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:
For a multi-step reaction, the overall rate of the reaction is determined by the slowest step, known as the Rate Determining Step (RDS).
Step 2: Detailed Explanation:
(a) Rate Law:
The first step is labeled as "slow". Therefore, it is the RDS.
The reactants in the RDS are \( NO_2Cl \). The rate law is written based on the stoichiometry of the RDS.
\[ \text{Rate} = k[NO_2Cl] \]

(b) Reaction Intermediate:
A reaction intermediate is a species that is produced in one step and consumed in a subsequent step, and does not appear in the overall balanced equation.
In the given mechanism, atomic chlorine \( Cl(g) \) is produced in step (i) and consumed in step (ii).
Therefore, \( Cl(g) \) is the reaction intermediate.
Step 3: Final Answer:
Rate Law: \( R = k[NO_2Cl] \). Intermediate: \( Cl \).