Which of the following is not an example of disproportionation reaction? (a) \(4 \mathrm{ClO}_{3}^{-} \rightarrow \mathrm{Cl}^{-}+3 \mathrm{ClO}_{4}^{-}\) (b) \(2 \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\) (c) \(2 \mathrm{NO}_{2}+2 \mathrm{OH}^{-} \rightarrow \mathrm{NO}_{2}^{-}+\mathrm{NO}_{3}^{-}+\mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{TiCl}_{4}+2 \mathrm{Mg} \rightarrow \mathrm{Ti}+2 \mathrm{MgCl}_{2}\)

Redox reactions, or reduction-oxidation reactions, are chemical processes where the oxidation state of one or more atoms changes due to the transfer of electrons. These reactions are essential for many biochemical and industrial processes. In these reactions, there are two half-reactions: reduction, where an element gains electrons and its oxidation state decreases, and oxidation, where an element loses electrons and its oxidation state increases.

For instance, in the reaction where hydrogen peroxide decomposes into water and oxygen gas \(2 \text{H}_2 \text{O}_2 \rightarrow 2 \text{H}_2\text{O} + \text{O}_2\), the oxygen in hydrogen peroxide, with an oxidation state of -1, is both reduced to -2 in water and oxidized to 0 in oxygen gas. Such chemical reactions where both reduction and oxidation occur simultaneously are the backbone of many natural and technological processes.

The oxidation state, also referred to as oxidation number, is a number assigned to an element in a chemical compound that represents the number of electrons lost or gained by an atom of that element in the compound. It helps in determining how electrons are distributed in a molecule and which atoms are undergoing oxidation or reduction in redox reactions.

An element can have different oxidation states in different compounds or even in the same molecule. For example, chlorine in \(\text{ClO}_3^-\) has an oxidation state of +5, and in \(\text{Cl}^-\) it has an oxidation state of -1, as seen in the disproportionation reaction \(4 \text{ClO}_3^- \rightarrow \text{Cl}^-+3\text{ClO}_4^-\). Understanding oxidation states is crucial for analyzing redox processes, such as disproportionation reactions.

Chemical equations are symbolic representations of chemical reactions, where reactants are transformed into products. These equations must be balanced, meaning that the number of atoms for each element must be the same on both sides of the equation. This reflects the law of conservation of mass.

For a disproportionation reaction to be properly represented, the chemical equation must illustrate the same element undergoing both oxidation and reduction, like in the reaction \(2 \text{NO}_2 + 2 \text{OH}^- \rightarrow\text{NO}_2^- + \text{NO}_3^- + \text{H}_2 \text{O}\) where nitrogen's oxidation states change from +4 to both +3 and +5. It's essential to understand how to read and write chemical equations to grasp the concept of disproportionation and other redox reactions.