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Oxidation numbers track electron transfer. Rules for assigning them include: elements in their standard state = 0, monatomic ions = their charge, O = −2 (except in peroxides), H = +1 (except in metal hydrides).
Practice assigning oxidation numbers to all atoms in a compound, then identify what is oxidized and reduced.
Forgetting exceptions to oxidation number rules (peroxides, metal hydrides).
From your work with oxidation numbers and balancing chemical equations, you already know the basic concept: atoms in compounds are assigned numbers that reflect how electrons are distributed. Oxidation states are a bookkeeping device — they track where electrons "belong" by assuming that all bonds are fully ionic, even when they are covalent. This artificial assignment lets you see at a glance which atoms have gained electron density and which have lost it, making it possible to identify redox reactions (reactions involving electron transfer) from the equation alone.
The rules for assigning oxidation numbers follow a clear hierarchy. Any element in its elemental form — O₂, Fe, N₂, S₈ — has an oxidation state of 0, because identical atoms share electrons equally. Monatomic ions take their charge as their oxidation state: Na⁺ is +1, Cl⁻ is −1, Ca²⁺ is +2. For compounds, fluorine is always −1 (it is the most electronegative element and always "wins" the electrons). Oxygen is −2 in most compounds, with the key exception of peroxides (like H₂O₂) where it is −1, because each oxygen shares a bond with the other. Hydrogen is +1 when bonded to nonmetals and −1 in metal hydrides (like NaH), where the metal is more electropositive and "gives" its electron to hydrogen. The sum of all oxidation numbers in a neutral compound must equal zero; in a polyatomic ion, it must equal the ion's charge.
To identify a redox reaction, assign oxidation numbers to every atom on both sides of the equation and look for changes. If an atom's oxidation number increases, it has been oxidized — it lost electron density. If it decreases, it has been reduced — it gained electron density. The mnemonic "OIL RIG" (Oxidation Is Loss, Reduction Is Gain) captures this. For example, in the reaction 2Fe + 3Cl₂ → 2FeCl₃, iron goes from 0 to +3 (oxidized) and chlorine goes from 0 to −1 (reduced). If no oxidation numbers change, the reaction is not a redox reaction — it might be an acid-base, precipitation, or other type.
This ability to identify what is oxidized and reduced is the gateway to writing half-reactions, which separate the oxidation and reduction processes and make it possible to balance complex redox equations systematically. The oxidation number rules may feel like arbitrary conventions, but they encode a real physical insight: electronegativity determines which atom in a bond controls the shared electrons, and the oxidation state reflects that assignment. Mastering these rules now pays off immediately in electrochemistry, corrosion chemistry, and metabolic biochemistry, where tracking electron flow is central to understanding how reactions work.