Suppose that a certain biologically important reaction is quite slow at physiological temperature (37 oc) in the absence of a catalyst. assuming arrhenius behavior, by how much must an enzyme lower the activation energy of the reaction to achieve a 1 x 105-fold increase in the reaction rate? (give your answer in kj)

30 kJ

Explanation:

Arrhenius equation is given by:

Here, k is rate constant, A is Pre-exponential factor, Ea is activation energy and T is temperature.

taking natural log of both side

ln k = ln A - Ea/RT

In Arrhenius equation, A, R and T are constant.

Therefore,

is the  lowering in activation energy by enzyme,

R = 8.314 J/mol.K

T = 37°C + 273.15 = 310 K

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