For an endothermic reaction,where $\Delta H$ represents the enthalpy of the reaction in $kJ/mol$,the minimum value for the energy of activation will be

For a first-order reaction $A \rightarrow P$,the rate constant $k$ depends on temperature $(T)$ according to the equation $\log \, k = - \frac{2000}{T} + 0.6$. The pre-exponential factor $A$ and the activation energy $E_a$ are,respectively:

For a reaction $A \rightarrow B$,the enthalpy of reaction is $-4.2 \ kJ \ mol^{-1}$ and the enthalpy of activation is $9.6 \ kJ \ mol^{-1}$. The correct potential energy profile for the reaction is shown in which option?

The reason for almost doubling the rate of reaction on increasing the temperature of the reaction system by $10 \, ^\circ C$ is

Define the following terms:
$(a)$ Activation energy
$(b)$ Activated complex

$A$ large increase in the rate of reaction for a small increase in temperature is due to: