Faradays

Faradays are historical non-SI units of electrical charge equal to the charge carried by one mole of electrons, approximately 96,485 coulombs. The unit is named after Michael Faraday, a foundational figure in electromagnetism and electrochemistry, and is related to the Faraday constant, F, used in various electrochemical calculations. The exact value of a faraday is the product of Avogadro’s number and the elementary charge: F = NA × e ≈ 96485 C per mole of electrons.

In practice, one faraday represents the total charge corresponding to one mole of electrons. As such, 1 faraday equals about 96,485 coulombs, and calculations often use the relation between charge and moles of electrons via the Faraday constant. The term is primarily encountered in historical or specialized electrochemistry literature; it is not part of the International System of Units (SI), where coulombs are the standard unit of electric charge.

Faradays are commonly discussed in the context of Faraday’s laws of electrolysis, which relate the amount of chemical change to the total electric charge passed. The amount of substance produced or consumed in an electrochemical reaction can be expressed as moles = Q / (nF), where Q is the total charge, n is the number of electrons transferred per formula unit, and F is the Faraday constant. From this, current, time, and charge can be interrelated for practical electrolysis and deposition processes.

See also Faraday, Michael; Faraday constant; electrolysis; electrochemistry.