Question

A chemical reaction occurs in two steps :
(i) \( NO_2Cl(g) \xrightarrow{\text{slow} NO_2(g) + Cl(g) }\)
(ii) \( NO_2Cl(g) + Cl(g) \xrightarrow{\text{fast}} NO_2(g) + Cl_2(g) \)
(a) Write down the rate law.
(b) Identify the reaction intermediate.

Hint:

Intermediate vs. Catalyst: An intermediate is formed then consumed. A catalyst is consumed then regenerated. Both don't appear in the final balanced equation.

Answer 1

Step 1: Understanding the Concept:
In a multi-step reaction mechanism, the overall rate of the reaction is determined by the slowest step, known as the Rate Determining Step (RDS).
Step 3: Detailed Explanation:
(a) Rate Law:
Step (i) is the slow step. Therefore, it determines the rate.
The rate depends on the concentration of the reactants in the slow step.
Rate law: \(Rate = k [NO_2Cl]\).
(b) Reaction Intermediate:
An intermediate is a species that is produced in one step and consumed in a subsequent step, and thus does not appear in the overall balanced equation.
In this mechanism, \(Cl(g)\) is formed in step (i) and used up in step (ii).
Intermediate: \(Cl(g)\) (atomic chlorine).
Step 4: Final Answer: The rate law is \(Rate = k [NO_2Cl]\) and the intermediate is \(Cl(g)\).