Either the actual charge of an atom (ion) in a substance, assuming the atom exists as a monatomic ion, or a hypothetical charge assigned by simple rules. 2) The charge an atom would have in a substance if the pairs of electrons in each bond belonged to the more electronegative atom. 3) The number of electrons that must be added or subtracted from an atom in a combined state to convert it to the elemental form. 

Generally the following rules for assigning oxidation numbers to an atom can be used:

1. In its elementary state the oxidation number of an atom is zero. For instance the oxidation number for Cl₂ or oxygen O₂ is zero.
2. All Group IA (alkali metals) elements have an oxidation number of +1 in any compound. All Group IIA (alkali earth metals) elements have an oxidation number of +2 in any compound.
3. has an oxidation number of -1 in all of its compounds.
4. Chlorine, bromine and iodine have an oxidation number of -1 in any compound of halogen with a less electronegative element.
5. Usually oxygen has an oxidation number of -2 in a compound. Peroxides, like H₂O₂ and Na₂O₂, are the major exceptions to this and in these cases oxygen has an oxidation number of -1.
6. has an oxidation number of +1 in most of its compounds. In hydrides (compounds like NaH), however, in which hydrogen is bonded to metallic elements, hydrogen has an oxidation number of -1.
7. The sum of the oxidation numbers of the atoms in a compound always equals zero. For polyatomic ions, the oxidation numbers of the atoms add up to the charge of the ion.