An Overview of Oxidation Number Rules

• The oxidation number, also known as the oxidation state, is said to be the charge possessed by an atom if all its bonds to other atoms were entirely ionic in nature.
• It is used to describe an atom’s degree of oxidation (electron loss) in a chemical molecule.
• The oxidation state can either be positive, negative, or zero.
• While totally ionic bonds do not exist in nature, many bonds display high ionicity, making oxidation numbers a helpful predictor of charge.
• An atom’s oxidation state does not represent its actual formal charge or other fundamental atomic attributes.
• This is true, especially in high oxidation states, where the ionisation energy necessary to generate a doubly positive ion is significantly larger than the energies available in chemical processes.
• Furthermore, the oxidation number of atoms in a molecule may differ depending on the electronegativity scale utilised in their computation.  
• The oxidation number is necessary for comprehending the nomenclature norms of inorganic compounds.
• In addition, numerous facts about chemical processes may be described at a fundamental level using oxidation states.

Assigning oxidation numbers

An oxidation number is defined as a positive or negative number assigned to an atom to represent its degree of oxidation or reduction. Oxidation state and oxidation number are considered synonymous. The transfer of electrons might be total or partial in a redox process. A partial electron transfer is a change in the electron density around an atom caused by a change in the other atoms to which it is covalently bound. The charge shift is determined by the atoms involved in the bond’s respective electronegativity values.

Generally, the oxidation number of an atom in a molecule is the charge that the atom would have if all polar covalent and ionic interactions culminated in a complete transfer of electrons, from the less electronegative atom to the more electronegative one. The Lewis structure of a specific chemical can be used to assign oxidation numbers, but in the case of many simple molecules, they can alternatively be allocated using the set of principles stated below.

Oxidation Number Rules

1. In a neutral free element, the oxidation number of an atom is zero. Any element in an uncombined form, whether monatomic or polyatomic, is termed a free element. For example, the oxidation number of each atom in Fe, Li, N2, Ar, and P4 is zero.
2. A monatomic (one-atom) ion’s oxidation number is the same as its charge. For example, K+, Se2+, and Au3+ have oxidation numbers of +1, +2, and +3, respectively.
3. In most compounds, the oxidation number of oxygen is -2. An exception is when oxygen is linked to fluorine, which is the only more electronegative element than oxygen.

In such compounds, oxygen possesses oxidation numbers of 0, +1, or +2. A second exception can be found in compounds containing two oxygen atoms that are linked to one another.

For example, in the case of the peroxide ion (O22-), each oxygen atom has an oxidation number of -1. Other examples of oxygen in the -1-oxidation state are sodium peroxide (Na2O2) and calcium peroxide (CaO2).

4.     In most compounds, the oxidation number of hydrogen is +1. The only exception is when hydrogen is bound to a metal in the form of a binary ionic complex known as a metal hydride. Sodium hydride (NaH) and magnesium hydride are two examples (MgH2). The hydrogen oxidation number in these compounds is -1.
5.     In all compounds, fluorine has an oxidation number of -1. Depending on the bonding environment, other halogens can have varying oxidation values.
6.     The total of the oxidation numbers of all atoms in a neutral molecule is zero. In H2O, for example, the oxidation numbers of H and O are +1 and -2, respectively. Since this formula contains two hydrogen atoms, the total of all oxidation numbers in H2O is 2(+1) + 1(-2) = 0.
7.     The sum of the oxidation numbers of all atoms in a polyatomic ion equals the ion’s overall charge. For example, in SO42-, the oxidation numbers of S and O are +6 and -2, respectively. The total of all oxidation numbers of the sulphate ion would be 1(+6) + 4(-2) = -2, which equals the ion’s charge.

 A look at the standards for assigning oxidation numbers indicates that several elements, such as nitrogen, Sulphur, and chlorine, have no precise rules. These elements, and others, can have varying oxidation values depending on the other atoms to which they are chemically bonded in a molecular molecule. Studying a few compounds can help you figure out the best way to assign oxidation values to other atoms.

Conclusion

The total number of electrons that have been removed from an element, added to an element to get it to its current state is the oxidation number of an atom. An increase in oxidation state is referred to as oxidation. Reduction is the process of lowering the oxidation state of a substance. Electrochemical reactions result in the movement of electrons. Mass and charge are preserved when balancing these reactions, but you must know which atoms are oxidised and which atoms are reduced.