Oxidation Number Rules for Chemistry Students

Rule 1: Elemental form

An element by itself, with no charge, has an oxidation number of 0. Examples: Na, Fe, Cl₂, N₂, P₄, and S₈.

Rule 2: Monatomic ions

The oxidation number of a monatomic ion equals its charge. Examples: Cl⁻ = −1; S²⁻ = −2; Al³⁺ = +3.

Rule 3: Metals in compounds

Group IA (Group 1) metals have an oxidation number of +1. Group IIA (Group 2)metals have an oxidation number of +2. Aluminum (Group IIIA/Group 13) has an oxidation number of +3.

Rule 4: Fluorine

Fluorine has an oxidation number of −1 in compounds.

Rule 5: Hydrogen

Hydrogen has an oxidation number of +1 in compounds, unless an earlier rule takes precedence.

Rule 6: Oxygen

Oxygen has an oxidation number of −2 in compounds, unless an earlier rule takesprecedence.

Note 1: Earlier rules take precedence

If two rules conflict, apply only the rule with the smaller number. Example: In OF₂ (oxygen difluoride), fluorine is −1 by Rule 4. The two fluorine atoms contribute −2 total, so oxygen must be +2 for the neutral compound to sum to zero. Rule 4 takes precedence over Rule 6.

Note 2: Binary ionic compounds with metals

In binary ionic compounds with metals, nonmetals have oxidation numbers equal to the charges on their anions. Example: In FeCl₃, each chloride has an oxidation number of −1. Three chloride ions contribute −3, so iron has an oxidation number of +3.