Question

A chemical reaction occurs in two steps:

\[\ce{NO2Cl_{(g)} ->[slow] NO2_{(g)} + Cl_{(g)}}\]

\[\ce{NO2Cl_{(g)} + Cl_{(g)} ->[fast] NO2_{(g)} + Cl2_{(g)}}\]

1. Write down the rate law.
2. Identify the reaction intermediate.

Answer 1

a. Since the first step is slow, it is the rate-determining step.

Slow step:

\[\ce{NO2Cl -> NO2 + Cl}\]

The rate depends on the reactant in the slow step.

Rate = k[NO₂Cl]

b. An intermediate is a species that is:

• Formed in one step
• Consumed in another step
• Does not appear in the overall reaction

Here, Cl (chlorine atom) is formed in step 1 and consumed in step 2.

Reaction intermediate = Cl_(g)