Faraday constant


In physical chemistry, the Faraday constant is a physical constant defined as the quotient of the total electric charge by the amount of elementary charge carriers in any given sample of matter: it is expressed in units of coulombs per mole.
As such, it represents the "molar elementary charge", that is, the electric charge of one mole of elementary carriers. It is named after the English scientist Michael Faraday. Since the 2019 revision of the SI, the Faraday constant has an exactly defined value, the product of the elementary charge and the Avogadro constant :

Derivation

The Faraday constant can be thought of as the proportionality factor between the charge in coulombs and the amount of substance in moles, and is therefore of particular use in electrochemistry, particularly in electrolysis calculations. Because the elementary charge is exactly, and there are exactly entities per mole, the Faraday constant is given by the product of these two quantities:
The value of was first determined in the 1800s by weighing the amount of silver deposited in an electrochemical reaction, in which a measured current was passed for a measured time, and using Faraday's law of electrolysis. Until about 1970, the most reliable value of the Faraday constant was determined by a related method of electro-dissolving silver metal in perchloric acid.

Other common units

  • 96.485 kJ per volt–gram-equivalent
  • 23.061 kcal per volt–gram-equivalent
  • 26.801 A·h/mol

    Faraday – a unit of charge

Related to the Faraday constant is the "faraday", a unit of electrical charge. Its use is much less common than of the coulomb, but is sometimes used in electrochemistry. One faraday of charge is the charge of one mole of elementary charges, that is,
Where N0 is Avogadro's number, the unitless counterpart to NA. Conversely, the Faraday constant F equals 1 faraday per mole..