The reaction is:

$$\ce{NO(g) + NO2(g) + O2(g) -> N2O5(g)}$$

From the experimental data, I have found the rate law to be: $$\ce{Rate~=~k[NO][NO2]}$$

This means the overall reaction does not depend on $\ce{O2(g)}$.

Can someone suggest how I go about building the elementary steps to form the reaction mechanism?

  • $\begingroup$ Based on the rate law, which species do you think must be involved in the rate determining step? $\endgroup$ – bon Sep 6 '15 at 17:37
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    $\begingroup$ @bon should be NO & NO2 $\endgroup$ – justbehappy Sep 6 '15 at 17:49
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    $\begingroup$ Correct. So you should be able to construct a two step mechanism using this information. $\endgroup$ – bon Sep 6 '15 at 17:56
  • $\begingroup$ Let me try....is the first step: NO(g)+NO2(g)-> N2O3 (g) then the 2nd step: N2O3(g)+O2(g)->N2O5 (g)? $\endgroup$ – justbehappy Sep 6 '15 at 18:03
  • $\begingroup$ @bon how does this reaction mechanism show that O2 is not part of the rate law? $\endgroup$ – justbehappy Sep 6 '15 at 18:04

First step - the rate determining step or the slow step (activation energy is high). $$\ce{NO(g) + NO2(g) -> N2O3(g)}$$

2nd step - the fast step. $$\ce{N2O3(g) + O2(g) -> N2O5(g)}$$

Since the rate determining step does not include $\ce{O2(g)}$, the rate is independent of the initial $\ce{O2(g)}$ concentration.

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