Question

The following reaction is an example of disproportionation reaction.
$2{H_2}{O_2} \to 2{H_2}O + {O_2}$
If true enter 1, else enter 0.

Answer 1

Hint – Here we will proceed by explaining the concept of disproportionation reaction to check whether the given equation is disproportionate or not.

Complete step-by-step answer:
The disproportionate reaction is a reaction in which a substance is simultaneously oxidized and reduced giving two different products. It is a disproportionation reaction when a multi atomic species whose pertinent element has a specific oxidation state gets oxidised and reduced in two separate half-reactions, yielding two other products containing the same pertinent element.
As oxidation means increase in oxidation number and reduction means decrease in oxidation number. Therefore, we can say that -
It is a redox reaction in which a species is simultaneously reduced and oxidized to form two different products
$2{H_2}{O_2} \to 2{H_2}O + {O_2}$
In this reaction ${H_2}{O_2}$ is reduced into ${O_2}$ because oxidation number of oxygen atom changes from
-1 to zero.
While the oxidation number of oxygen changes from -1 to -2, so ${H_2}{O_2}$ is oxidised into ${H_2}O$. So this reaction is a disproportionation reaction.
The reaction, $2{H_2}{O_2} \to 2{H_2}O + {O_2}$ is a disproportionation reaction.
One molecule of hydrogen peroxide is oxidised to oxygen and the second molecule of hydrogen peroxide is reduced to water.
Therefore, the answer is 1.

Note – In this particular question, one must know that a disproportionate reaction is a special type of redox reaction in which an element simultaneously gives electrons and accepts electrons to form different products. This simple means that oxidation and reduction occur simultaneously with an atom, element, or ion acting both as a reducing and an oxidising agent.