The activation energy for a reaction which doubles the rate when the temperature is raised from $298 \ K$ to $308 \ K$ is ........... $kJ \ mol^{-1}$

Write the Arrhenius equation representing the relationship between the rate constants $k_1$ and $k_2$ at two different temperatures $T_1$ and $T_2$.

Two reactions $R_1$ and $R_2$ have identical pre-exponential factors. Activation energy of $R_1$ exceeds that of $R_2$ by $10 \, kJ \, mol^{-1}.$ If $k_1$ and $k_2$ are rate constants for reactions $R_1$ and $R_2$ respectively at $300 \, K,$ then $\ln (k_2/k_1)$ is equal to :
$(R=8.314 \, J \, mol^{-1} \, K^{-1})$

The activation energy for the reaction $2 HI_{(g)} \rightarrow H_{2(g)} + I_{2(g)}$ is $209.5 \ kJ \ mol^{-1}$ at $581 \ K$. Calculate the fraction of molecules of reactants having energy equal to or greater than activation energy?

The rate of reaction increases with a rise in temperature because of:

What is threshold energy?