Question

A reaction of importance in the formation of smog is that between ozone and nitrogen monoxide described by O3(g)+NO(g)⟶O2(g)+NO2(g) O3(g)+NO(g)⟶O2(g)+NO2(g) The rate law for this reaction is rate of reaction=k[O3][NO] rate of reaction=k[O3][NO] Given that k=4.09×106 M−1⋅s−1k=4.09×106 M−1⋅s−1 at a certain temperature, calculate the initial reaction rate when [O3][O3] and [NO][NO] remain essentially constant at the values [O3]0=5.84×10−6 M[O3]0=5.84×10−6 M and [NO]0=8.65×10−5 M,[NO]0=8.65×10−5 M, owing to continuous production from separate sources.

Answer 1

Answer:

Initial rate of reaction is $2.07\times 10^{-3}M.s^{-1}$.

Explanation:

It is a second order reaction.

Initial rate of reaction = $k[O_{3}]_{0}[NO]_{0}$   , where k is rate constant, $[O_{3}]_{0}$ is the initial concentration of $O_{3}$ and $[NO]_{0}$ is the initial concentration of NO.

Here, k = $4.09\times 10^{6}M^{-1}.s^{-1}$, $[O_{3}]_{0}=5.84\times 10^{-6}M$ and $[NO]_{0}=8.65\times 10^{-5}M$

So, initial rate of reaction = $(4.09\times 10^{6}M^{-1}.s^{-1})\times (5.84\times 10^{-6}M)\times (8.65\times 10^{-5}M)$

= $2.07\times 10^{-3}M.s^{-1}$

So, initial rate of reaction is $2.07\times 10^{-3}M.s^{-1}$