Oxidation Number Calculator

About the Oxidation Number Calculator

Type a formula or an ion — KMnO4, Cr2O7^2-, NH4+, Fe3O4, CuSO4·5H2O, K4[Fe(CN)6] or a name such as potassium permanganate — and get the oxidation state of every element, each with the rule that set it: uncombined elements 0, fluorine −1, group 1 and 2 metals +1 and +2, hydrogen +1 (or −1 in metal hydrides), oxygen −2 (or −1 in peroxides), chlorine, bromine and iodine −1 unless bonded to oxygen, nitrogen or a lighter halogen, and algebra for the rest, because the states always add up to the charge. Polyatomic ions and brackets are worked out on their own, so ammonium nitrate gives nitrogen −3 and +5 rather than their average.

Where atoms of one element are not alike, the calculator says so instead of hiding it: Fe₃O₄ is +8/3 on average because one iron is +2 and two are +3, tetrathionate’s sulfurs average +5/2, and a formula that cannot decide between two metals lists the combinations that fit. Known exceptions such as pyrite (FeS₂, with S₂²⁻), bleaching powder (CaOCl₂, chlorine +1 and −1) and peroxo compounds are recognised, and a state an element cannot have is flagged. In a reaction it finds what is oxidised and what is reduced, the oxidising and reducing agents and, from the balanced equation, the number of electrons transferred, with each nitrogen of NH₄NO₃ and each iron of Fe₃O₄ compared in its own state rather than as an average.

With a Pro pass, copy the result or download it as a CSV, or unlock this one result. Without a pass the free preview shows the key oxidation state or what is oxidised and reduced in full, with the first rows of each table and the first step of the working; other figures hidden.

Common uses

• Chemistry homework on oxidation numbers and redox, with the rule for every element
• Checking which element is oxidised or reduced before balancing a redox equation
• Coordination chemistry: the oxidation state of the metal in a complex
• Spotting mixed-valence compounds and peroxides in an unfamiliar formula

The rules, in order

1. An uncombined element is 0 (O₂, S₈, Fe), a monatomic ion has its charge (Fe³⁺ +3).
2. The oxidation states of all the atoms add up to the charge of the molecule or ion.
3. Fluorine is −1 in its compounds; chlorine and bromine (and iodine) are −1 when not bonded to a lighter halogen, oxygen or nitrogen.
4. Group 1 metals are +1 and group 2 metals +2.
5. Hydrogen is +1, but −1 when bonded to metals or metalloids (hydrides: NaH, LiAlH₄, NaBH₄).
6. Oxygen is −2, except when bonded to oxygen (peroxides −1, superoxides −½) or to fluorine (OF₂ +2).

An earlier rule wins over a later one: in Na₂O₂ sodium stays +1 and oxygen gives way to −1. Whatever is left comes from rule 2. When two unknown elements are left, the more electronegative takes its usual negative state (HCN: N −3, so C +2), which is what the IUPAC ionic approximation does.

Averages, fractions and exceptions

A fraction is an average over atoms that are not alike, or the share of equivalent atoms: Fe₃O₄ (+8/3) holds Fe²⁺ and two Fe³⁺, S₄O₆²⁻ (+5/2) two sulfurs at +5 and two at 0, KO₂ (−½) two equivalent oxygens, Fe₀.₉₅O (+40/19) a solid that is short of iron. Carbon in organic molecules is averaged too: ethanol’s −2 is the mean of −3 and −1. The formula alone cannot place the atoms; the notes say what is known. Pyrite, FeS₂, contains disulfide S₂²⁻ (iron +2, sulfur −1), and CrO₅, H₂SO₅ and peroxodisulfates hide O–O groups: algebra with oxygen at −2 would give impossible states such as S +7, which the calculator flags.

In a reaction each atom keeps its own state: NH₄NO₃ → N₂O + 2H₂O is redox (−3 and +5 meet at +1) although nitrogen averages +1 on both sides, and red lead, Pb₃O₄ (two Pb²⁺ and one Pb⁴⁺), dissolving in nitric acid is not, though its +8/3 average would suggest it.

Oxidation and reduction

• Oxidised: the oxidation state goes up (electrons are lost); the species it was in is the reducing agent
• Reduced: the oxidation state goes down (electrons are gained); its species is the oxidising agent
• Disproportionation: one element is both oxidised and reduced (Cl₂ → Cl⁻ + ClO⁻); comproportionation: two states meet in one (IO₃⁻ + I⁻ → I₂)
• The electrons given up equal the electrons taken up in the balanced equation; to balance redox equations step by step, use the Chemical Equation Balancer

Sources

• P. Karen, P. McArdle and J. Takats, Comprehensive definition of oxidation state, IUPAC Recommendations, Pure Appl. Chem. 88, 831: the definition (the charge of an atom after the ionic approximation of its bonds) and the order of the simple rules
• OpenStax, Chemistry 2e, §4.2 Classifying Chemical Reactions: the textbook rules, oxidation and reduction, oxidising and reducing agents
• Electronegativities (Pauling) and common oxidation states: the PubChem periodic table, as in the Periodic Table

Limitations

• Oxidation states come from the formula and these rules, not from the structure: atoms of one element in different positions get their average unless the formula is a known case or splits into known ions.
• Where two elements have almost the same electronegativity (P and H, N and Cl, S and C) sources differ, and the note says so.
• In reactions, an element in several states on both sides (bleaching powder with hydrochloric acid) is matched with the least change of state the balanced equation allows; when the equation cannot be balanced, its atoms are not paired, and the page says so.
• Radicals, metal clusters and other unusual bonding can need the full IUPAC treatment with the structure.

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