The rate constant is doubled when temperature increases from $27^{\circ} C$ to $37^{\circ} C$. What is the activation energy in $kJ \ mol^{-1}$?

The rate of a reaction quadruples when the temperature changes from $293 \ K$ to $313 \ K$. Calculate the energy of activation of the reaction assuming that it does not change with temperature.

The rate of a reaction doubles,when the temperature is changed from $300 \ K$ to $310 \ K$. Activation energy of the reaction is....... $(R=8.314 \ J \ K^{-1} \ mol^{-1}, \log 2=0.301)$

For the reaction,$aA + bB \rightarrow cC + dD$,the plot of $\log k$ vs $\frac{1}{T}$ is given below. The temperature at which the rate constant of the reaction is $10^{-4} \ s^{-1}$ is ............... $K$. (Rounded-off to the nearest integer) [Given: The rate constant of the reaction is $10^{-5} \ s^{-1}$ at $500 \ K$.]

The variation of the rate constant with temperature is given by the Arrhenius equation $k = A e^{-E_a / (RT)}$. If $T \to \infty$,the rate constant $k$ will be equal to:

The minimum energy a molecule should possess in order to enter into a fruitful collision is known as