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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 whilst expressing complex half-reaction equations involved in oxidation/reduction reactions.
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 +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 formulae, 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 -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 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 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 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 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 where instead it is -1.
With the example, Cr(OH)3, oxygen has an oxidation number of -2 (no fluorine, O-O bonds or peroxide present), and hydrogen has a state of +1 (not a metal hydride). 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.
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