Which of the given reactions is not an example of disproportionation reaction?

Disproportionation = same element simultaneously oxidised and reduced.

Step 1 — Check each option:

(1) $\mathrm{2H_2O_2 \rightarrow 2H_2O + O_2}$: O in $\mathrm{H_2O_2}$: $-1$; in $\mathrm{H_2O}$: $-2$ (reduced); in $\mathrm{O_2}$: $0$ (oxidised) Same element (O) both oxidised and reduced → Disproportionation ✓

(2) $\mathrm{2NO_2 + H_2O \rightarrow HNO_3 + HNO_2}$: N in $\mathrm{NO_2}$: $+4$; in $\mathrm{HNO_3}$: $+5$ (oxidised); in $\mathrm{HNO_2}$: $+3$ (reduced) Same element (N) both oxidised and reduced → Disproportionation ✓

(3) $\mathrm{MnO_4^{-} + 4H^{+} + 3e^{-} \rightarrow MnO_2 + 2H_2O}$: This is a half-reaction (reduction only). Only one species of Mn is present and it only gets reduced. No simultaneous oxidation of Mn → NOT disproportionation ✗

(4) $\mathrm{3MnO_4^{2-} + 4H^{+} \rightarrow 2MnO_4^{-} + MnO_2 + 2H_2O}$: Mn in $\mathrm{MnO_4^{2-}}$: $+6$; in $\mathrm{MnO_4^{-}}$: $+7$ (oxidised); in $\mathrm{MnO_2}$: $+4$ (reduced) → Disproportionation ✓

Answer: Option (3)