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title: "The gas-phase oxidation of nitric oxide proceeds according to the following balanced equation: \\[2\\text{NO}(g) + \\text{O}_2(g) \\rightarrow 2\\text{NO}_2(g)\\] The experimentally determined rate law for the reaction is \\(\\text{rate} = k[\\text{NO}]^2[\\text{O}_2]\\). To explain this rate law, two different reaction mechanisms are proposed:  Mechanism 1 Step 1 (fast equilibrium): \\(\\text{NO}(g) + \\text{O}_2(g) \\rightleftharpoons \\text{NO}_3(g)\\) Step 2 (slow): \\(\\text{NO}_3(g) + \\text{NO}(g) \\rightarrow 2\\text{NO}_2(g)\\)  Mechanism 2 Step 1 (fast equilibrium): \\(\\text{NO}(g) + \\text{NO}(g) \\rightleftharpoons \\text{N}_2\\text{O}_2(g)\\) Step 2 (slow): \\(\\text{N}_2\\text{O}_2(g) + \\text{O}_2(g) \\rightarrow 2\\text{NO}_2(g)\\)  Both mechanisms are consistent with the overall stoichiometry and the experimental rate law. Which of the following pieces of experimental evidence would provide the most decisive support for Mechanism 2 over Mechanism 1?"
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url: "https://nerd-notes.com/ubq/120005/"
date_modified: "2026-08-21T08:31:49+00:00"
---

# The gas-phase oxidation of nitric oxide proceeds according to the following balanced equation:
\[2\text{NO}(g) + \text{O}_2(g) \rightarrow 2\text{NO}_2(g)\]
The experimentally determined rate law for the reaction is \(\text{rate} = k[\text{NO}]^2[\text{O}_2]\). To explain this rate law, two different reaction mechanisms are proposed:

Mechanism 1
Step 1 (fast equilibrium): \(\text{NO}(g) + \text{O}_2(g) \rightleftharpoons \text{NO}_3(g)\)
Step 2 (slow): \(\text{NO}_3(g) + \text{NO}(g) \rightarrow 2\text{NO}_2(g)\)

Mechanism 2
Step 1 (fast equilibrium): \(\text{NO}(g) + \text{NO}(g) \rightleftharpoons \text{N}_2\text{O}_2(g)\)
Step 2 (slow): \(\text{N}_2\text{O}_2(g) + \text{O}_2(g) \rightarrow 2\text{NO}_2(g)\)

Both mechanisms are consistent with the overall stoichiometry and the experimental rate law. Which of the following pieces of experimental evidence would provide the most decisive support for Mechanism 2 over Mechanism 1?

The gas-phase oxidation of nitric oxide proceeds according to the following balanced equation:
\[2\text{NO}(g) + \text{O}_2(g) \rightarrow 2\text{NO}_2(g)\]
The experimentally determined rate law for the reaction is \(\text{rate} = k[\text{NO}]^2[\text{O}_2]\). To explain this rate law, two different reaction mechanisms are proposed:

Mechanism 1
Step 1 (fast equilibrium): \(\text{NO}(g) + \text{O}_2(g) \rightleftharpoons \text{NO}_3(g)\)
Step 2 (slow): \(\text{NO}_3(g) + \text{NO}(g) \rightarrow 2\text{NO}_2(g)\)

Mechanism 2
Step 1 (fast equilibrium): \(\text{NO}(g) + \text{NO}(g) \rightleftharpoons \text{N}_2\text{O}_2(g)\)
Step 2 (slow): \(\text{N}_2\text{O}_2(g) + \text{O}_2(g) \rightarrow 2\text{NO}_2(g)\)

Both mechanisms are consistent with the overall stoichiometry and the experimental rate law. Which of the following pieces of experimental evidence would provide the most decisive support for Mechanism 2 over Mechanism 1?

- **A.** Measuring the overall reaction rate at several different temperatures to determine the activation energy of the reaction.
- **B.** Determining the initial rate of the reaction when the partial pressure of \(\text{O}_2(g)\) is varied at constant \(\text{NO}(g)\).
- **C.** Measuring the standard enthalpy change of the reaction using calorimetry and comparing it to calculated bond energies.
- **D.** Detecting the transient presence of \(\text{N}_2\text{O}_2(g)\) molecules in the reaction mixture using fast spectroscopic methods.

*The answer key and step-by-step explanation are available to logged-in users at https://nerd-notes.com/ubq/120005/*