Oxidation numbers that add up.

• In reactions between covalent molecules, such as H₂ + Cl₂ → 2HCl, no electron is fully transferred; the shared pair only shifts toward the more electronegative atom. To treat such reactions as redox anyway, chemists assign each atom an oxidation number.
• Oxidation number is the charge an atom would carry in a compound if every bonding pair were handed entirely to the more electronegative partner.
• Rule 1: an element in the free state has oxidation number 0, whatever its form (H₂, O₂, Cl₂, O₃, S₈, P₄, Na, Mg, Al).
• Rule 2: a monatomic ion has an oxidation number equal to its charge. In their compounds alkali metals are always +1, alkaline earth metals +2 and aluminium +3.
• Rule 3: oxygen is usually −2, but −1 in peroxides (H₂O₂, Na₂O₂), −½ in superoxides (KO₂, RbO₂), +2 in OF₂ and +1 in O₂F₂. It takes positive values only when bonded to fluorine.
• Rule 4: hydrogen is usually +1, but −1 in binary hydrides of active metals such as LiH, NaH and CaH₂.
• Rule 5: fluorine is −1 in all its compounds. Chlorine, bromine and iodine are −1 in halides but take positive values when combined with oxygen, as in oxoacids and oxoanions.
• Rule 6: the oxidation numbers of all atoms in a neutral compound add to zero; in a polyatomic ion they add to the ion's charge. In CO₃²⁻, x + 3(−2) = −2 gives carbon +4.
• For main-group elements the highest oxidation number usually equals the group number for groups 1 and 2, and the group number minus 10 for groups 13 to 17. Across period 3 it rises from +1 in NaCl to +7 in HClO₄.

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