Download Oxidation Number Rules

Chemistry 1211
Rules for Establishing Oxidation Numbers
1. Any uncombined atom, or any atom in a molecule of an element, is assigned an
oxidation number of zero.
Examples: N2, Cl2, C, Sn, S8
2. The oxidation number of a simple, monatomic ion is the same as the charge on the
ion.
Examples: Na+ is +1, Cu+2 is +2, Cu+ is +1, F¯ is -1.
3. The oxidation numbers of some common atoms are:
a. Fluorine, the most electronegative element, is -1 in all fluorine containing
compounds.
b. In most oxygen containing compounds oxygen is -2. In peroxides (i.e. H2O2)
each oxygen has an oxidation number of -1. In the compound OF2, the
oxygen atom has an oxidation number of +2.
c. Hydrogen usually has an oxidation number of +1 except in metallic hydrides
where it then has an oxidation number of -1
Examples: HCl, hydrogen is +1; NaH, hydrogen is -1.
d. The halogens, unless bonded to an element with a higher electronegativity,
have an oxidation number of -1.
Examples: NaCl, chlorine is -1; HClO4, chlorine is +7.
4. Since chemical compounds are neutral, it stands to reason that the algebraic sum of
their oxidation numbers must equal zero.
Examples:
NaCl, sodium is +1, chlorine is -1, total is 0
CO2, carbon is +4, oxygen is -2 (so -4 for two), total is 0.
5. The sum of the oxidation numbers of atoms comprising a polyatomic ion equals the
charge on the ionic group.
Examples:
SO4-2, sulfur is +6, oxygen is -2 (so -8 for four), total is -2.
NH4+, nitrogen is -3, hydrogen is +1 (so +4 for four), total is +1.