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## Summary

In physical chemistry, the Faraday constant, denoted by the symbol F and sometimes stylized as ℱ, is the electric charge of one mole of electrons. It can be thought of as the conversion factor between the mole (used in chemistry) and the coulomb (used in physics and in practical electrical measurements), and is therefore of particular use in electrochemistry. Named after Michael Faraday, it has the currently accepted value of

F = 96485.33212... C·mol−1.

Since 1 mol electrons is 6.022 × 1023 electrons (Avogadro's number), and a coulomb is the (negative) charge of 6.241×1018 electrons,  the Faraday constant is given by the quotient of these two numbers:

$F={\frac {6.022\cdot 10^{23}\;{\mathsf {e^{-}/mol}}}{6.241\cdot 10^{18}\;{\mathsf {e^{-}/C}}}}=96485.3\;{\mathsf {C/mol}}$ One common use of the Faraday constant is in electrolysis calculations. One can divide the amount of charge in coulombs (the current integrated over time in amp-hours divided by 3600) by the Faraday constant in order to find the chemical amount (in moles) of a substance that has been electrolyzed.

The value of F was first determined 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.

## 2019 redefinition

Since the 2019 redefinition of SI base units, which introduced exactly defined values for the elementary charge and the mole, the Faraday constant is exactly

e × (1 mol) mol−1 = 1.602176634×10−19 C × 6.02214076×1023 mol−1 = 96485.3321233 C·mol−1.

## Other common units

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