REDOX reactions can take 2 forms; reduction & oxidation.
Oxidation is the loss of electrons (or the gain of hydrogen) whereas Reduction is the gain of electrons (or
the loss of a hydrogen). You can remember this through the term OILRIG.
Reducing Agents (denoted by [H]) reduce other species, so are in themselves oxidised when they do so
Oxidising Agents (denoted by [O]) oxidise other species, so are in themselves reduced when they do so
This means that REDUCTION & OXIDATION must always occur simultaneously, as if one species is reduced,
another one must be oxidised
Oxidation States:
Oxidation states refer to the charge that would be present on a species if it was ionic. There are some
given oxidation states that are always true;
 Group 1 metals are always +1
 Group 2 metals are always +2
 Hydrogen is always +1 except in metal hydrides where it is -1 (NaH)
 Oxygen is always -2 except in peroxides where it is -1 (HOOH)
 Halogens are usually (But not always) -1
In addition to these given oxidation states, there are a few simple rules to follow:
 In elements or elemental compounds, the oxidation state is always 0
 The total of the oxidation states must equal the overall charge on the molecule
 Simple ions have the oxidation state of the charge on the given ion
 Most electronegative ion can be assumed the negative oxidation state
To find the oxidation state, you should set up an equation where you add up all the oxidation states (taking
into account subscripts) and set equal to the overall charge on the molecule. Then solve this equation to
find the required oxidation state of the given species.
Generally, if the OS has increases (more positive), electrons have been lost, so oxidation has occurred. If
the OS has decreased (more negative), electrons have been gained, so reduction has occurred.
Disproportionation Reactions:
These reactions are quite rare. They occur when a single species is both oxidised and reduced at the same
time. An example is chlorine in the reaction of chlorine with water;

In this reaction, chlorine has an OS of 0 in Cl2, but an OS of +1 in HOCl and an OS of -1 in HCl, therefore has
been both oxidised and reduced, so is a disproportionation reaction.
REDOX half equations:
 Write out the initial and final species (final species is what is being formed – asked for in the
equation, therefore your initial species are what are used to form the final species)
 Find out the oxidation states of each atom in the species
 Balance for Oxygen if present by adding water
 Balance for Hydrogen if present by adding H+
 Balance for charge by adding electrons
 Add state symbols and check charge is conserved
Note, you multiply charges by any coefficients of the species in the equation.