An electrochemical cell is a setup used for creating an electromotive force in a conductor separating two reactions. The current is caused by the reactions releasing and accepting electrons in to the different ends of the conductor. The most common example of an electrochemical cell is a standard 1.5-volt battery.

In each half-cell is a chemical undergoing either oxidation or reduction. In a full electrochemical cell one side must be losing electrons (oxidation) in to its electrode, while the other half-cell gains electrons (reduction.) If the atoms/ions involved in the reaction are metal, the same metal is used for each electrode. If the atoms/ions involved in the reaction at each half-cell are not metal, obviously no electrode can be constructed out of it. Nonreactive metals such as platinum are used as a substitute. Finally a salt bridge is necessary to provide electrical contact between the cells--but without the solutions mixing. This can simply be a strip of filter paper soaked in saturated potassium nitrate (V) solution.

Different choices of substances for each half cell results in varying potential differences. Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions -- when equilibrium is reached the cell cannot provide further voltage. In the half-cell which is undergoing oxidation, the further the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide. Similarly, in the reduction reaction, the further the equilibrium lies to the ion/atom with the more negative oxidation state the higher the potential.

This potential can be predicted quantitatively through the use of electrode potentials (the voltage measured when the substance is connected to hydrogen.) The difference in voltage between electrode potentials gives a prediction for the potential measured.

See also: electrochemical potential