Guidelines for Assigning Oxidation Numbers:
1. Each atom in a pure element has an oxidation number of 0.
For example, the oxidation number of Cu in metallic copper is 0 and is zero for each atom in I2,
O2, or S8.
2. For ions consisting of a single atom, the oxidation number is equal to the charge on
the ion. Elements of periodic Groups 1A-3A form monatomic ions with a positive charge and
oxidation number equal to the group number. Aluminum therefore forms Al3+, and its oxidation
number is +3.
3. Fluorine is always -1 in compounds with other elements. This is a direct consequence of
the fact that fluorine is the most electronegative element.
4. Cl, Br, and I are always -1 in compounds except when combined with O of F.
This means that Cl has an oxidation number of -1 in NaCl
(in which Na is +1, as predicted by the fact that it is an element of Group IA).
In the ion ClO-, however, the Cl atom has an oxidation number of +1
(and O has an oxidation number of -2; see Guideline 5).
5. The oxidation number of H is +1.
Although this statement applies to many, many compounds, a few important exceptions occur.

When H forms a binary compound with a metal, the metal forms a positive ion and H
becomes a hydride ion, H- . Thus, in CaH2 the oxidation number of Ca is +2 (equal to the
group number) and that of H is -1.
6. O is -2 in most compounds.
Oxygen can have an oxidation number of -1 in a class of compounds called peroxides,
compounds based on the O2-2 ion. For example, in H2O2 , hydrogen peroxide, H is assigned its
usual value of +1 and so O is -1.
7. The algebraic sum of the oxidation numbers in a neutral compound must be zero; in a
polyatomic ion, the sum must be equal to the ion charge.