Electrolysis and Faraday’s Law Lab

Electrolysis uses electricity to force a chemical change that would not happen on its own — splitting compounds, plating metals, producing chlorine and aluminium. And unlike most of chemistry, the amounts involved are exactly predictable.

What happens in the cell

Two electrodes sit in a liquid containing ions. Connect a supply and:

• The cathode (negative) attracts positive ions, which gain electrons and are reduced — metal plates out, or hydrogen forms.
• The anode (positive) attracts negative ions, which lose electrons and are oxidised — often giving off a gas such as oxygen or chlorine.

A mnemonic that has survived generations: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons).

Faraday’s law

The mass deposited depends only on how much charge passed:

m = M · I · t / (n · F)

• M — molar mass of the substance (g/mol)
• I — current in amperes, t — time in seconds
• n — electrons needed per ion (Cu²⁺ needs 2, Ag⁺ needs 1)
• F — the Faraday constant, 96485 C/mol

The consequences are clean and testable: double the current → double the mass. Double the time → double the mass. But switch from silver (n = 1) to copper (n = 2) and you get half as many moles for the same charge, because each ion costs twice as many electrons.

Try this

• Double the current and confirm the mass exactly doubles.
• Compare silver with copper at the same charge — silver delivers more moles, though copper is a lighter atom.
• Drop the efficiency to 70% and see the gap between the theoretical and real deposit. Side reactions steal electrons.

Why it matters

Electroplating puts chrome on taps and gold on connectors. Electrolysis produces aluminium (the Hall–Héroult process), chlorine and sodium hydroxide, and — run in reverse, in a fuel cell or battery — the same chemistry stores and releases energy.