Oxidation number

From Open Encyclopedia

The oxidation state or oxidation number is defined as the sum of negative and positive charges in an atom, which indirectly indicates the number of electrons it has accepted or donated. The oxidation number is a convenient conceptual approximation when working with complex electrochemical reactions that eases the tracking of electrons and helps verify that they have been conserved. This is especially useful when expressing complex half-reaction equations involved in oxidation/reduction reactions. Covalent compounds such as CO₂ are treated as "fictitiously ionic".

Atoms are defined as having an oxidation number 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 electrons 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. In summary, if an atom or ion donates an electron in a reaction its oxidation state is increased by one, if an element accepts an electron its oxidation state is decreased by one.

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 state of +3 is expressed as iron(III). Manganese with an oxidation state of +4 present in manganese dioxide is given the name manganese(IV) oxide. The motive for placing oxidation numbers in names is only to distinguish between different compounds of the same elements.

In an older system, species with two common oxidation numbers are named according to size of their oxidation number. The species with a larger oxidation number has the suffix "-ic" and the smaller one has "-ous". For example, Fe_{2_{O_{3_{, in which iron has an oxidation number of 3, is called ferric oxide and FeO, in which iron has an oxidation number of 2, is called ferrous oxide.}}}}

In chemical formulas, the oxidation number of ions is placed in superscript after the element's symbol. For example, oxygen(-II) is written as O^2-. Oxidation numbers of neutral atoms are not expressed. The following formula describes the element I₂ accepting two electrons to gain an oxidation number of -1.

I₂ + 2e^- → 2I^-

When dealing with oxidation-reduction or "redox" reactions, the following rules define oxidation number:

The atom with the greater electronegativity of dissimilar atoms sharing an electron is counted as receiving the electron.
Identical atoms sharing an electron are each credited with one half of the electron.

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

There are a number of rules that can be used in determining the oxidation number of a molecule or ion:

The oxidation number of (neutral) atoms and molecules of an element equal zero.
The oxidation number of a monatomic ion equals the charge of that ion.
In neutral molecules, the sum of the oxidation numbers adds up to zero.
The sum of oxidation numbers on a polyatomic ion must equal the charge of that ion.
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, including peroxide and superoxide, where one oxygen must neutralize the other's charge.
Group I ions have an oxidation number equal to +1 within compounds.
Group II ions 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, sometimes nitrogen, or other halogens, where the oxidation numbers can be positive.
Hydrogen always has an oxidation number of +1 in compounds with the more electronegative elements, namely C, N, O, F, S, Cl, Se, Br, and I. With all others, it is -1.

With the example, Cr(OH)₃, oxygen has an oxidation number of -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 * (-2 + 1) = -3. As the compound is neutral, Cr has to have a charge of +3.

Most elements have more than one possible oxidation number, with carbon having nine:

-4: CH₄

-3: C₂H₆

-2: CH₃F

-1: C₂H₂

0: CH₂F₂

+1: C₂H₂F₄

+2: CHF₃

+3: C₂F₆

+4: CF₄

For an exhaustive list of the possible oxidation states of each element, see http://www.scescape.net/~woods/periodic.html