The rate constants for the first-order decomposition of a compound are 5.22× 10–4 s–1 at 43°c and 2.91 × 10–3 s–1 at 62°c. what is the value of the activation energy for this reaction? (r = 8.31 j/(mol · k)) a. 79.5 kj/mol b. 34.5 kj/mol c. 0.751 kj/mol d. 0.87104 kj/mol e. 2 kj/mol

Activation energy of the reaction is 79.5 kJ/mol

Explanation:

According to Arrhenius equation for a reaction-

where k is the rate constant, A is the Arrhenius constant, is the activation energy and T is temperature in kelvin

For the given two different set of condition, we can write-

at , ............(1)

at , ............(2)

gives-

Solving this equation we get

So activation energy of the reaction is 79.5 kJ/mol

i know insulin is primary but i dont know the others im sorry

explanation: