Question

The rate constant for the reaction below was determined to be 3.241×10-5 s–1 at 800 K....

The rate constant for the reaction below was determined to be 3.241×10-5 s–1 at 800 K. The activation energy of the reaction is 215 kJ/mol. What would be the value of the rate constant at 9.10×102 K? N2O(g) --> N2(g) + O2(g)

I'm having trouble calculating the rate constant with the arrhenius equation that deals with two temps, could you show me the step by step how to do this?

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Homework Answers

Answer #1

Ans :- 1.6 x 10-3 s-1

Explanation :-

Given,

rate constant at 800 K = k1 = 3.241 x 10-5 s-1

Temperature = T1 = 800 K

rate constant at 9.10 x 102 K = k2 =?

Temperature = T2 = 9.10 x 102 K

Activation energy = Ea = 215 KJ = 215000 J/mol

From the Arrhenius equation

ln (k2/k1) = (Ea/R) . (1/T1 - 1/T2)

ln (k2/k1) = 215000 J/mol / 8.314 J K-1 mol-1 .(9.10 x 102 - 800 K / (800 K)(9.10 x 102)

ln (k2/k1) = 3.9074

k2 = k1 x exp ( 3.9075)

k2 = 3.241 x 10-5 x 49.7694

k2 = 1.61 x 10-3 s-1

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