AQA A Level chemistry - AS Unit 2: Section 3.2.4 Redox Reactions

3.2.4 Redox Reactions - Oxidation states

Students should:

know and be able to apply the rules for assigning oxidation states in order to work out the oxidation state of an element in a compound from its formula
understand oxidation and reduction reactions of s and p block elements

Oxidation state

This is the apparent valency of an atom within a compound. It is usually considered as if the element were bonded ionically to allow the apparent number of electrons gained or lost to be assessed.

The sum of all the oxidation numbers in a species must add up to the overall charge of the species.

By convention, the oxidation number is written as a Roman numeral in the name immediately after the element to which it refers, eg. iron(II) sulfate, sulfur(VI) oxide.

The oxidation number of an uncombined element is always zero (0)

Calculating the oxidation state

There are some elements that virtually always have the same oxidation number and these can be used to calculate the oxidation numbers of other atoms to which they are bonded.

Hydrogen, for example, always has an oxidation number of -1 when bonded to a metal (more electropositive element) and +1 when bonded to a more electronegative element (non-metal).

Oxygen is always -2 (except when in the form of the peroxide, -1, or superoxide, -0.5, ions).

Group 1 and 2 metals always have oxidation numbers of 1+ and 2+ respectively.

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