Faraday's Laws of Electrolysis

Chemists began passing electric current through solutions in the early 19th century; early results included the isolation of most of the alkali and alkaline earth metals by Humphry Davy. Davy's assistant, Michael Faraday, went on to study the quantitative aspects of electrolysis. By the 1830s,

Faraday had accumulated enough data to summarize his findings in two laws: (1) The mass of a substance liberated at or deposited on an electrode is proportional to the electric charge passed through the electrolyte. (2) The mass of such a substance is proportional to the molar mass of the substance, adjusted for the number of electrons required for the oxidation or reduction process. (Faraday's work was 40 years before the discovery of the electron so he worded the second law somewhat differently.)

The charge on 1 mol of electrons is the Faraday constant, 1F = 96485 C/mol. When current is passed through a solution of CuSO₄, Cu²⁺ is reduced at the cathode and water is oxidized at the anode:

If 1.000F of charge is passed through the solution, 29.5g (0.500 mol) of metallic copper is deposited on the cathode and 8.00 g (0.250 mol) of O₂(g) is liberated at the anode. In general, reduction occurs at the cathode and oxidation at the anode.

Faraday's laws can be summarized by:

• 1 faraday (F) = 96, 485 C/mol electrons;

• Mole relationships apply to electrons in

balanced oxidation-reduction half-reactions.

مواضيع ذات صلة

Concentration Cells and the Glass Electrode

Galvanic Cells and Standard Half-Cell Potentials

Half-Reaction Method

Oxidation State Method

Calculating weight, particles, and moles

Defining the mole

Having it both ways with rechargeable batteries

Electrolytic cells: Getting chemical reactions from electricity

Galvanic cells: Getting electricity from chemical reactions

Exploring Electrochemical Cells

Balancing Redox Equations

Oxidation numbers