Oxidation numbers

In an electrochemical process, electrons are transferred between reacting species. This is in contrast to acid-base reactions, in which protons are transferred. The species that loses electrons is said to be oxidized and the species that gains electrons is reduced. Reactions of this sort, are therefore called oxidation-reduction or redox reactions.

The oxidation number of an atom is a kind of fictitious charge (not necessarily the same as the ``formal charge'' introduced in chapter 3) that can be used to describe the chemical transformations that occur in redox reactions, where electrons are transferred. The rules for assigning oxidations numbers to atoms in a molecule are as follows:

1.

The oxidation numbers of the atoms in a molecule must add up to the total charge of the molecule (0 for neutral molecules and the appropriate positive or negative charge of an ion).

2.

Alkali metal (group I) atoms have oxidation number +1 in their compounds, and alkaline earth (group II) atoms have oxidation number +2 in their compounds.

3.

Fluorine always has oxidation number -1 in its compounds.

4.

Other halogens (group VII) have oxidation number -1 in their compounds except those compounds with oxygen and other halogens, where the oxidation number can be positive.

5.

Hydrogen is assigned oxidation number +1 in its compounds. The exception is in metal hydrides, such as LiH, where rule 2 takes precedence and hydrogen is given an oxidation number -1.

6.

Oxygen is assigned an oxidation number -2 in its compounds. Exceptions are compounds with fluorine, in which rule 3 takes precedence and compounds that contain O-O bonds, where rules 2 and 5 take precedence. In superoxides, such as KO , oxygen has an oxidation number .

Examples:

1.

N O. By rule 6, oxygen has an oxidation number -2, so nitrogen must have oxidation number +1.

2.

OF . By rule 3, fluorine has an oxidation number -1, so oxygen here must have oxidation number +2.

3.

H O . By rule 5, hydrogen has an oxidation number +1. Thus oxygen must have an oxidation number -1.

Oxidation and reduction reactions are characterized by changes in the oxidation numbers of reacting species. Consider, for example, the reaction:

According to the rules of assigning oxidation numbers, oxygen must have an oxidation number -2 in MgO, hence Mg has an oxidation number of +2. In Mg(s) the oxidation number of Mg is clearly 0, and the same for oxygen in O (g). We indicate this in the reaction by

Thus, Mg loses two electrons (giving it the net positive oxidation number in MgO), hence it is oxidized, changing its oxidation state from 0 to +2. The oxygen accepts these electrons, changing its oxidation state from 0 to -2. Hence it is reduced.