Question

Out of the following redox reactions, disproportionation is not shown in:
$I$: ${NH}_4{NO}_3\underset{}{\overset{\Delta }{\to }}{N}_{2}O+2H_{2}O$
$II$: ${NH}_4{NO}_2\underset{}{\overset{\Delta }{\to }}{N}_2+2H_{2}O$
$III$: $P{Cl}_5\underset{}{\overset{\Delta }{\to }}P{Cl}_3+{Cl}_2$

• A. $I,II$
• B. $II,III$
• C. $I,III$
• D. $III$

Answer 1

The correct option is D $III$

Disproportionation is a specific type of redox reaction in which an element from a reaction undergoes both oxidation and reduction to form two different products.

1. Reaction I: nitrogen on hand oxidises from oxidation state -3 (${{NH}_4}^{+}$)to +1 ($N_{2}O$). On other hand reduces from +5 (${{NO}_3}^{-}$) to +1 ($N_{2}O$). Hence, disproportionation reaction.
2. Reaction II: nitrogen on hand oxidises from oxidation state -3 (${{NH}_4}^{+}$)to 0 ($N_2$). On other hand reduces from +3 (${{NO}_2}^{-}$) to 0 ($N_2$). Hence, disproportionation reaction.
   
3. Reaction III doesnt show disproportionation reaction as since phosphorus undergoes reduction and chlorine undergoes oxidation.

Hence, $D$ is the right answer.