Faraday's Laws of electrolysis played a great part in putting *electrochemistry* on a sound quantitative scientific base. Two experimental discoveries embody these laws i) the amount of chemical change during an electrochemical reaction is proportional to the amount of electricity passed, ii) for the same amount of electricity passed, the amount of chemical change is proportional to the molecular weight of the substance divided by the number of electrons, n, required to change one molecule. This number of electrons is usually 1 or 2. The amount of electricity passed is the electrical charge, Q, defined as the integral of current passed with respect to time, t, i.e.

Faraday's Laws can be expressed in the equation

where F is Faraday's constant of proportionality. If the current in a given amount of time, t, is constant, then the amount of chemical change is

Faraday's Laws apply separately to the reactions occurring at both electrodes in an electrochemical cell, i.e. to the formation of both oxidation and reduction products, and they apply equally well to *galvanic (spontaneous) reactions* and to *electrolytic (non-spontaneous or driven) reactions*. Furthermore Faraday's Laws apply when more than one reaction takes place at an electrode. For example, in the *electroplating* of a metal, M, at a cathode a secondary (side) reaction can be the formation of hydrogen gas

The total amount of material produced (in mols) by a charge Q is

where

The Faraday's constant is equivalent to the charge associated with a unit amount of electrons and is equal to the product of the fundamental charge on a single electron, Q_{e}, and the Avogadro constant, N_{A}, i.e.,