Oxidation Numbers
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 atoms.
 
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.
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