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This is an old revision of this page, as edited by Dwmyers (talk | contribs) at 14:06, 2 October 2002 (when oxidation state rules break down...). The present address (URL) is a permanent link to this revision, which may differ significantly from the current revision.

Just seen new post of oxidation state. Going to attempt to combine them... user:sodium

Attempt completed. The two articles are displayed below, in case I have made any errors or left out text. I did not include some of the bottom of oxidation state as it was not directly relevant, having been taken out of the electrochemistry article (actually some of the text was written by me!) user:sodium

(former oxidation state):

Oxidation state (also referred to as oxidation number) is a convenient conceptual aproximation when working with complex electrochemical reactions. It eases tracking the electrons and verifying that they have been conserved when expressing complex half reaction equations involved in oxidation-reduction reactions.

The oxidation state of an ion is the number of electrons it appears to have compared to its neutral state (which is defined as having an oxidation state of 0). If an atom or ion donates an electron in a reaction its oxidation state is increased, if an element accepts an electron its oxidation state is decreased.

The following rules define oxidation number:

1.) The atom with the greater Electronegativity of disimiliar atoms sharing an electron is counted as receiving the electron.

2.) Identical atoms sharing an electron are each credited with one/half of the electron.


For example when sodium (electronegativity of 0.9) reacts with chlorine (electronegativity of 3.0), sodium donates one electron and gains an oxidation state of +1. Chlorine accepts the electron and gains an oxidation state of -1. The sign of the oxidation state (positive/negative) actually corresponds to the value of each ion's electronic charge. The attraction of the differently charged sodium and chlorine ions is the reason they form an ionic bond with each other.

The loss of electrons of a substance is called oxidation, and the gain of electrons is reduction. This can be easily remembered through the use of the mnemonic, OIL RIG: Oxidation Is Loss Reduction Is Gain. The substance which loses electrons is also known as the reducing agent (reductant,) and the substance which accepts the electrons is called the oxidising agent (oxidant.)

A reaction in which both oxidation and reduction is occuring is called a redox reaction. These are very common; as one substance loses electrons the other substance accepts them. Redox reactions are the basis for ionic bonding.

Despite the name, an oxidation reaction does not necessarily need to involve oxygen. In fact, even fire can be fed by an oxidant other than oxygen: Fluorine fires are often unquenchable, as fluorine is an even stronger oxidant (it has a higher electronegativity) than oxygen.

(former oxidation number):

The oxidation state or oxidation number is defined as the sum of the negative and positive charges in an atom i.e., the number of electrons it has accepted or donated.

Atoms are considered to have oxidation numbers of zero, meaning that they are electrically neutral. The positive protons in the nucleus balance the negative electron cloud surrounding it, there being equal numbers of both. If an atom donates an electron it has more protons than neutrons and becomes positive. This ion is said to have an oxidation number of +1. Conversely if an atom accepts an electron it becomes negatively charged, gaining an oxidation number of -1.

Oxidation numbers are denoted in chemical names by bracketed roman numerals placed immediately after the relevant element. For example, an iron ion, with an oxidation of +3 is expressed as iron (III). Manganese with an oxidation state of +7 present in manganese oxide is given the name manganese (VII) oxide. The motive for placing oxidation numbers in names is only to distinguish between different compounds of the same elements. The actual charge (positive/negative) of the ion is not expressed because it is not necessary for this purpose.

In chemical formulas, the oxidation number of ions is placed in superscript after the element's symbol. For example, oxygen (II) is written as O2-. Oxidation numbers of neutral numbers are not expressed. The following formula describes the element I2 accepting two electrons to gain an oxidation number of -2. 

I2 + 2e- -> 2I-

Sometimes it is not immediately evident what the oxidation number of ions in a molecule are from its molecular formula. For example, given Cr(OH)3, no oxidation numbers are present yet it is clear that ionic bonding is occuring.

There are a number of rules that can be used in determining an ions atomic number:

  • The oxidation number of (neutral) atoms equal zero.
  • In neutral molecules, the sum of the oxidation numbers adds up to zero.
  • Fluorine always has a -1 oxidation number within compounds.
  • Oxygen has an oxidation number of -2 in compounds, except (i) in the presence of fluorine, in which fluorine's oxidation number takes precedence; (ii) in oxygen-oxygen bonds, where one oxygen must neutralize the other's charge; (iii) in peroxide compounds, in which it takes an oxidation number of -1.
  • Group I ions have an oxidation number equal to +1 within compounds.
  • Group II have an oxidation number of +2 within compounds.
  • Halogens, besides fluorine, generally have -1 oxidation numbers in compounds. This rule can be broken in the presence of oxygen or other halogens, where the oxidation numbers can be positive.
  • Hydrogen always has an oxidation number of +1 oxidation number in compounds, except in metal hydrides.
See also: Electrochemistry

The problem with rules this hard coded is that there will be exceptions. For example, what's the oxidation number of the oxygen in superoxide?

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