Oxidation-reduction reactions are some of the most important chemical reactions. Redox reactions, as they are called, are the energy producing reactions in industry as well as in the body. The core of a redox reaction is the passing of one or more electrons from one species to another. The species that loses electrons is said to be oxidized, and the species gaining electrons is reduced. These are old terms, but they are still used today. Oxidation and reduction occur simultaneously.
Oxidation numbers are assigned to each element in a chemical reaction to help us learn which element is oxidized and which is reduced. If, in a reaction, the oxidation number of an element increases (becomes more positive), the element is being oxidized. On the other hand, if the oxidation number of an element decreases, the element is being reduced. The changes in oxidation numbers are also used to balance redox equations. The goal is to keep the total number of electrons lost in the oxidation equal to the total number gained in the reduction.
Ions have an electrical charge—negative if they have gained electrons and positive if they have lost electrons. The existence of ions suggests a transfer of electrons from one atom to another giving rise to the positive and negative charges. A useful extension of this concept is to assign hypothetical charges called
oxidation numbers to atoms with polar covalent bonds. The general idea is to assign the shared electrons in each bond to the more electronegative element.
As an example, use the water molecule in a standard Lewis diagram, as shown in Figure
1 .
Figure 1
Lewis structure for H2O.
Because oxygen is more electronegative than hydrogen, for the purpose of assigning oxidation numbers, it is assumed that all 4 electrons in the 2 covalent bonds are associated completely with the oxygen atom. (See Figure
2 .)
Figure 2
Assignment of oxidation numbers.
The resulting hypothetical electrical charges are the oxidation numbers, which are shown in parentheses to remind you that they are conceptual rather than real. The atoms in H2O are not ions.
Four rules apply when assigning oxidation numbers to atoms. First, the oxidation number of each atom in a pure element is defined as zero. Therefore, elemental carbon (graphite or diamond) has an oxidation number of 0, as does an atom in metallic iron, or each of the 2 hydrogen atoms in the H2 molecule:
A single-atom ion is assigned an oxidation number equal to its electrical charge. Examples are sodium or iron ions, the latter occurring in two oxidation states:
A multiple-atom molecule or ion must have oxidation numbers that sum to the electrical charge of the group of atoms. A neutral molecule has oxidation numbers adding to zero. Therefore, the oxidation numbers of the 1 nitrogen and 3 hydrogen atoms of the neutral NH3 ammonia molecule sum to 0:
whereas the oxidation numbers of the 1 carbon and 3 oxygens of the charged
carbonate ion sum to −2:
For each bond between two different elements, the shared electrons are assigned to the element of greater electronegativity, which was nitrogen in the NH3 example and oxygen in the
example. Oxygen usually has an oxidation number of −2, the halogens are commonly −1, hydrogen is almost always +1, and the alkali metals are +1.
Problem 1: What is the oxidation number of nitrogen in magnesium nitride (Mg3N2) and in nitric acid (HNO3)?
Elements
Oxidation‐Reduction Reactions
carbonate ion sum to −2:
