The isomerization reaction CH3NC → CH3CN obeys the first-order rate law, rate = k[CH3NC], in the presence of an excess of argon. Measurement at 500. K reveals that in 492 seconds, the concentration of CH3NC has decreased to 72% of its original value. Calculate the rate constant (k) of the reaction at 500. K.

Respuesta :

Answer:

         [tex]\large\boxed{\large\boxed{k=6.68\times 10^{-4}s^{-1}}}[/tex]

Explanation:

The first-order rate law is:

              [tex]\dfrac{d[CH_3NC]}{[CH_3NC]}=-kt[/tex]

The solution of that differential equation is:

          [tex][CH_3NC]=[CH_3NC]_0e^{-kt}[/tex]

You are given:

       [tex]\dfrac{CH_3NC]}{[CH_3NC]_0}=0.72[/tex]

       [tex]t=492s[/tex]

And you want to determine k.

Substitute into the solution of the differential equation and solve for t:

               [tex]0.72=e^{-k\times 492s}[/tex]

               [tex]k=-ln(0.72)/492s\\\\k=6.6769\times10^{-4}s^{-1}\approx6.68\times 10^{-4}s^{-1}[/tex]