Question

For a reaction:

\( 2 \text{H}_2\text{O}_2 \xrightarrow{\text{I}} 2 \text{H}_2\text{O} + \text{O}_2 \)

The proposed mechanism is as given below:

(I) \( \text{H}_2\text{O}_2 \xrightarrow{\text{slow}} \text{H}_2\text{O} + \text{IO}^- \) (slow)

(II) \( \text{H}_2\text{O}_2 + \text{IO}^- \xrightarrow{\text{fast}} \text{H}_2\text{O} + \text{I}^+ + \text{O}_2 \) (fast)

Answer 1

(I) The rate-determining step is the first reaction, where hydrogen peroxide decomposes to form water and IO⁻. The rate law for the reaction will be dependent on the concentration of H₂O₂.

(II) In the second step, the intermediate IO⁻ reacts with H₂O₂ in a fast step to produce water, iodine ions, and oxygen gas. This step does not affect the rate law, as it is not rate-determining.

(1) Rate law: The rate law is determined by the slow step, so the rate law is:

Rate = \( k[\text{H}_2\text{O}_2] \)

where \( k \) is the rate constant.

(2) Overall order and molecularity: The overall order of the reaction is 1, as the rate law depends on the concentration of only one reactant, H₂O₂. The molecularity of the reaction is 2, as the rate-determining step involves the collision of two molecules of H₂O₂.