The above diagram gives the potential energy profile of the catalyst and uncatalysed reactions of X(g) + Y(g) → XY(g). Deduce the respective activation ener...

Answer Details

The activation energy is the minimum energy required for a chemical reaction to occur. In the given diagram, the activation energy for the uncatalyzed reaction is the difference between the energy of the reactants and the maximum point on the curve representing the transition state. Similarly, the activation energy for the catalyzed reaction is the difference between the energy of the reactants and the maximum point on the curve representing the transition state in the presence of the catalyst. From the diagram, the activation energy for the uncatalysed reaction is 500 kJ (difference between the energy of the reactants and the maximum point on the curve representing the transition state for the uncatalysed reaction). The activation energy for the catalysed reaction is 300 kJ (difference between the energy of the reactants and the maximum point on the curve representing the transition state in the presence of the catalyst). Therefore, the activation energy for the uncatalysed reverse reaction (XY(g) → X(g) + Y(g)) would also be 500 kJ (same as the uncatalysed forward reaction). Similarly, the activation energy for the catalysed reverse reaction (XY(g) → X(g) + Y(g)) would also be 300 kJ (same as the catalysed forward reaction). So, the respective activation energies in kJ of the catalysed and uncatalysed reverse reactions are 300 and 500 kJ respectively.