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In coordination chemistry, the oxidation number of a central atom in a coordination compound is the charge that it would have if all the ligands were removed along with the electron pairs that were shared with the central atom[1] (in a coordination compound, all shared electron pairs are donated by the ligands). Imagine the central atom of a molecule being stripped of all its appendages, and the intervening electrons going with the ligands, leaving the central atom 'naked.' So, for example, in water, the central atom is oxygen, and the ligands are the two hydrogens which came to the oxygen to share its electrons. In stripping, the hydrogens go away along with the two electrons they came to the oxygen for, leaving the oxygen minus 2 electrons. Thus, oxygen has an oxidation number of +2 (the absence of 2 negative charges, the electrons, is effectively equivalent to the 'presence' of 2 positive charges).

The oxidation number is used in the nomenclature of inorganic compounds. It is represented by a Roman numeral; the plus sign is omitted for positive oxidation numbers. The oxidation number is placed either as a right superscript to the element symbol, e.g. FeIII, or in parentheses after the name of the element, e.g. iron(III): in the latter case, there is no space between the element name and the oxidation number.

The oxidation number is usually numerically equal to the oxidation state and so the terms oxidation state and oxidation number are often used interchangeably. To be more precise, however, oxidation number is used in coordination chemistry with a slightly different meaning since the rules used for counting electrons are different: every electron belongs to the ligand, regardless of electronegativity. Also, oxidation numbers are conventionally represented with Roman numerals while oxidation states use Arabic numerals. The oxidation state can differ from the oxidation number in a few cases where the ligand atom is less electronegative than the central atom (e.g., in iridium phosphine complexes), resulting in a formal oxidation state that is different from the oxidation number.

Spectroscopic oxidation states

Although formal oxidation numbers can be helpful for classifying compounds, they are unmeasurable and their physical meaning can be ambiguous. Formal oxidation numbers require particular caution for molecules where the bonding is covalent, since the formal oxidation numbers require the heterolytic removal of ligands, which essentially denies covalency. Spectroscopic oxidation states, as defined by Jorgenson and reiterated by Wieghardt, are measurables that are bench-marked using spectroscopic and crystallographic data.[2]

See also


  1. ^ International Union of Pure and Applied Chemistry. "oxidation number". Compendium of Chemical Terminology Internet edition.
  2. ^ Bill, E.; Bothe, E.; Chaudhuri, P.; Chlopek, K.; Herebian, D.; Kokatam, S.; Ray, K.; Weyhermueller, T.; Neese, F.; Wieghardt, K. (2005). "Molecular and electronic structure of four- and five-coordinate cobalt complexes containing two o-phenylenediamine- or two o-aminophenol-type ligands at various oxidation levels: An experimental, density functional, and correlated ab initio study". Chemistry - A European Journal 11: 204–224.  

Study guide

Up to date as of January 14, 2010

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Oxidation numbers are the numbers of electrons that are released or gained from the outer shell electronic configuration of an atom. They refer to how an atom will bond with another. Due to this, they help in balancing oxidation-reduction reaction equations. There are several rules for how these are assigned.


  • 1. The oxidation number (o.n.) of an uncombined element is 0.
  • 2. The total o.n. of all of the atoms in a neutral compound is 0.
  • 3. The total o.n. of all of the atoms in a polyatomic ion is equal to its charge.
  • 4. In compounds, group one alkali metals have an o.n. of +1 and group two alkali earth metals have an o.n. of +2.
  • 5. In compounds, fluorine atom has an o.n. of -1, hydrogen atom has an o.n. of +1, and oxygen atom has an o.n. of -2.
  • 6. In two element compounds with metals, group 17 elements have an o.n. of -1, group 16 elements have an o.n. of -2, and group 15 elements have an o.n. of -3.
  • 7. The o.n. of hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals
  • 8. The o.n. of fluorine is -1 in all compounds. The other halogens have an o.n. of -1 in most binary compounds When combined with oxygen, as in oxyanions, however, they have a positive oxidation state.


What are the oxidation numbers of the following atoms:

1. Oxygen in O2?

  • Oxygen is an uncombined element, so by rule #1, the oxidation number is 0.

2. Carbon in CO2?

  • Because CO2 is a neutral compound, the overall oxidation number must be 0. According to rule #5, oxygen has an oxidation number of -2. Therefore, carbon must have an oxidation number of +4.

3. Oxygen in H2O?

  • Oxygen in a combined element is -2, by rule #5.

4. Sulfur in PbSO4?

  • Oxygen has an oxidation number of -2. Because we know that lead must have an oxidation number of +2, being an ion, the oxidation number of sulfur must then be +8. (-2 * 4 oxygen = +8 * 1 sulfur)

5. Oxygen in H2O2?

  • In this rare case, where the overall o. n. must still equal 0, hydrogen is still given an oxidation number of +1, forcing oxygen to have a charge of -1, in order for it to balance to 0.


1. Petrucci, Harwood, and Herring. General Chemistry: Principles and Modern Applications. Eighth edition.


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