The species which does not undergo disproportionation reaction is :

<p>$\textbf{Oxidation States:}$</p> <p><p>In $\mathrm{ClO_4^-}$ (perchlorate), chlorine is in the +7 oxidation state.</p></p> <p><p>In $\mathrm{ClO_3^-}$ (chlorate), chlorine is in the +5 oxidation state.</p></p> <p><p>In $\mathrm{ClO_2^-}$ (chlorite), chlorine is in the +3 oxidation state.</p></p> <p><p>In $\mathrm{ClO^-}$ (hypochlorite), chlorine is in the +1 oxidation state.</p></p> <p>$\textbf{Disproportionation Reaction:}$</p> <p>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.</p> <p><p>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:</p> <p>$3\,\mathrm{ClO^-} \rightarrow 2\,\mathrm{Cl^-} + \mathrm{ClO_3^-}$</p></p> <p><p>In $\mathrm{ClO_3^-}$, chlorine is in an intermediate oxidation state (+5) that, under certain conditions, can undergo disproportionation.</p></p> <p><p>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.</p></p> <p>$$ \boxed{\mathrm{ClO_4^-} \text{ (perchlorate ion) does not undergo disproportionation.}} $$</p>