---
title: "A student investigates the gas-phase reaction represented by the following balanced equation:  \\[ \\text{NO}_2\\text{(g)} + \\text{CO(g)} \\rightarrow \\text{NO(g)} + \\text{CO}_2\\text{(g)} \\]  Experimental initial-rate data show that the reaction is second order with respect to \\(\\text{NO}_2\\) and zero order with respect to \\(\\text{CO}\\), giving the rate law \\(\\text{Rate} = k[\\text{NO}_2]^2\\). Two proposed reaction mechanisms are shown below.  Mechanism I: Step 1: \\(\\text{NO}_2\\text{(g)} + \\text{NO}_2\\text{(g)} \\rightarrow \\text{NO}_3\\text{(g)} + \\text{NO(g)}\\) (slow) Step 2: \\(\\text{NO}_3\\text{(g)} + \\text{CO(g)} \\rightarrow \\text{NO}_2\\text{(g)} + \\text{CO}_2\\text{(g)}\\) (fast)  Mechanism II: Step 1: \\(\\text{NO}_2\\text{(g)} + \\text{CO(g)} \\rightarrow \\text{NO}_2\\text{CO(g)}\\) (slow) Step 2: \\(\\text{NO}_2\\text{CO(g)} \\rightarrow \\text{NO(g)} + \\text{CO}_2\\text{(g)}\\) (fast)  Which mechanism is consistent with the experimental rate law, and why?"
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date_modified: "2026-09-28T12:01:54+00:00"
---

# A student investigates the gas-phase reaction represented by the following balanced equation:

\[ \text{NO}_2\text{(g)} + \text{CO(g)} \rightarrow \text{NO(g)} + \text{CO}_2\text{(g)} \]

Experimental initial-rate data show that the reaction is second order with respect to \(\text{NO}_2\) and zero order with respect to \(\text{CO}\), giving the rate law \(\text{Rate} = k[\text{NO}_2]^2\). Two proposed reaction mechanisms are shown below.

Mechanism I:
Step 1: \(\text{NO}_2\text{(g)} + \text{NO}_2\text{(g)} \rightarrow \text{NO}_3\text{(g)} + \text{NO(g)}\) (slow)
Step 2: \(\text{NO}_3\text{(g)} + \text{CO(g)} \rightarrow \text{NO}_2\text{(g)} + \text{CO}_2\text{(g)}\) (fast)

Mechanism II:
Step 1: \(\text{NO}_2\text{(g)} + \text{CO(g)} \rightarrow \text{NO}_2\text{CO(g)}\) (slow)
Step 2: \(\text{NO}_2\text{CO(g)} \rightarrow \text{NO(g)} + \text{CO}_2\text{(g)}\) (fast)

Which mechanism is consistent with the experimental rate law, and why?

A student investigates the gas-phase reaction represented by the following balanced equation:

\[ \text{NO}_2\text{(g)} + \text{CO(g)} \rightarrow \text{NO(g)} + \text{CO}_2\text{(g)} \]

Experimental initial-rate data show that the reaction is second order with respect to \(\text{NO}_2\) and zero order with respect to \(\text{CO}\), giving the rate law \(\text{Rate} = k[\text{NO}_2]^2\). Two proposed reaction mechanisms are shown below.

Mechanism I:
Step 1: \(\text{NO}_2\text{(g)} + \text{NO}_2\text{(g)} \rightarrow \text{NO}_3\text{(g)} + \text{NO(g)}\) (slow)
Step 2: \(\text{NO}_3\text{(g)} + \text{CO(g)} \rightarrow \text{NO}_2\text{(g)} + \text{CO}_2\text{(g)}\) (fast)

Mechanism II:
Step 1: \(\text{NO}_2\text{(g)} + \text{CO(g)} \rightarrow \text{NO}_2\text{CO(g)}\) (slow)
Step 2: \(\text{NO}_2\text{CO(g)} \rightarrow \text{NO(g)} + \text{CO}_2\text{(g)}\) (fast)

Which mechanism is consistent with the experimental rate law, and why?

- **A.** Mechanism I, because \(\text{NO}_3\text{(g)}\) acts as a catalyst in Step 2, which increases the reaction rate without altering the reaction order of \(\text{CO(g)}\).
- **B.** Mechanism I, because the rate-determining Step 1 involves a bimolecular collision between two \(\text{NO}_2\) molecules, yielding \(\text{Rate} = k_1[\text{NO}_2]^2\) independent of \([\text{CO}]\).
- **C.** Mechanism II, because \(\text{CO(g)}\) is a reactant in the overall balanced equation and must participate in the rate-determining step for the reaction to occur.
- **D.** Mechanism II, because the \(1:1\) stoichiometric ratio of \(\text{NO}_2\text{(g)}\) to \(\text{CO(g)}\) in Step 1 matches the coefficients in the overall balanced chemical equation.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/123653/*