The rate of a reaction:

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 :
$(R=8.314 \, J \, mol^{-1} \, K^{-1})$

Which of the following options is incorrect?

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

The activation energy for the forward reaction is $150 \ kJ \ mol^{-1}$ and that of the reverse reaction is $260 \ kJ \ mol^{-1}$. The enthalpy change for the reaction is:

The activation energy of a reaction is $0$. The rate constant of the reaction: