Oxidation numbers

Oxidation numbers are bookkeeping numbers. They allow chemists to do things such as balance redox equations.

Oxidation numbers are positive or negative numbers, but don’t confuse them with charges on ions or valences. Chemists assign oxidation numbers to elements using these rules:

✓ For free elements: The oxidation number of an element in its free (uncombined) state is zero (for example, Al(s) or Zn(s)). This is also true for elements found in nature as diatomic (two-atom) elements (H₂, O₂, N₂, F₂, Cl₂, Br₂, or I₂) and for sulfur, found as S₈.

✓ For single-atom ions: The oxidation number of a monatomic

(one-atom) ion is the same as the charge on the ion (for example, Na⁺ = +1, S^2– = –2).

✓ For compounds: The sum of all oxidation numbers in a neutral compound is zero. The sum of all oxidation numbers in a polyatomic (many-atom) ion is equal to the charge on the ion. This rule often allows chemists to calculate the oxidation number of an atom that may have multiple oxidation states, if the other atoms in the ion have known oxidation numbers.

✓ For alkali metals and alkaline earth metals in compounds:

The oxidation number of an alkali metal (IA family) in a compound is +1; the oxidation number of an alkaline earth metal (IIA family) in a compound is +2.

✓ For oxygen in compounds: The oxidation number of oxygen in a compound is usually –2. If, however, the oxygen is in a class of compounds called peroxides (for example, hydrogen peroxide, or H2O2), then the oxygen has an oxidation number of –1. If the oxygen is bonded to fluorine, the number is +1.

✓ For hydrogen in compounds: The oxidation state of

hydrogen in a compound is usually +1. If the hydrogen is part of a binary metal hydride (compound of hydrogen and some metal), then the oxidation state of hydrogen is –1.

✓ For halogens: The oxidation number of fluorine is always –1. Chlorine, bromine, and iodine usually have an oxidation number of –1, unless they’re in combination with an oxygen or fluorine. (For example, in ClO^–, the oxidation number of oxygen is –2 and the oxidation number of chlorine is +1; remember that the sum of all the oxidation numbers in ClO^– has to equal –1.)

These rules give you another way to define oxidation and reduction — in terms of oxidation numbers. For example, consider this reaction, which shows oxidation by the loss of electrons:

Zn(s) → Zn²⁺ + 2e^–

Notice that the zinc metal (the reactant) has an oxidation number of zero (the first rule), and the zinc cation (the product) has an oxidation number of +2 (the second rule). In general, you can say that a substance is oxidized when there’s an increase in its oxidation number.

Reduction works the same way. Consider this reaction:

Cu²⁺ + 2e^– → Cu(s)

The copper is going from an oxidation number of +2 to zero. A substance is reduced if there’s a decrease in its oxidation number.