Nitric oxide is oxidized in air through the multistep mechanism outlined below.
Step 1 2 NO ⇌ N₂O₂ (fast equilibrium)
Step 2 N₂O₂ + O₂ → 2 NO₂ (slow)
Step 3 NO₂ + NO → N₂O₃ (fast)
Assume that the observed rate of NO oxidation is governed by the elementary process occurring in Step 2.
(a) For each elementary step, state its molecularity.
(b) Write the rate law expression associated with the rate-determining step.
(c) Determine the individual reaction orders with respect to N₂O₂ and O₂ that appear in the rate law from part (b), and state the overall reaction order for the oxidation process.
(d) Experimental studies sometimes report a rate law that is first order in NO and first order in O₂ for the overall oxidation of NO. Propose a chemical explanation, based on the mechanism above, for how this experimentally observed rate law can be reconciled with the rate law you obtained in part (b).