To solve this problem, we need to analyze the potential energy profile for the reaction $A \to B. \newline$Given:$\newline$• Enthalpy of reaction $(\Delta H) = -4.2 \,$kJ mol$^{-1} \newline$• Enthalpy of activation $(E_a) = 9.6 \,$kJ mol$^{-1} \newline$The enthalpy of reaction is negative, indicating that the reaction is exothermic.$\newline$In an exothermic reaction, the energy of the products (B) is lower than that of the reactants (A).$\newline$The activation energy is the energy difference between the reactants and the transition state.$\newline$Let's analyze the options:$\newline$• Option 1: The energy of B is higher than A, indicating an endothermic reaction.$\newline$• Option 2: The energy of B is lower than A, indicating an exothermic reaction.$\newline$• Option 3: The energy of B is equal to A, indicating no net change in energy.$\newline$• Option 4: The energy of B is higher than A, indicating an endothermic reaction.$\newline$Since the reaction is exothermic, the correct potential energy profile is shown in Option 2, where the energy of B is lower than A.