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title: "The gas-phase oxidation of nitric oxide is represented by the following overall equation:  \\[ 2\\,\\text{NO}(g) + \\text{O}_2(g) \\rightarrow 2\\,\\text{NO}_2(g) \\]  A researcher proposes the following two-step mechanism for the reaction:  Step 1 (fast equilibrium): \\(2\\,\\text{NO}(g) \\underset{k_{-1}}{\\overset{k_1}{\\rightleftharpoons}} \\text{N}_2\\text{O}_2(g)\\)  Step 2 (slow): \\(\\text{N}_2\\text{O}_2(g) + \\text{O}_2(g) \\xrightarrow{k_2} 2\\,\\text{NO}_2(g)\\)  The rate constants for the elementary steps are \\(k_1 = 4.0 \\times 10^3\\text{ M}^{-1}\\text{s}^{-1}\\), \\(k_{-1} = 2.0 \\times 10^4\\text{ s}^{-1}\\), and \\(k_2 = 3.0 \\times 10^2\\text{ M}^{-1}\\text{s}^{-1}\\). Based on the proposed mechanism, what is the initial rate of the reaction when \\([\\text{NO}] = 0.020\\text{ M}\\) and \\([\\text{O}_2] = 0.050\\text{ M}\\)?"
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url: "https://nerd-notes.com/ubq/119706/"
date_modified: "2026-08-21T08:12:06+00:00"
---

# The gas-phase oxidation of nitric oxide is represented by the following overall equation:

\[ 2\,\text{NO}(g) + \text{O}_2(g) \rightarrow 2\,\text{NO}_2(g) \]

A researcher proposes the following two-step mechanism for the reaction:

Step 1 (fast equilibrium): \(2\,\text{NO}(g) \underset{k_{-1}}{\overset{k_1}{\rightleftharpoons}} \text{N}_2\text{O}_2(g)\)

Step 2 (slow): \(\text{N}_2\text{O}_2(g) + \text{O}_2(g) \xrightarrow{k_2} 2\,\text{NO}_2(g)\)

The rate constants for the elementary steps are \(k_1 = 4.0 \times 10^3\text{ M}^{-1}\text{s}^{-1}\), \(k_{-1} = 2.0 \times 10^4\text{ s}^{-1}\), and \(k_2 = 3.0 \times 10^2\text{ M}^{-1}\text{s}^{-1}\). Based on the proposed mechanism, what is the initial rate of the reaction when \([\text{NO}] = 0.020\text{ M}\) and \([\text{O}_2] = 0.050\text{ M}\)?

The gas-phase oxidation of nitric oxide is represented by the following overall equation:

\[ 2\,\text{NO}(g) + \text{O}_2(g) \rightarrow 2\,\text{NO}_2(g) \]

A researcher proposes the following two-step mechanism for the reaction:

Step 1 (fast equilibrium): \(2\,\text{NO}(g) \underset{k_{-1}}{\overset{k_1}{\rightleftharpoons}} \text{N}_2\text{O}_2(g)\)

Step 2 (slow): \(\text{N}_2\text{O}_2(g) + \text{O}_2(g) \xrightarrow{k_2} 2\,\text{NO}_2(g)\)

The rate constants for the elementary steps are \(k_1 = 4.0 \times 10^3\text{ M}^{-1}\text{s}^{-1}\), \(k_{-1} = 2.0 \times 10^4\text{ s}^{-1}\), and \(k_2 = 3.0 \times 10^2\text{ M}^{-1}\text{s}^{-1}\). Based on the proposed mechanism, what is the initial rate of the reaction when \([\text{NO}] = 0.020\text{ M}\) and \([\text{O}_2] = 0.050\text{ M}\)?

- **A.** \(1.2 \times 10^{-3}\text{ M}\cdot\text{s}^{-1}\)
- **B.** \(3.0 \times 10^{-2}\text{ M}\cdot\text{s}^{-1}\)
- **C.** \(6.0 \times 10^{-2}\text{ M}\cdot\text{s}^{-1}\)
- **D.** \(2.4 \times 10^1\text{ M}\cdot\text{s}^{-1}\)

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/119706/*