Question

For a reaction
$2H_2{O}_2\underset{\text{alkaline medium}}{\overset{I^{-}}{\to }}2H_{2}O+O_2$

The proposed mechanism is given below:

(1) $H_2{O}_2+I^{-}\to H_{2}O+I{O}^{-}\left(slow\right)$
(2) $H_2{O}_2+I{O}^{-}\to H_{2}O+I^{-}+O_2\left(fast\right)$

(i) Write rate law for the reaction.
(ii) Write the overall order of the reaction.
(iii) Out of steps (1) and (2), which one is the rate determining step?

Answer 1

The rate of the decomposition of hydrogen peroxide depends on the concentration of the catalyst [$I^{-}$] also. The reaction is catalysed by iodide ion in alkaline medium.
(i) The rate law for the reaction is:
rate = $\frac{-d}{dt}\left[H_2{O}_2\right]=k\left[H_2{O}_2\right]\left[I^{-}\right]$
(ii) The reaction is first order with respect to both $\left[H_2{O}_2\right]$ and $\left[I^{-}\right]$. So, the overall order is 2.
The reaction takes place in two steps, both are bimolecular elementary reactions. $I{O}^{-}$ is the intermediate formed in the first step which is not in the overall reaction.
(iii) Step 1, being slow, is the rate determining step.
The formation of the intermediate determines the rate of the reaction.