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(B) disproportionation reaction is a redox reaction in which the same element in a single reactant is both oxidized and reduced.$(i) \ \stackrel{0}{Cl

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$(ii)$ and $(iv)$

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(B) disproportionation reaction is a redox reaction in which the same element in a single reactant is both oxidized and reduced.$(i) \ \stackrel{0}{Cl}_{2} + 2KI \longrightarrow 2KCl + \stackrel{0}{I}_{2}$: This is a displacement reaction where $Cl$ is reduced and $I$ is oxidized. Not disproportionation.$(ii) \ \stackrel{0}{Cl}_{2} + 2OH^{-} \longrightarrow \stackrel{+1}{ClO}^{-} + \stackrel{-1}{Cl}^{-} + H_2O$: Here, $Cl$ is oxidized from $0$ to $+1$ and reduced from $0$ to $-1$. This is a disproportionation reaction.$(iii) \ \stackrel{0}{Mg} + 2\stackrel{+1}{H}Cl \longrightarrow \stackrel{+2}{Mg}Cl_{2} + \stackrel{0}{H}_{2}$: This is a displacement reaction where $Mg$ is oxidized and $H$ is reduced. Not disproportionation.$(iv) \ 2H_2\stackrel{-1}{O}_{2} \longrightarrow 2H_2\stackrel{-2}{O} + \stackrel{0}{O}_{2}$: Here, $O$ is reduced from $-1$ to $-2$ and oxidized from $-1$ to $0$. This is a disproportionation reaction.Therefore, reactions $(ii)$ and $(iv)$ are disproportionation reactions.

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