Oxidation state

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% ionic. Oxidation states are represented by Arabic numerals and can be positive, negative, or zero. Thus, H+ would have an oxidation state of +1.

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.

Formal vs. spectroscopic oxidation states
Although formal oxidation states can be helpful for classifying compounds, they are unmeasureable and their physical meaning can be ambiguous. Formal oxidation states require particular caution for molecules where the bonding is covalent, since the formal oxidation states require the heterolytic removal of ligands, which essentially denies covalency. Spectroscopic oxidation states, as defined by Jorgenson and reiterated by Wieghart, are measureables that are bench-marked using spectroscopic and crystallographic data. Like many concepts in chemistry, spectroscopic oxidation states are powerful but require collateral measurements. Formal oxidation states, on the other hand, result from arithmetic rules, not bonding. Skill in assigning formal oxidation states is considered essential, especially in inorganic chemistry.

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.

For more about issues with calculating atomic charges, see partial charge.

From a Lewis structure

When a Lewis structure of a molecule is available, the oxidation states may be assigned unambiguously by computing the difference between the number of valence electrons that a neutral atom of that element would have and the number of electrons that "belong" to it in the Lewis structure. For purposes of computing oxidation states, electrons in a bond between atoms of different elements belong to the most electronegative atom; electrons in a bond between atoms of the same element are split equally, and electrons in lone pair belong only to the atom with the lone pair.

For example, consider acetic acid:



The methyl group carbon atom has 6 valence electrons from its bonds to the hydrogen atoms because carbon is more electronegative than hydrogen. Also, 1 electron is gained from its bond with the other carbon atom because the electron pair in the C–C bond is split equally, giving a total of 7 electrons. A neutral carbon atom would have 4 valence electrons, because carbon is in group 14 of the periodic table. The difference, 4 – 7 = –3, is the oxidation state of that carbon atom. That is, if it is assumed that all the bonds were 100% ionic (which in fact they are not), the carbon would be described as C3-.

Following the same rules, the carboxylic acid carbon atom has an oxidation state of +3 (it only gets one valence electron from the C–C bond; the oxygen atoms get all the other electrons because oxygen is more electronegative than carbon). The oxygen atoms both have an oxidation state of –2; they get 8 electrons each (4 from the lone pairs and 4 from the bonds), while a neutral oxygen atom would have 6. The hydrogen atoms all have oxidation state +1, because they surrender their electron to the more electronegative atoms to which they are bonded.

In the reaction of acetaldehyde with the Tollens' reagent to acetic acid it can be seen that in this reaction the carbonyl carbon atoms changes its oxidation state from +1 to +3 (oxidation). At the same time two equivalents of silver Ag+ are reduced to Ago.
 * [[Image:Redox Tollens Oxidationszahlen C.svg|600px|Change in oxidation state in Tollens reaction]]

Without a Lewis structure 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. This fact, combined with the fact that some elements almost always have certain oxidation states, allows one to compute the oxidation states for atoms in simple compounds.

Other rules and guidelines

 * Hydrogen has an oxidation state of +1 except when bonded to more electropositive elements such as sodium, aluminium, and boron, as in NaH, NaBH4, LiAlH4
 * Oxygen has an oxidation state of &minus;2 except where it is &minus;1 in peroxides, &minus;1/2 in superoxides, and of +2 in oxygen difluoride, OF2,+1 in O2F2.
 * Alkali metals have an oxidation state of +1 in virtually all of their compounds (exception, see alkalide).
 * Alkaline earth metals have an oxidation state of +2 in virtually all of their compounds.
 * Halogens have an oxidation state of &minus;1 except when they are bonded to oxygen or with another halogen.

Example

With the example, Cr(OH)3, oxygen has an oxidation state of &minus;2 (no fluorine, O-O bonds present), and hydrogen has a state of +1 (bonded to oxygen). So, the triple hydroxide group has a charge of 3 × (&minus;2 + 1) = &minus;3. As the compound is neutral, Cr has an oxidation state of +3.

Elements with multiple oxidation states
Most elements have more than one possible oxidation state &mdash; with carbon having nine, as follows below:


 * 1) –4: CH4
 * 2) –3: C2H6
 * 3) –2: CH3F
 * 4) –1: C2H2
 * 5)  0: CH2F2
 * 6) +1: C2H2F4
 * 7) +2: CHF3
 * 8) +3: C2F6
 * 9) +4: CF4

For an exhaustive list of the possible oxidation states of each element, see the Standard Periodic Table.

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 then 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