![]() ![]() ![]() One common use of the Faraday constant is in electrolysis calculations. Because 1 mole contains exactly 6.022 140 76 ×10 23 entities, and 1 coulomb contains exactly C / e = 10 19 / 1.602 176 634 elementary charges, the Faraday constant is given by the quotient of these two quantities:į = N A / 1/ e = 9.648 533 212 331 001 84 ×10 4 C⋅mol −1. The Faraday constant 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. Since the 2019 redefinition of SI base units, the Faraday constant has an exactly defined value, the product of the elementary charge ( e, in coulombs) and the Avogadro constant ( N A, in reciprocal moles):į = e × N A = 1.602 176 634 ×10 −19 C × 6.022 140 76 ×10 23 mol −1 = 9.648 533 212 331 001 84 ×10 4 C⋅mol −1. It is named after the English scientist Michael Faraday. In physical chemistry, the Faraday constant (symbol F, sometimes stylized as ℱ) is a physical constant defined as the quotient of the total electric charge ( q) by the amount ( n) of elementary charge carriers in any given sample of matter: F = q/ n it is expressed in units of coulombs per mole (C/mol).Īs such, it represents the " molar elementary charge", i.e., the electric charge of one mole of elementary carriers (e.g., protons). ![]()
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