Question

A proposed mechanism for one of the pathways for the destruction of ozone in the atmosphere is:

step 1 slow:   O₃ + NO ---->NO₂ + O₂

step 2 fast:   NO₂ + O -----> NO + O₂

(1) What is the equation for the overall reaction? Use the smallest integer coefficients possible. If a box is not needed, leave it blank.

_____+_____----->_____
(3) Which species acts as a reaction intermediate? Enter formula. If none, leave box blank:(2) Which species acts as a catalyst? Enter formula. If none, leave box blank:

(4) Complete the rate law for the overall reaction that is consistent with this mechanism.

(Use the form k[A]^m[B]ⁿ... , where '1' is understood (so don't write it) for m, n etc.)

Rate = _______?

Answer 1

Answer to the first part

In order to make the two equations to one single form, we need to add both the equations

That is, now the equation will be

            O₃ + NO+ NO₂ + O ---->NO₂ + O₂ +NO + O₂

As per the updated system of balancing, the same compounds on either side can be taken out.

After balancing, the final equation will be

O_3(g)+ O(g) ---->2O_2(g)

Answer to the second and third question

3) NO (nitrous oxide) is the catalyst in this reaction

NO + O₃ -> NO₂+ O₂

2) NO₂ reaction intermediate in this reaction

O₃ + NO ---->NO₂ + O₂

Answer to the fourth question

Lets split the reaction into two,

O₃----> O₂ +O ( Rapid equilibrium. ) K 1

O_3(g)+ O(g) ---->2O_2(g) ( rate determining step) K2

As the slowest step is the rate determining step we are considering second step only.

r = r2 = K2[O3][O].

r = K2.[O3]. K . [O3]/ [O2] the concentration of the nascent oxygen is eliminated as it is an intermediate.

r = K.K2.[O3]² / [O2]                           

This reaction is an example of first order reaction.