Question

Which of the following is an example of disproportionation redox reaction?

• A. $N_2(g)+ O _2(g) \rightarrow 2 NO (g)$
• B. $2 H _2(g)+ O _2(g) \rightarrow 2 H _2 O (l)$
• C. $2 Pb \left( NO _3\right)_2(s) \rightarrow 2 PbO (s)+4 NO _2(g)+ O _2(g)$
• D. $NaH (s)+ H _2 O (l) \rightarrow NaOH (a q)+ H _2(g)$
• E. $2 NO _2(g)+2 OH ^{-} \rightarrow NO2 ^{-}( aq )+ NO _3^{-}( aq )+ H _2 O ( l )$

Answer 1

Answer: (E)

Disproportionation redox reaction involves simultaneous oxidation and reduction of atoms of same element from one oxidation state to two different oxidation states.
$2 \overset{+4}{NO _2}(g)+2 OH ^{-} \rightarrow \overset{+3}{NO _2^{-}}+ \overset{+5}{NO _3^{-}}+ H _2 O$
In this reaction, $N$ of nitrogen dioxide is of $+4$ oxidation state but after the reaction, $N$ is present in $+3$ and $+5$ oxidation state simultaneously. So, this reaction is an example of disproportionation redox reaction.