QJEEM25Ch69b6: JEE Main 2025 Chemistry - Class 11 JEEM25Ch69b6 Solution: The species which does not undergo disproportionation reaction…

$\textbf{Oxidation States:}$

In $\mathrm{ClO_4^-}$ (perchlorate), chlorine is in the +7 oxidation state.

In $\mathrm{ClO_3^-}$ (chlorate), chlorine is in the +5 oxidation state.

In $\mathrm{ClO_2^-}$ (chlorite), chlorine is in the +3 oxidation state.

In $\mathrm{ClO^-}$ (hypochlorite), chlorine is in the +1 oxidation state.

$\textbf{Disproportionation Reaction:}$

A disproportionation reaction is one in which a species simultaneously undergoes oxidation and reduction. For this to occur, the element must be in an intermediate oxidation state such that it can be oxidized to a higher state and reduced to a lower one.

In $\mathrm{ClO^-}$ and $\mathrm{ClO_2^-}$ , chlorine is in lower oxidation states (+1 and +3, respectively), making them susceptible to disproportionation. For example, hypochlorite can disproportionate in basic solution as follows:

$3\,\mathrm{ClO^-} \rightarrow 2\,\mathrm{Cl^-} + \mathrm{ClO_3^-}$

In $\mathrm{ClO_3^-}$ , chlorine is in an intermediate oxidation state (+5) that, under certain conditions, can undergo disproportionation.

However, in $\mathrm{ClO_4^-}$ , chlorine is in its highest possible oxidation state (+7) and cannot be oxidized further. Since disproportionation requires one part of the species to be oxidized and the other reduced, $\mathrm{ClO_4^-}$ is thermodynamically stable and does not undergo disproportionation.

$\boxed{\mathrm{ClO_4^-} \text{ (perchlorate ion) does not undergo disproportionation.}}$