Oxidation State, A Long-Standing Issue

The oxidation state is the simplest attribute of an element in a compound. It is taught early in the chemistry curriculum as a convenient electron-counting scheme for redox reactions. Its applications range from descriptive chemistry of elements to nomenclature and electrochemistry, or as an independent variable in plots and databases of bonded-atom properties (such as radius, bond-valence parameter, standard reduction potentials, spectral parameters, or spin). The history of the oxidation state goes back about 200 years when it described the stepwise increase in the amount of oxygen bound by elements that form more than one oxide. In his 1835 textbook Unorganische Chemie, Wçhler speaks of such an “oxydationsstufe” (an older German spelling for oxidation grade). This expression remains in use for oxidation state in several languages. The equivalent term oxidation number is also common; in English this refers more to redox balancing than to the chemical systematics of an element. Under the entry for oxidation number, the IUPAC “Gold Book” gives a defining algorithm for the oxidation state of a central atom as the charge it obtains after removal of its ligands along with the shared electron pairs. The entry for oxidation state in Ref. [3] complements this with a set of charge-balance rules and of postulated oxidation states for oxygen and hydrogen with exceptions. Details vary from textbook to textbook. Some list the rules according to decreasing priority to avoid the explicit exceptions; here is an example: 1. Atoms in an element have oxidation state 0. 2. The sum of the oxidation states for atoms in a compound is 0. 3. Fluorine in compounds has the oxidation state ¢1. 4. Alkaline metals in compounds have the oxidation state + 1, alkaline-earth metals + 2. 5. Hydrogen in compounds has the oxidation state + 1. 6. Oxygen in compounds has the oxidation state ¢2.