Calculate the activation energy of a reaction,whose rate constant doubles on raising the temperature from $300 \ K$ to $600 \ K$.

Based on the provided energy profile diagram,how many steps are in the mechanism of the reaction?

The first order rate constant for a certain reaction increases from $1.667 \times 10^{-6} \ s^{-1}$ at $727 \ ^oC$ to $1.667 \times 10^{-4} \ s^{-1}$ at $1571 \ ^oC$. The rate constant at $1150 \ ^oC$,assuming constancy of activation energy over the given temperature range is [Given : $\log \ 19.9 = 1.299$ ]

$A$ reaction having equal activation energies for forward and reverse reaction has:

For a first order chemical reaction,

In a laboratory,a biochemical reaction performed outside the human body in the absence of an enzyme proceeds at a rate $10^{-6}$ times slower. If this reaction is performed in the presence of an enzyme,what will be the activation energy of the reaction?