The correct answer is D (There are two intermediates and three transition states in the overall reaction).
The graph depicts the reaction pathway from reactant A to product D. Reactant A is less stable than product D, as indicated by its lower position on the y-axis. The peaks represent activated complexes, which are less stable than both the reactants and products. Points B and C denote transition states, which have lower energy than the activated complexes but are less stable than the products in terms of energy expenditure.
Specifically, an activated complex forms transiently during collisions and is entirely unstable, capable of either dissociating back into reactants or forming products. Transition states are characterized by the highest energies along the reaction profile, while intermediates are relatively stable species involved in the reaction mechanism but are neither reactants nor appear in the reaction rate law. According to the profile, the reaction mechanism involves three transition states (points of maximum energy) and two intermediates (points B and C).
Thus, option (D) is correct. Activation energy represents the minimum collision energy required for a chemical reaction to initiate, acting as a barrier separating reactants from products. Only a fraction of colliding molecules possess adequate kinetic energy to surmount this barrier. In a reaction profile, the activation energy for a forward reaction equals the energy difference between an activated complex and a reactant.
Option (A) is incorrect because the enthalpy change (ΔH) of the reaction reflects the energy difference between products and reactants, indicating an exothermic reaction (ΔH < 0) due to lower product energy. Option (B) is incorrect because the fastest step in the reaction pathway is C → D, not B → C, owing to its lower activation energy. Option (C) is incorrect because the rate-determining step in the overall reaction is A → B, not B → C, due to its higher activation energy.