are looked at as redox processes in an upcoming section of this chapter.
As discussed, reactions involving ions are relatively simple to recognize as
redox (or not) because it is straightforward to identify changes in charges if they
occur. It’s not quite so simple when redox reactions involve molecular species.
For these types of reactions, as well as those involving ionic species, a related
but noteworthily different term is used to describe the responsibility particles
have for the electrons around them as a chemical reaction unfolds. To keep track
of the transfer of electrons in all formulas (ionic or molecular), chemists have
devised a system of electron bookkeeping called oxidation states (or oxidation
numbers.) In this system, an oxidation state is assigned to each member of a
formula using rules that recognize the degree to which electrons practically
belong to a particular element in a given substance from an ionic bonding
perspective. Basically, elements with high electronegativities are given
responsibility for the electrons in a bond, ionic or covalent, and the change in this
responsibility will be noted by changes in their oxidation states. In this regard, the
oxidation states system assumes an ionic perspective for all bonding, i.e.,
electrons are not shared but belong to one element or the other in a chemical
bond. Although we know that electrons are shared in a large number of chemical
bonds, particularly those described as covalent (or molecular), assigning
responsibility for the electrons in this way allows for easy recognition of how the
accountability for electrons changes during all chemical processes.
Oxidation states are designated by a small number superscript preceded by a
plus or minus sign. This is not to be confused with the ionic charges we have been
using thus far that are shown as plus or minus signs to the right of the magnitude
of ionic charge as a superscript.
The Rules for Assigning an Oxidation State
The basic rules for assigning an oxidation state to an element in a substance’s
formula are given below. By applying these simple rules, you can assign
oxidation states to elements in practically all substances you may encounter as a
beginning chemistry student. To apply these rules, remember that the sum of the
oxidation states must be zero for an electrically neutral compound. For a
polyatomic ion, the sum of the oxidation states must be equal to the charge on
the ion.
- The oxidation state of an element of an atom in an element is zero.
Examples: 0 for Na(s), O 2 (g), and Hg(l) - The oxidation state of a monatomic ion is the same as its charge.
Examples: +1 for Na+ and –1 for Cl−. - The oxidation state for fluorine is −1 in its compounds. Examples: HF,
hydrogen fluoride, has one H at +1 and one F at −1. PF 3 has one P at +3 and