrode potentials and Gibbs free energy only indicate the thermodynamic feasibility or spontaneity of a reaction. They give no information about the kinetics or rate of reaction, which may be very slow due to a high activation energy.