|An oxidation number is the charge an atom in a compound would
have if the electron pairs in the bonds belonged entirely to the more electronegative
|In the example below, we assign the numbers of +1 to H and to
the chlorine we assign -1 in a molecule of HCl. We know that in covalently
bonded molecules such as HCl, that the atoms never carry more than partial
positive or negative charges. Nevertheless, oxidation numbers are assigned
as if each compound were ionic. It is important to remember, therefore, that
the oxidation numbers assigned to atoms in compounds do not have to
correspond to the actual charges of the atoms -- sometimes they do, but often
they do not.
|A redox reaction is a chemical reaction in which changes in
oxidation numbers occur.
|Any element, when not combined with atoms of a different element,
has an oxidation number of zero.
eg. Na = 0, Cl2 = 0
|Any simple monatomic ion (one-atom ion) has an oxidation number
equal to its charge.
Na+ = +1, S-2 = -2
|The sum of the oxidation numbers of all of the atoms in a formula
must equal the charge written for the formula.
|eg. SO4-2 = S has an oxidation number
of +6 in this ion. Each oxygen has a charge of -2. Therefore +6-2-2-2-2 =
-2 is the overall charge on the ion. The oxidation number of oxygen is always
-2. If you look at the oxidation number of S on the periodic table it can
be in many different states. Its the +6 of the S that is the most important
in this ion.
|In compounds, the oxidation number of any Group IA metal is
always +1, the oxidation number of any Group IIA elements is always +2, and
the oxidation number of aluminum is always +3. Check your periodic table
to verify this.
|In binary compounds with metals, the oxidation number of a nonmetal
is equal to the charge of its simple monatomic anion. eg. FeBr2 The
bromine is -1. therefore the Fe must be +2 and is therefore the iron(II) cation
or the ferrous cation.
|In compounds F is always -1. O is almost always -2 and H is almost always +1.|