Oxidation is a type of reaction in which the reactant releases one or more electrons. It is therefore said that the reactant becomes oxidized. Oxidation is the opposite of reduction, and these two reactions always occur concurrently in a so-called reduction oxidation reaction, often called a redox reaction . The oxidation part of a redox reaction will always have one or more ‘free’ electrons among the products, as electrons are said to be delivered in this type of reaction. An example of an oxidation reaction is the oxidation of metallic sodium (Na) to sodium ions Na+.
Na → Na⁺ + e‾
As sodium releases a negatively charged electron, the ion gets a positive charge. There are also cases where more than one electron is given as for example aluminum (Al).
Na → Al3⁺ + 3e‾
What ions an element is formed depends on the nature of the element and its location in the periodic table .
In a redox reaction, electrons are transmitted from one of the reactants to another. In order to detect the electron transfer, thus matching a redox reaction, it is necessary to know the oxidation levels of atoms which is also known as oxidation steps. The oxidation number in ion compounds is the same as the charges of the incoming ions if the ions are monoatomic, that is, if the ions consist only of one element.
In molecules, the atoms bind together by covalent bonds where one or more electron pairs are divided between the two atoms in the bond. Therefore, the oxidation number of molecules should be regarded as the hypothetical charge the atoms would have. And that; if the electron pairs shared in the bonds were assigned the most electronegative atom in the bond (as in ion bonds). One can also say that the oxidation number is the charge that each atom in a compound would have if the compound was split into monoatomic ions.
Determination of oxidation number
The following rules apply for the determination of oxidation figures.
- The oxidation number of a free element such as metallic iron (Fe), oxygen atom in oxygen (O i O2) or Argon (Ar) is always equal to 0.
- The oxidation number of monoatomic ions is always equal to the charge of the ion. Ie the oxidation number is Na⁺ equal to +1, for Al3⁺ equals +3, etc.
- The sum of all oxidation numbers is for a neutral compound always equal to 0, as is the sum of the oxidation numbers of the atoms in a composite ion equal to the charge of the compound ion.
- Fluorine (F) always has the oxidation number -1, but excluding difluoro F2, where it is 0.
- In almost all cases, the hydrogen (H) oxidation number +1 and oxygen (O) has the oxidation number -2.
In addition to these rules, the periodic table can be used to determine oxidation rates for some metals. For example, metals in the 1st main group always have an oxidation number of +1 and metals in the 2nd main group of +2 when they form part of a compound.
We now look at a few examples where the oxidation numbers of the atoms in different compounds are determined.
According to the rules, H has an oxidation number of +1 and oxygen of -2. The compound is neutrally charged and therefore the sum of the oxidation numbers must give 0. With 2 hydrogen atoms and 4 oxygen atoms, 2 · (+1) + 4 · (-2) = -6 are obtained. The sulfur atom (S) must therefore have an oxidation number of 6 to give the sum 0.
For this compound ion, the sum must yield -1, and since oxygen ‘always’ has an oxidation number of -2, the total oxidation number of the oxygen atoms is 6. Therefore, the chlorometry (Cl) must necessarily have an oxidation number of +5.
Potassium is in the 1st main group and therefore has +1 and oxygen again -2. By multiplying the number of the two elements, you get -12. Therefore, since the compound contains two chromatomers (Cr), they both have an oxidation number of 6, so that the sum gives 0.
A case where oxygen does not have an oxidation number of -2 is hydrogen peroxideH2O2. Instead, oxygen here has an oxidation number of -1. Similarly, hydrogen has an oxidation number of -1 in metal hydrides as NaH.