Minimum energy required to start a reaction is called

• A. Catalysis
• B. Activation energy
• C. Collision energy
• D. Transition state energy

Answer: (B)

Activation energy is the minimum energy required to start a reaction. It represents the energy barrier that reactants must overcome to reach the transition state and proceed to products. On a potential energy diagram, this is the difference between the energy of the reactants and the highest-energy point along the reaction path—the transition state. Not every collision provides enough energy, and only those with energy at least equal to the activation energy can lead to reaction; temperature affects how many molecules meet this threshold, and this is captured by the Arrhenius relationship. Catalysts work by offering an alternative pathway with a lower transition-state energy, thereby lowering the activation energy and increasing the reaction rate. The energy of the transition state itself is higher than the reactants by the amount of the activation energy, which is why Ea is the key quantity describing the minimum energy needed to initiate the reaction.