Table 6.1 Typical half-cell reactions
Reductions Oxidations
deposition of metals
e.g. Ag+ + e– = Ag^
formation of hydrogen gas
2H+ + 2e– = H 2dissolution of metals
e.g. Cd = Cd2+ + 2e–^
formation of oxygen gas
2H 2 O = O 2 + 4H+ + 4e–^
oxidation of halogens
e.g. 2Cl– = Cl 2 + 2e–
change of oxidation state
e.g. Fe3+ + e– = Fe2+ Sn2+ = Sn4+ + 2e–Reversible Cell
One in which the half-cell reactions are reversed by reversing the current flow; such a cell is said to be
in thermodynamic equilibrium.
Standard Hydrogen Electrode (SHE)
This consists of a platinum electrode coated with platinum black to catalyse the electrode reaction and
over the surface of which hydrogen at 760 mm of mercury is passed. The electrode is in contact with a
solution of hydrogen ions at unit activity (1.228 M HCl at 20°C) and its potential is arbitrarily chosen to
be zero at all temperatures.
Electrode Potential E
The potential of an electrode measured relative to a standard, usually the SHE. It is a measure of the
driving force of the electrode reaction and is temperature and activity dependent (p. 230). By
convention, the half-cell reaction must be written as a reduction and the potential designated positive if
the reduction proceeds spontaneously with respect to the SHE, otherwise it is negative. If the sign of the
potential is reversed, it must be referred to as an oxidation potential.
Standard Electrode Potential,
Electrode potential measured in solutions where all reactants and products are at unit activity (p. 231).
Theoretical Cell Potential
The algebraic sum of the individual electrode potentials of an electrochemical cell at zero current, i.e.
Ecell = Ecathode + Eanode. In practice, when current flows in a cell or a liquid junction is present (vide infra),
and for certain electrode systems or reactions, the cell potential departs from the theoretical value.
Liquid-junction Potential
A potential developed across a boundary between electrolytes differing in