Faraday Constant

The Faraday constant (F) is the electric charge carried by one mole of electrons, equal to about 96,485 coulombs per mole. It links the amount of substance in a chemical reaction to the amount of electric charge, and it appears in the equations of electrochemistry and membrane physiology.

F is the product of the charge of one electron, about 1.602 × 10⁻¹⁹ coulombs, and Avogadro's number, 6.022 × 10²³ per mole. It is named for Michael Faraday, whose laws of electrolysis in the 1830s showed that the mass of a substance deposited at an electrode is proportional to the charge passed. In the SI system, since 2019 the value is exact because the elementary charge and Avogadro's number are defined constants.

The Faraday constant is in the relation between voltage and free energy, ΔG = −nFE, where n is the number of electrons transferred. It appears in the Nernst equation as RT/zF, which gives the voltage across a cell membrane for an ion, and it connects the gas constant to the Boltzmann constant through the elementary charge. The energy change in terms of Gibbs free energy can then be found from measured potentials.

In oxidative phosphorylation the electron transport chain moves electrons through a potential drop of about 1.1 volts from NADH to oxygen, and the energy available can be calculated with F. The same constant is used to convert the proton-motive force of chemiosmosis from volts to energy, and to relate the flow of ions in a neuron to electrical current.

• Nernst Equation — an equation that uses it
• Gibbs Free Energy — the energy it links to voltage
• Boltzmann Constant — a related fundamental constant