Oxidation State

Oxidation state is a formal value assigned to an atom by distributing bonding electrons according to an ionic approximation. It describes the atom’s degree of oxidation rather than its actual electrical charge in a substance. Oxidation states provide a systematic way to classify compounds and follow oxidation–reduction reactions. The term oxidation number is largely synonymous in English. (goldbook.iupac.org)

Definition and electron assignment

In the IUPAC definition, oxidation state is the charge an atom would have after an ionic approximation of its bonds to atoms of different elements. For ordinary bonds, the bonding electrons are assigned entirely to the partner with higher electronegativity. Bonds between atoms of the same element are divided equally, while nonbonding electrons remain assigned to their original atom. This accounting procedure does not imply that a covalent bond physically becomes ionic. (goldbook.iupac.org)

Using a Lewis structure, the calculation can be expressed as

\mathrm{OS}=N_{\mathrm{valence}}-N_{\mathrm{assigned}},

where $N_{\mathrm{valence}}$ is the number of valence electrons in the neutral atom and $N_{\mathrm{assigned}}$ is the number assigned after the ionic approximation. An equivalent method adds the atom’s formal charge to its signed heteronuclear bond orders: positive where the atom is the less electronegative partner and negative where it is the more electronegative partner. Specialized coordination bonds can require exceptions to the simple electronegativity rule. (goldbook.iupac.org)

For example, in water, H₂O, assigning each O–H bond’s electrons to oxygen gives oxygen an oxidation state of −2 and each hydrogen +1. Water nevertheless consists of neutral molecules, not separate O²⁻ and H⁺ ions. (openstax.org)

Common assignment rules

For many compounds, oxidation states can be obtained without explicitly distributing every electron:

An uncombined element has oxidation state 0. Examples include metallic iron, H₂, O₂, and elemental sulfur.
A monatomic ion has an oxidation state equal to its charge. Na⁺ is +1; Cl⁻ is −1.
Fluorine is −1 in its compounds.
Hydrogen is usually +1, but is −1 in hydrides of sufficiently electropositive elements, such as NaH.
Oxygen is usually −2. It is −1 in peroxides such as H₂O₂; in compounds with fluorine it can be positive, as in OF₂, where it is +2.
The sum of all atomic oxidation states equals the overall charge. It is zero for a neutral compound and equals the charge of a polyatomic ion. (openstax.org)

The last rule can be written

\sum_i n_i x_i=Q,

where $n_i$ is the number of atoms assigned oxidation state $x_i$ , and $Q$ is the total charge in units of the elementary charge. In nitrate, NO₃⁻, oxygen contributes $3(-2)$ , so nitrogen must be +5:

x+3(-2)=-1.

These rules are shortcuts to electron accounting, not unconditional properties of an element. (openstax.org)

Distinction from other charge descriptions

Oxidation state must be distinguished from several related quantities:

Formal charge divides bonding electrons equally, regardless of electronegativity. In the usual Lewis structure of water, all three atoms have formal charge zero, although their oxidation states are nonzero. (goldbook.iupac.org)
Partial atomic charge describes a partition of the actual electron density. Its value depends on the partitioning method and need not be an integer. It is not interchangeable with oxidation state. (arxiv.org)
Net electric charge belongs to the complete atom, ion, or molecule. For a monatomic ion it coincides numerically with oxidation state; for a polyatomic species it equals the sum of the constituent atoms’ oxidation states. (openstax.org)

Thus, a high positive oxidation state does not mean that an atom inside a covalent compound carries an equally large localized positive charge. (arxiv.org)

Role in oxidation–reduction reactions

Oxidation increases oxidation state; reduction decreases it. Oxygen need not participate. When hydrogen reacts with chlorine to form hydrogen chloride, hydrogen changes from 0 to +1 and chlorine from 0 to −1. This is a redox reaction even though no oxygen is present. (openstax.org)

Changes in oxidation state express electron equivalents and help establish stoichiometric relationships. For example, oxidation of copper metal to Cu²⁺ represents a two-electron change:

\mathrm{Cu \rightarrow Cu^{2+}+2e^-}.

In balancing a redox equation, oxidation and reduction half-reactions are multiplied so that their electron counts cancel. Atoms and charge must also balance; aqueous reactions may require water, H⁺, or OH⁻. (openstax.org)

Oxidation state also tracks changes in organic chemistry. Applying bond-electron accounting to ordinary carbon compounds, each C–H bond contributes −1 to carbon, each C–O single bond +1, and a C=O double bond +2; C–C bonds contribute zero. For neutral carbon atoms with zero formal charge, this gives −4 in methane and +4 in carbon dioxide. These values are derived bookkeeping results, not measurements of carbon’s partial charge. (goldbook.iupac.org)

Chemical nomenclature and electron counting

In Stock nomenclature, oxidation states are indicated by Roman numerals in parentheses: FeCl₂ is iron(II) chloride, whereas FeCl₃ is iron(III) chloride. The numeral identifies the oxidation state, not the number of atoms or the charge of the entire compound. (publications.iupac.org)

In coordination chemistry, oxidation state helps organize metal electron configurations. For many transition-metal complexes, the formal $d$ -electron count is

d^{\,g-\mathrm{OS}},

where $g$ is the metal’s group number. Iron, in group 8, is formally $d^6$ as Fe(II) and $d^5$ as Fe(III). Such counts are useful classifications, but do not specify the complete electronic structure of a covalently bonded complex. (iupac.org)

Average values and ambiguous assignments

A chemical formula can determine an average oxidation state without uniquely identifying the state of every atom. For example, applying the usual hydrogen and oxygen values to ethanol, C₂H₆O, gives an average carbon oxidation state of −2. Applying bond accounting separately to its CH₃ and CH₂OH carbons instead gives −3 and −1. This derivation illustrates why atoms of the same element need not have identical oxidation states. (goldbook.iupac.org)

More difficult assignments arise when bonding is delocalized, structural information is incomplete, or the ionic approximation does not clearly distinguish the partners. IUPAC recognizes cases requiring estimates, additional measurements, or explicitly nominal oxidation states. (iupac.org)

In metal complexes containing redox-active ligands, changes in electronic structure may involve both the metal and its ligands. Such non-innocent ligands can make the metal’s oxidation state less straightforward to assign. Oxidation-state notation remains useful, but a single integer need not capture the underlying distribution of electrons. (pmc.ncbi.nlm.nih.gov)

Historical development

The word oxidation originally described reactions involving oxygen. Its meaning expanded to encompass electron transfer and changes in oxidation number. (openstax.org)

The oxidation-state concept developed before the modern electronic theory of chemical bonding. IUPAC’s 2014 technical report examined its definitions and ambiguities; the subsequent 2016 recommendations established a comprehensive definition based on ionic approximation, connecting traditional assignment rules with explicit bond-electron accounting. (doi.org)