The reaction of NO with O2 to give NO2 is an important step in the commercial production of HNO3. — Kinetics Chemistry Question
Problem Context
The reaction of NO with O2 to give NO2 is an important step in the commercial production of HNO3.
Describe an experiment to measure the rate of this reaction.
Model Answer
The stoichiometry of the reaction is: 2NO(g) + O2(g) → 2NO2(g) so there is a change in the number of moles of gas as the reaction goes forward. This reaction can be monitored by measuring the total pressure of the system as a function of time. Alternatively, the appearance of the red color of NO2 can be measured (e.g. with a spectropohotometer.)
If the rate equation is found to be Rate = k[NO]2[O2], give the effect on the rate of tripling the concentration of.
i. NO
Model Answer
Because the rate law is Rate = k[NO]2[O2], tripling the concentration of NO will cause the rate to increase by a factor of 9.
ii. O2
Model Answer
Tripling the concentration of O2 will cause the rate to increase by a factor of 3.
These two mechanisms have been proposed for this reaction,
I 2NO + O2 → 2NO2
II 2NO → N2O2
N2O2 + O2 → 2NO2
i. State and explain which of the two mechanisms is more likely.
Model Answer
Mechanism is II more likely because mechanism one involves a trimolecular collision. Such a collision is uncommon. By contrast, II has a pair of bimolecular reactions which are considerably more likely to occur.
ii. State and explain which of the two steps in mechanism II must be the slow step if this mechanism is to be consistent with the rate law in b.
Model Answer
Step 2 must be the slow step because it would have a rate law of Rate=k[N2O2][O2], but N2O2 is an intermediate whose concentration arises from the first step. With a slow second step, the first step achieves equilibrium, so [N2O2] = K[NO]2 and the overall rate law would be, Rate = k[NO]2[O2].