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*Revision Materials and Past Papers* 2.1.1 Atomic structure and isotopes 2.1.2 Compounds, formulae and equations 2.1.3 Amount of substance 2.1.4 Acids 2.1.5 Redox 2.2.1 Electron structure 2.2.2 Bonding and structure 3.1.1 Periodicity 3.1.2 Group 2 3.1.3 The halogens 3.1.4 Qualitative analysis 3.2.1 Enthalpy 3.2.2 Reaction Rates 3.2.3 Chemical equilibrium 4.1 Basic concepts and hydrocarbons 4.1.2 Alkanes 4.1.3 Alkenes 4.2.1 Alcohols 4.2.2 Haloalkanes 4.2.3 Organic synthesis 4.2.4 Analytical techniques 5.1.1 How fast? 5.1.2 How far? 5.1.3 Acids, bases and buffers 5.2.1 Lattice enthalpy 5.2.2 Enthalpy and entropy 5.2.3 Redox and electrode potentials 5.3.1 Transition elements 5.3.2 Qualitative analysis 6.1.1 Aromatic compounds 6.1.2 Carbonyl compounds 6.1.3 Carboxylic acids and esters 6.2.1 Amines 6.2.2 Amino acids, amides and chirality 6.2.3 Polyesters and polyamides 6.2.4 Carbon–carbon bond formation 6.2.5 Organic synthesis 6.3.1 Chromatography and qualitative analysis 6.3.2 Spectroscopy Required Practicals

2.1.5 Redox

Oxidation number Redox reactions

Oxidation Numbers

Specification Reference 2.1.5 (a)–(c)

Quick Notes

  • Oxidation number (state): Represents the "charge" an atom would have if the bonding in the compound was fully ionic.
    • We assume electrons in bonds 'belong' to the more electronegative atoms, even if no actual electron transfer has occurred.
  • Rules for oxidation states:
    • Elements in their natural state have an oxidation number of 0.
    • Oxygen is usually −2 (except in peroxides where it is −1).
    • Hydrogen is usually +1 (except in metal hydrides where it is −1).
    • Group 1 metals are +1, Group 2 metals are +2.
    • The sum of oxidation states in a neutral compound is 0.
    • The sum of oxidation states in an ion equals the charge of the ion.
  • Roman numerals in names indicate the oxidation number of an element
    (e.g. iron(III) = +3).

Full Notes

Oxidation Numbers

Oxidation numbers help track electron transfer in reactions. It is straightforward to see how atoms have lost or gained electrons when ions get formed, however it can be harder to see how atoms have lost or gained electron density when dealing with molecules.

Example Carbon combustion

Carbon is oxidised to form carbon dioxide when combusted. However, no ions get formed, meaning it isn’t immediately clear how electrons are involved.

Edexcel A-Level Chemistry diagram showing oxidation of carbon to carbon dioxide in combustion.

To help, we consider each atom to have an ‘imaginary’ charge, described as its oxidation number (or state).

Rules for Assigning Oxidation States

Element / Case Oxidation State
Uncombined elements (e.g., O2, N2, Fe) 0
Group 1 metals +1
Group 2 metals +2
Oxygen (usual) −2
Oxygen in peroxides (O22−) −1
Oxygen in OF2 +2
Hydrogen (usual) +1
Hydrogen in metal hydrides (e.g. NaH) −1
Neutral compound Sum of oxidation states = 0
Polyatomic ion Sum of oxidation states = charge of ion

Using these rules, we can see how carbon gets oxidised from an oxidation state of 0 in C(s) to +4 in CO2(g).

Edexcel A-Level Chemistry diagram showing electron transfer and change in oxidation state of carbon from 0 to +4.

An increase in oxidation number (gets more positive) means oxidation has occurred. A decrease in oxidation number (gets more negative) means reduction has occurred.

Worked Example

Assign oxidation states in H2SO4 (sulfuric acid).

  1. H = +1 (there are 2 H, total +2).
  2. O = −2 (there are 4 O, total −8).
  3. The total charge must be 0, so S must be +6 to balance the equation:
    2(+1) + S + 4(−2) = 0 → S = +6.

Roman Numerals in Names

Oxidation numbers are shown in Roman numerals in the names of compounds, particularly for transition metals and other elements with variable oxidation states.

Examples

Summary