A redox reaction can be separated spatially so that electrons travel through an external conductor while ions move through an electrolyte. This arrangement is an electrochemical cell.

At the anode, oxidation produces electrons. At the cathode, reduction consumes electrons. These definitions apply whether the cell operates spontaneously or is driven externally.

Each electrode reaction has an equilibrium electrode potential measured relative to a reference electrode. Tabulated standard reduction potentials, written $E^\circ$, refer to the defined chemical standard states. For a cell assembled from two tabulated reduction half-reactions,

$$\boxed{E_{\mathrm{cell}}^\circ=E_{\mathrm{cathode}}^\circ-E_{\mathrm{anode}}^\circ}.$$

A positive $E_{\mathrm{cell}}^\circ$ for the chosen forward reaction indicates a thermodynamic tendency to operate in that direction when all species are in their standard states.

The scale can also be related to charge: Faraday's constant $F$ is the electric charge carried by one mole of singly charged particles. A cell voltage is therefore an energy-per-charge measure of electrochemical driving force.

Changing composition, gas pressure or temperature generally changes an electrode's equilibrium potential, so $E^\circ$ is not the operating voltage under arbitrary conditions. Quantitative treatment away from standard state requires the thermodynamics of chemical mixtures and equilibrium.

Thermodynamic favorability does not determine the observed current by itself. Charge-transfer kinetics, mass transport and surface films can make a favorable reaction proceed slowly.