What is the activation energy $(kJ \, mol^{-1})$ for a reaction if its rate constant doubles when the temperature is raised from $300 \, K$ to $400 \, K$ ? $(R = 8.314 \, J \, mol^{-1} \, K^{-1})$

Rate constant of a reaction is given as $\log_{10} K = -\frac{2000}{T} + 6.0$. Activation energy is ....... $kcal$.

For the reaction $A_2 + B_2 \rightleftharpoons 2 AB$,the activation energies $(E_a)$ for the forward and backward reactions are $180 \ kJ \ mol^{-1}$ and $200 \ kJ \ mol^{-1}$ respectively. If a catalyst lowers the $E_a$ for both reactions by $100 \ kJ \ mol^{-1}$,which of the following statements is correct?

Activation energy for the hydrolysis of sucrose by acid is $X \ kJ \ mol^{-1}$ whereas activation energy for the hydrolysis of sucrose by sucrase is $Y \ kJ \ mol^{-1}$. $X$ and $Y$ respectively are

In most cases,for a rise of $10 \ K$ temperature,the rate constant is doubled or tripled. This is due to the reason that

Find the correct equation among the following.