Two practical examples: the battery and the corrosion of iron

The Daniell cell: converting chemistry into electricity

A Daniell cell consists of a zinc plate immersed in a zinc sulphate solution and a copper plate immersed in a copper sulphate solution, with the two compartments connected by a salt bridge. A conductor connects the two plates via an external circuit.

The two half-reactions:

Zn = Zn2+ + 2 e-      (oxidation, at the anode)
Cu2+ + 2 e- = Cu       (reduction, at the cathode)

Check: for Zn²⁺ + 2e⁻, left-hand side 0, right-hand side (+2) + 2 × (-1) = 0. For Cu²⁺ + 2e⁻ = Cu, left-hand side: (+2) + 2 × (-1) = 0; right-hand side: 0. Both half-reactions are balanced.

Overall equation (balance):

Zn + Cu2+ -> Zn2+ + Cu

Charges: left 0+2=+2; right +2+0=+2. Balanced.

The electrons released by the zinc flow through the wire, creating an electric current: it is the energy from this redox reaction that powers a calculator or a watch.

Iron corrosion: a costly redox reaction

In humid air, iron oxidises spontaneously according to two coupled half-reactions:

Fe = Fe2+ + 2 e-
O2 + 2 H2O + 4 e- = 4 OH-

Checking the second half-reaction: oxygen, 2 (in O₂) + 2 (in 2 H₂O) = 4; on the right, 4 (in 4 OH⁻); hydrogen, 4 (in 2 H₂O) = 4 (in 4 OH⁻); charges: 0+0+4×(-1) = -4 on the left, 4×(-1) = -4 on the right.

Multiplying the first half-equation by 2 and adding the two sides gives:

2 Fe + O2 + 2 H2O -> 2 Fe2+ + 4 OH- -> 2 Fe(OH)2

This compound oxidises further in air to form rust, a mixture of iron(III) oxides and hydroxides. Unlike the protective oxide layer that forms on aluminium, rust is porous: it allows oxidation to continue deep into the metal, which explains why unprotected iron eventually becomes completely corroded.