We can keep track of how many electrons an atom has using its oxidation number, a convention used to describe the charge of an atom. The rules for assigning oxidation numbers (also known as oxidation states) are described below.
In rule 1 we see that atoms in their elemental form always have oxidation
numbers (“ON”) of zero. For example, each atom in
Rule 3 states that the sum of all of the oxidation numbers in a compound equals the overall charge of a molecule or polyatomic ion. An example is the hydroxide ion (OH-). The oxidation number of hydrogen in the compound is +1 and the oxidation number of oxygen is -2 (as we would expect from rule 4). The sum of these two oxidation numbers equals the overall charge of the hydroxide ion, -1.
Rule 4 describes a set of “common” oxidation numbers that can be used to determine the oxidation numbers of atoms in compounds. These rules are presented roughly in order of priority, starting with the group I metals (e.g., Li, Na, etc.). The group I metals, when found in a compound, always have oxidation numbers of +1. Similarly, group II metals (e.g., Mg, Ca, etc.) in compounds always have oxidation numbers of +2. Hydrogen usually has an oxidation number of +1. The only exceptions to this are hydride compounds, which are compounds that consist only of a metal and hydrogen; in the context of a hydride hydrogen has an oxidation number of -1. Fluorine always has an oxidation number of -1 when found in a compound. Oxygen almost always has an oxidation number of -2. The only exceptions are peroxides and superoxides (which contain oxygen-oxygen bonds) and compounds that contain oxygen-fluorine bonds. The other halogen elements (chlorine, bromine, and iodine) usually have oxidation numbers of -1 unless they are found to atoms that are more electronegative than they are (e.g., O or F).
Table 1: Common Oxidation Numbers for Elements
| Element/Group | Common Oxidation Number* |
|---|---|
| Group 1A | +1 |
| Group 2A | +2 |
| F | -1 |
| H | +1 |
| O | -2 |
| Group 7A | -1 |
| Group 6A | -2 |
| Group 5A | -3 |
*These are simply a list of common oxidation numbers. We will see examples that do not follow this set of common numbers.
The general idea behind oxidation numbers is that they are determined by assigning electrons in covalent bonds to whichever atom in the bond is more electronegative. Once all of the covalent bonds are broken in this way, the oxidation number of each atom is equal to the charge of the resulting monatomic ions. The example below walks through the how to do this with water.
Example: Let’s consider the O-H bonds in water. In these bonds O is more electronegative than H, so let’s break these two bonds and in both cases we’ll give the electrons in the bonds to oxygen. In the second row we see that this results in two hydrogen ions with +1 charges. The charges on these ions tell us the oxidation number of hydrogen in the compound (+1). Breaking bonds in this fashion also resulted in an oxygen with a -2 charge; this tells us that the oxidation number of oxygen in water is -2.
This method can be applied to more complicated molecules as well.