Question

Given the following two-step mechanism:

Step 1) NO₂(g) ＋Cl₂(g) → ClNO₂(g)＋ Cl (g)
Step 2) NO₂(g) ＋Cl₂(g) → ClNO₂(g)
a) Identify the intermediate
b) What is the overall reaction?
c) If the first step is the slow step, what is the predicted rate law?
d) If the second step is the slow step, what is the predicted rate law?

Answer 1

The intermediate in the two-step mechanism is ClNO₂(g), and the overall reaction is 2NO(g) + Cl₂(g) → 2NOCl(g). Rate laws depend on the slow step: rate = k₁[NO][Cl₂] for a slow first step, and rate = k₂[NOCl][NO] for a slow second step.

Step-by-step explanation:

In the given two-step mechanism, we can identify the intermediate and the overall reaction. The intermediate is a species that is produced in one step and consumed in another. Here, ClNO₂(g) acts as the intermediate because it is formed in step 1 and consumed in step 2. By adding the two steps, the intermediate cancels out, yielding the overall reaction:

2NO(g) + Cl₂(g) → 2NOCl(g)

If the first step is slow, the rate law is determined by the reactants involved in that step, hence it would be:

rate = k₁[NO][Cl₂]

If the second step is slow, the rate law would be:

rate = k₂[NOCl][NO]

However, since the intermediate is typically not included directly in the rate law, and additional information or assumptions are needed to express the rate in terms of reactants only.