What is the term for the amount of energy required to bring all molecules in a reaction to the reactive state?

Activation energy is the term that describes the minimum amount of energy needed to initiate a chemical reaction. It represents the energy barrier that reactants must overcome for a reaction to occur. When a sufficient amount of energy is provided to the reactants, their molecules can achieve an activated state, where they can undergo the necessary transformations to form products.

This concept is critical in understanding how reactions proceed, as factors that influence activation energy (such as temperature, catalysts, and concentration of reactants) can significantly affect the reaction rate. For example, higher temperatures can provide molecules with more kinetic energy, increasing the likelihood that they reach the activated state.

In contrast, the other terms listed, like reaction energy, thermal energy, and potential energy, refer to different concepts within thermodynamics and kinetics. Reaction energy does not specifically denote the energy needed to reach the activated state. Thermal energy relates to the temperature and kinetic energy of particles, but does not directly connect to the concept of overcoming an energy threshold for a reaction. Potential energy refers to the stored energy in a system, not the specific energy required to initiate a reaction. Thus, activation energy is the most precise term for describing the energy threshold necessary for a reaction to proceed.