Electrochemical Equivalent Weight
Calculate the electrochemical equivalent weight, EW = M ÷ n, dividing the element's molar mass M (g/mol) by the number of electrons exchanged n (valence). The result, in g/eq, is the mass associated with transferring one mole of electrons and appears in nearly every electrochemical calculation: Faraday's law, electrochemical corrosion rate, electrodeposition and anode sizing. For divalent iron (Fe²⁺), for example, EW = 55.85 ÷ 2 ≈ 27.9 g/eq. Enter the molar mass and the number of electrons exchanged.
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Electrochemical equivalent weight
The equivalent weight (or equivalent mass) of an element in an electrochemical reaction is the mass that corresponds to the transfer of one mole of electrons. It follows from dividing the molar mass by the number of electrons exchanged in the reaction: EW = M ÷ n. For iron oxidizing to Fe²⁺ (n = 2), EW = 55.85 ÷ 2 ≈ 27.9 g/eq; for aluminium going to Al³⁺ (n = 3), EW = 26.98 ÷ 3 ≈ 9.0 g/eq. The concept is the bridge between electric charge and mass across all of electrochemistry. By Faraday's law, passing 96,485 coulombs (1 faraday, the charge of one mole of electrons) deposits or consumes exactly one equivalent weight of the substance. That is why EW shows up in every calculation of electroplating (how much metal plates out at a given current), of electrochemical corrosion rate (relating corrosion current to mass loss) and of anode sizing. Identifying n correctly (the oxidation state involved) is the step that causes most of the mistakes. Enter the molar mass and the number of electrons exchanged.
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