Previously, we have examined the processes involved in writing, balancing, and interpreting reactions and looked at examples of several types of reactions. One type of reaction we did not examine closely was the oxidation-reduction reaction (sometimes called redox reaction). Even though this type of reaction is very important in the chemistry of drug molecules, it is beyond the scope of our instruction to study them in detail. However, a basic understanding of this process will be valuable to you in understanding many of the incompatibilities, storage problems, and some disease states that you will encounter later.
Review of Valence. Before these reactions are studied, valence should be reviewed briefly. The following two valence concepts are especially important in oxidation-reduction reactions:
- All elements in their free and uncombined state are considered to have a valence of zero. This holds even for those elements that are diatomic molecules in their free state.
- All atoms can exist in a number of valence states. The common valences which you learned previously are the preferred and most stable valences under normal conditions, but other valences can and do occur.
- These two concepts are important because oxidation-reduction reactions always involve a change in the valence numbers of some of the elements involved in the reaction.
Oxidation. Oxidation, in inorganic chemistry, is defined as the loss of electrons or an increase in the valence of an element. Consider, for example, the oxidation of elemental iron:
Fe0 -2e- → Fe+2
Iron in its free state has a valence of zero and is very reactive since its common valence state is +2 or +3. It loses two electrons to become the ferrous ion. The valence has gone from 0 to +2, thus iron has been oxidized. It can undergo further oxidation to the +3 valence state:
Fe+2 -1e- → Fe+3
Here the ferrous ion has lost another electron to become a ferric ion.
Reduction. In inorganic chemistry, reduction is defined as the gain of electrons or a decrease in the valence of an element. Consider the reduction of elemental oxygen:
O2 + 4e- → 2O -2
Observe that oxygen is a diatomic molecule in its free elemental form and has a valence of zero. Since the most common valence state of oxygen is -2, oxygen accepts electrons readily to become the oxygen anion. The valence of each oxygen atom has gone from 0 to -2, thus oxygen had been reduced. If the valence is made smaller (reduced), reduction has occurred.
Oxidizing and Reducing Agents. For all practical purposes, it is impossible to simply add or subtract electrons from an element except in an electrolytic cell. In fact, the oxidation of one element and the reduction of another always occur simultaneously. One element loses the electrons; the other element gains the electrons that are lost by the first. Consider these two reactions when they are combined:
2Fe - 4e → 2Fe+2
O2 + 4e- → 2O-2
2Fe + O2 → 2FeO
This is an oxidation-reduction reaction that is very common in our industrialized society. The oxidation of iron by atmospheric oxygen gives us iron oxide, commonly known as rust. In this reaction, oxygen was reduced, going from a zero to a -2 state by receiving electrons from iron. Because it accepted the electrons from iron and allowed the iron to oxidize, oxygen is called an oxidizing agent. Iron, which gave up electrons, is called the reducing agent. General characteristics of reducing and oxidizing are shown in the following table.
|REDUCING AGENT||OXIDIZING AGENT|