Oxidation state of sulfur

In chemistry, the oxidation state is an indicator of the degree of oxidation of an atom in a chemical compound. The formal oxidation state is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100 percent ionic. Oxidation states are represented by Arabic numerals and can be positive, negative, or zero.

The increase in oxidation state of an atom is known as an oxidation; a decrease in oxidation state is known as a reduction. Such reactions involve the formal transfer of electrons, a net gain in electrons being a reduction and a net loss of electrons being an oxidation. In this manner, the concept of oxidation state helps scientists understand oxidation-reduction (redox) chemical reactions.

History

The concept of oxidation state in its current meaning was introduced by W. M. Latimer in 1938. Oxidation itself was first studied by Antoine Lavoisier, who held the belief that oxidation was literally the results of reactions of the elements with oxygen and that the common bond in any salt was based on oxygen.

Official definition and rules

Here is the definition of the oxidation state, according to the International Union of Pure and Applied Chemistry (IUPAC):

Oxidation state: A measure of the degree of oxidation of an atom in a substance. It is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed-upon set of rules: (l) the oxidation state of a free element (uncombined element) is zero; (2) for a simple (monatomic) ion, the oxidation state is equal to the net charge on the ion; (3) hydrogen has an oxidation state of 1 and oxygen has an oxidation state of -2 when they are present in most compounds. (Exceptions to this are that hydrogen has an oxidation state of -1 in hydrides of active metals, e.g. LiH, and oxygen has an oxidation state of -1 in peroxides, e.g. H2O2; (4) the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. For example, the oxidation states of sulfur in H2S, S8 (elementary sulfur), SO2, SO3, and H2SO4 are, respectively: -2, 0, +4, +6 and +6. The higher the oxidation state of a given atom, the greater is its degree of oxidation; the lower the oxidation state, the greater is its degree of reduction.

Calculation of formal oxidation states

There are two common ways of computing the oxidation state of an atom in a compound. The first one is used for molecules when one has a Lewis structure, as is often the case for organic molecules, while the second one is used for simple compounds (molecular or not) and does not require a Lewis structure.

It should be remembered that the oxidation state of an atom does not represent the "real" charge on that atom: this is particularly true of high oxidation states, where the ionization energy required to produce a multiply positive ion are far greater than the energies available in chemical reactions. The assignment of electrons between atoms in calculating an oxidation state is purely a formalism, albeit a useful one for the understanding of many chemical reactions.