# Structural Biochemistry/Oxidation states
## Background Information

There are many kinds of oxidation states of metals especially transition metals. The stability of a given oxidation state is a comparison between the ionization energy required to remove electrons from the valence orbitals and the solvation or ligation energy obtained by surrounding the metal cation with solvent or ligand atoms. Oxidation-Reduction reactions, otherwise known as redox reactions, are involved in a wide range of vital natural processes. Examples include the rusting of iron, respiration process, and the browning of certain foods. Redox reactions are also used in the operation of batteries.
Oxidation is the loss of electrons, and reduction is the gain of electrons. Each atom in a molecule is assigned an oxidation number or the oxidation state. Besides, the oxidation number for each element in a binary ionic compound equals the ionic charge. The oxidation number for each element in a covalent compound is not obvious because the atoms do not have whole charges. [The Molecular Nature of Matter and Change].
Oxidation states are assigned to metals based on a couple of different rules. These numbers that are assigned are useful in the quest of describing and determining oxidation/reduction reactions and redox equation balancing. It is important to remember that oxidation states do not suggest the compound's charge, but rather for use in balancing reactions. Some metals have fixed oxidation states while others such as iron have numerous oxidation states that are possible!

## History

The current concept of "oxidation state" was introduced by W. M. Latimer in 1938. Oxidation itself was first studied by Antoine Lavoisier, who believed that oxidation was always the result of reactions with oxygen, thus the name. Although Lavoisier's idea has been shown to be incorrect, the name he proposed is still used, albeit more generally. Oxidation states were one of the intellectual "stepping stones" Mendeleev used to derive the modern periodic table.

## Oxidation States

In order to determine whether a chemical reaction is a oxidation-reduction reaction, we must look at all the oxidation numbers (oxidation states) at all the elements involved in the reaction. To do so there are a set of rules to follow when assigning oxidation numbers.
* An atom in its pure/elemental form, the oxidation state is always zero.
* Monoatomic ions have an oxidation state equal to the charge of the ion. For example, K+ would have an oxidation number of +1, S^2- would have an oxidation state of -2.
* keep in mind that alkali metal ions always have a +1 charge which means all group 1A metals always have an oxidation number of +1. The same rule applies to Group 2A having a +2 state and Aluminum possessing a +3 oxidation number.
* Nonmetals will usually have a negative oxidation number, however there are exceptions:
* Oxygen usually has a -2 oxidation number in both ionic and molecular compounds. However, compounds with peroxide, O2^2-, gives each oxygen an oxidation state of -1.
* Hydrogen has a +1 charge when bonding to nonmetals and a -1 charge when bonding to metals.
* The oxidation state of Fluorine is -1 in all compounds.
* The sum of all oxidation numbers within a neutral compound must be zero. The sum of the oxidation numbers of a polyatomic ion must equal the number of the charge.
For example, consider the reaction: zinc metal added to a strong acid
Zn(s) + 2H+(aq) -> Zn2+(aq) + H2(g) The oxidation state of Zinc metal = 0. The oxidation state of H+ = +1. The Oxidation state of Zn2+ = +2. The oxidation state of H2 = 0. The sum is equal to zero because there are 2 H+ ions.

## References

Gary Mieesler; Donald A. Tarr; ınorganic Chemistry 3. Edition, 2004.
Theodore L. Brown; H. Eugene Lemay, Jr.; Bruce E. Bursten; Catherin J. Murphy; Chemistry: the Central Science 11. Edition, 2009.
* ↑ Oxidation State, October 28, 2012