A unit of electric charge, symbol, F. This entry also treats the Faraday constant, symbol F. Their magnitudes are numerically identical and are determined experimentally. According to the 2006 CODATA values, F= 96 485.3399 ± 0.0024 coulombs; F = 96 485.3399 ± 0.0024 coulombs per mole.

Shortly after Volta invented the battery, it was found that when an electric current passes through water, the water decomposes into oxygen and hydrogen.

Michael Faraday investigated electrochemistry using a solution of copper sulfate. The electric current plated out the copper ions on the electrode, and through before and after weighings of the electrode Faraday discovered that the mass of copper plated out depended solely on the strength of the current and the length of time it flowed. This principle leads naturally to the idea that the quantity of matter liberated can be used as a measure of the quantity of charge.

One faraday was defined as the quantity of electricity flowing through an electrolyte that will liberate one gram-equivalent of any substance at each electrode.

When 96,494 coulombs are passed through a solution, 1 gram-equivalent weight of any ion is liberated. Hence, 96,494 coulombs, called a faraday, are associated with the charge on one equivalent of an ionic substance.¹

A definition of the same period from an American Standards document: “A faraday is the number of coulombs (96,500) required for an electrochemical reaction involving one chemical equivalent.”² A few paragraphs later in the same standard the faraday is used as a unit: “An electrochemical equivalent of an element, compound, radical or ion is the weight of that substance involved in a specified electrochemical reaction during the passage of a specified quantity of electricity, such as a faraday, ampere-hour or coulomb.” Virtually the same definitions were repeated 42 years later³, except that the number of coulombs is refined to 96485.

The definition mentions “chemical” equivalent because the faraday had slightly different values depending on whether one used the physicists' or the chemists' definition of atomic weight (96521.9 C vs 96495.7 C, respectively). The “gram-equivalent weight” became obsolete when the IUPAC and IUPAP agreed to the unified atomic mass unit (1959/1960) and the CIPM defined the mole (1967). By 1969 textbooks carry such definitions as “1 Faraday is the charge on one mole of single ionized atoms = 9.65219 × 10⁴ coulombs”.³ Actually, 1 faraday is now the charge on a mole of any single charged entity, whether particle (e⁻, H+), ion, atom, or molecule. In terms of Michael Faraday's experiments, 1 faraday of charge will deposit on the negatively-charged electrode 1 mole of silver (since silver ions are single ionized), or ½ mole of copper (since copper ions are doubly ionized), or 1/3 mole of aluminum.

Such statements as “1 mole of electrons is called a Faraday” or “A mole of electrons is given a special name: 1 Faraday” are mistaken. The faraday is a unit of electric charge, not a count of electrons.

The Faraday constant is a ratio of charge in coulombs to amount of substance, singly charged, in moles. It is thus the product of Avogadro's number by the charge on the electron.

In biochemical studies of ion transport across membranes, the Faraday constant is defined as the number of calories released as one mole of ions moves down a voltage gradient of 1 volt, about 23062 calories/mole/equivalent.

1. M. Cannon Sneed and J. Lewis Maynard.
General College Chemistry.
New York: D. Van Nostrand Company, Inc, 1944.
Page 304.

2. American Standard Definitions of Electrical Terms.
ASA C42-1941.
New York: American Institute of Electrical Engineers, (no date, but 1942).
Paragraph 60.05.330

3. Frank Jay, Editor in Chief.
IEEE Standard Dictionary of Electrical and Electronic Terms. Third Edition.
ANSI/IEEE Std 100-1984.
New York: The Institute of Electrical and Electronics Engineers in coop. with Wiley-Interscience, 1984.
Pages 342, 302.

4. Robert A. Carman.
Numbers and Units for Physics.
New York: John Wiley and Sons, 1969.

Page 207.