Faraday – IGCSE Chemistry Definition

A Faraday (symbol F) is the quantity of electric charge carried by one mole of electrons. Its value is approximately 96,500 coulombs per mole (C/mol). It is named after Michael Faraday, who established the quantitative laws of electrolysis. At IGCSE Supplement level, the Faraday constant allows candidates to calculate the mass of substance deposited or liberated during electrolysis.

The key relationships

Charge (Q) in coulombs = current (I) in amperes x time (t) in seconds.

Moles of electrons = charge (Q) / Faraday constant (96,500).

Once you know the moles of electrons, the half-equation tells you the moles of substance produced. For example, Cu²⁺ + 2e⁻ → Cu shows that 2 moles of electrons deposit 1 mole of copper. So if 0.20 mol of electrons are transferred, 0.10 mol of copper is deposited.

Worked calculation

If a current of 2.0 A flows through copper sulfate solution for 4825 seconds, the charge passed is Q = 2.0 x 4825 = 9650 C. Moles of electrons = 9650 / 96,500 = 0.10 mol. From Cu²⁺ + 2e⁻ → Cu, moles of Cu = 0.10/2 = 0.050 mol. Mass of Cu = 0.050 x 64 = 3.2 g.

Exam context

Quantitative electrolysis calculations are a Supplement topic. Paper 4 questions typically provide the current, time, and relevant half-equation, and ask for the mass of product. The Faraday constant (96,500 C/mol) is usually given in the question or on the data sheet. Candidates must know the formula Q = It and how to connect charge to moles of electrons.

Worked exam question

A current of 0.50 A is passed through molten lead(II) bromide for 3860 seconds. Calculate the mass of lead deposited at the cathode. (F = 96,500 C/mol, Ar: Pb = 207) (3 marks)

Q = I x t = 0.50 x 3860 = 1930 C [1]. Moles of electrons = 1930 / 96,500 = 0.020 mol. From Pb²⁺ + 2e⁻ → Pb, moles of Pb = 0.020 / 2 = 0.010 mol [1]. Mass = 0.010 x 207 = 2.07 g [1].

Common mistakes

Candidates frequently forget to convert time to seconds before calculating charge. Minutes must be multiplied by 60. Another error is dividing moles of electrons by the wrong number. Always check the half-equation to see how many electrons are needed per atom or molecule of product. For example, Al³⁺ + 3e⁻ → Al requires 3 moles of electrons per mole of aluminium.