Change in Oxidation Number

The imaginary charge that an atom would have if all of its bonds were entirely ionic is known as its oxidation number (or oxidation state). The oxidation state of an element in a compound is determined using a series of principles based on the assumption that the electron pair in a covalent bond belongs fully to the more electronegative element.

Any chemical’s oxidation number can change as a result of interacting with another highly reactive molecule. The oxidation number/state of a metal in a compound is sometimes expressed using the notation developed by Alfred Stock, a German scientist. Stock notation is the most common name for it. The oxidation number is indicated in the molecular formula by placing a Roman numeral denoting the oxidation number in parenthesis after the metal symbol.

How and when the oxidation number of a compound changes?

Whenever a reaction takes place between two or more reactants, one of the reactants gets oxidised and another gets reduced. This oxidation and Reduction cause a change in oxidation number of the compound as well as a metal. 

Consider the reaction below between elemental iron and copper sulphate:

Fe + CuSO₄ → FeSO₄ + Cu(Precipitated)

In the course of the reaction, the oxidation number of Fe increases from zero to +2. The oxidation number of copper decreases from +2 to 0. This result is very similar to the activity series of elements. Iron is above copper in the series, so will be more likely to form Fe²⁺ while converting the Cu²⁺ to metallic copper (Cu⁰).

Basic Terms involved in change of oxidation number of any compound

• Oxidation: An increment in the oxidation number of any element in the given compound or only an element is called oxidation of the element.

• Reduction: An decrement in the oxidation number of any element in the given compound or only an element is called oxidation of the element.

• Oxidising agent : A compound which can increase the oxidation number of an element in a given substance. These reagents are often termed as oxidants. 

• Reducing agent: A reagent which is capable of lowering down the oxidation number of any other element in a given substance or an idle element is called Reducing agent. These reagents are also termed as reductants. 

• Redox reactions:  Reactions which involve change in oxidation number of the interacting species.

When Oxidation number decrease and when does it increase?

An increase in oxidation number relates to the loss of negatively charged electrons, whereas a drop in oxidation number corresponds to the gain of electrons. As a result, the oxidised element or ion experiences an increase in oxidation number. The oxidation number of the atom or ion that is lowered decreases.

When we assign oxidation numbers to atoms in the species XY, we believe we are BREAKING the XY bond, and the two electrons that make up the bond are assigned to the more electronegative atom. We acquire formal oxidation values of Y(I) and X(+I) if Y is more electronegative than X. We utilize oxidation numbers to balance redox equations based on electron transfer and to keep track of electron transfer in coordination compounds (the key word here is “formal”).

Dichromate, Cr₂O₇^2-, and permanganate, MnO₄^–, are COMMON inorganic OXIDANTS i.e. these ACCEPT electrons, and the common reduction half-equations with associated colour changes are.

MnO₄^– + 8H^-1 + 5e^– → Mn⁺² + 4H₂O (l) 

Cr₂O₇^2- + 14H⁺ + 6e^–  → 2Cr⁺³ + 7H₂O(l) 

Common Examples of change in Oxidation Number

H₂S + Cl₂ → 2HCl + S

The O.N. of S increases from -2 to 0. So it is undergoing oxidation and O.N. of Cl₂ is decreased from 0 in reactant to -1 in product form.

CuO + H₂ → Cu + H₂O 

In this particular redox reaction, CuO is getting reduced to Cu. Hence, the change of CuO to Cu is a reduction reaction. H₂ is getting oxidised to H₂O. Hence, it is an oxidation reaction.

Fe_(s) + CuSO_4(aq) → FeSO_4(aq) + Cu_(s)

Oxidation state of Fe changes from 0 to +2 and the oxidation state of Cu changes from +2 to 0.

MnO₄^– + I^– → MnO₂ + I₂

Here I- is being oxidised to I₂ and MnO₄^– is being reduced to MnO₂.

Pb + PbO₂ + 2H₂SO4 → 2PbSO₄ + 2H₂O

In this reaction, Pb is getting oxidised to PbSO₄ and PbO₂ is reduced to PbSO₄.

Fe2O₃ + 3CO  → 2Fe  + 3CO₂

In this reaction, Pb is getting oxidised to PbSO₄ and PbO₂ is reduced to PbSO₄.

2Mg + O₂ → 2MgO

In this reaction, Mg is getting oxidised to MgO and O₂ is reduced to MgO.

2Fe + O₂ → 2FeO

In this reaction, Fe is getting oxidised to FeO and O₂ is reduced to MgO.

Cu + I₂  → CuI₂

In this reaction, Cu is getting oxidised to CuI₂ and I₂ is reduced to CuI₂.

H₂ + Cl₂ → HCl

In this reaction, H₂ is getting oxidised to HCl and Cl₂ is reduced to HCl.

Conclusion

Products are formed when atoms and molecules react. The nature of the change on the reactants to generate products is used to classify reactions into several categories. Reactant and product atoms/ions in a reaction may have the same or different numbers of valence electrons, depending on the reaction type.

Redox reactions are defined as those in which the number of valence electrons in the reactant atom/ion differs from the number on the product side. During the reaction, the atom/ion may have received or lost electrons. As a result, an atom/ion is either oxidised or reduced.