A proposed mechanism for one of the pathways for the destruction of ozone in the
ID: 484915 • Letter: A
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 (dASHES MEAN ARROW)
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.
____+____ - _____ + ______
(2) Which species acts as a catalyst? Enter formula. If none, leave box blank
______ + _____ - ______ + ______ (dashes mean arrow)
3) Which species acts as a reaction intermediate? 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= ____
Explanation / Answer
Q.
Overall reaction is written by addition of two half reactions.
O3+NO --- > NO2 + O2
NO2 + O --- > NO + O2
-----------------------------
NO + O3 + O + NO2 --- > NO2 + 2 O2 + NO
From above reaction we cancel the common terms and we will get
O3 + O -- > 2 O2
(Overall reaction)
Q. 2
The species that is used in first step and re appeared in the last step is called as catalyst.
Here NO is used in first step and re appeared in second ( last ) step so it’s Catalyst.
Answer: The catalyst is NO
Q. 3
Reaction intermediate
The species is formed in one step and reacted in another step is called as an intermediate.
Answer : NO2 ( formed in first step and utilized in second step)
Q. 4
Rate law of this reaction depends on the slow step (rate determine step)
Since reaction intermediate is formed in first step which is utilized in second step and so second is faster and first is slow step (rate determining step)
We use the coefficients in the elementary steps to write the rate law.
Since the coefficient of O3 is 1 and that of the NO is 1 so the rate law is
Rate = k [O3][NO]
k is rate constant.
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