To determine which of the given reactions is not an example of a disproportionation reaction, we must understand what a disproportionation reaction is. In a disproportionation reaction, a single substance is simultaneously oxidized and reduced, resulting in the formation of two different substances.
Consider the reaction \(2H_2O_2 → 2H_2O + O_2\). Here, the oxygen in hydrogen peroxide (\(H_2O_2\)) goes from an oxidation state of -1 to 0 in \(O_2\) (oxidation) and from -1 to -2 in \(H_2O\) (reduction). Hence, it is a disproportionation reaction.
For the reaction \(2NO_2 + H_2O → HNO_3 + HNO_2\), the nitrogen in \(NO_2\) is in the +4 oxidation state. In \(HNO_3\) it is in the +5 state (oxidation) and in \(HNO_2\) it is in the +3 state (reduction). Thus, it is also a disproportionation reaction.
In the reaction \(MnO_4^– + 4H^+ + 3e^– → MnO_2 + 2H_2O\), the oxidation state of manganese changes from +7 to +4. There are no two simultaneous changes in oxidation state, thus it is not a disproportionation reaction.
For the reaction \(3MnO_4^{2–} + 4H^+ → 2MnO_4^– + MnO_2 + 2H_2O\), manganese in \(MnO_4^{2-}\) has an oxidation state of +6. It goes to +7 in \(MnO_4^-\) (oxidation) and +4 in \(MnO_2\) (reduction), which constitutes a disproportionation reaction.
In summary, the reaction \(MnO_4^– + 4H^+ + 3e^– → MnO_2 + 2H_2O\) is the only one that does not involve simultaneous oxidation and reduction of the same element. Thus, it is not a disproportionation reaction.
