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How does a potential-energy surface represent the path of a chemical reaction?

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A graph showing a potential energy curve for a chemical reaction, with the y-axis labeled potential energy and the x-axis labeled progress of reaction. The curve rises from the initial state (reactants) to a peak labeled transition state (activated complex), indicating activation energy, then drops to the final state (products). — chemical kinetics In the Arrhenius equation, the activation energy (E) represents the minimum amount of energy required to transform reactants into products in a chemical reaction. On a potential energy curve, the value of the activation energy is equivalent to the difference in potential energy between particles in an intermediate configuration (known as the activated complex, or transition state) and particles of reactants in their initial state. The activation energy thus can be visualized as a barrier that must be overcome by reactants before products can be formed.

In chemical kinetics, a potential-energy surface serves as a model that illustrates how the potential energy of a reaction system changes in relation to critical interatomic distances. The progression of an elementary chemical reaction is represented by the movement of the system across this surface.1

As the system transitions from its initial state to the final product, it follows a reaction path that passes through an intermediate configuration known as the activated complex. At this specific point, the potential energy reaches a maximum value, creating an energy barrier. The difference between the energy of this transition state and the initial state is closely related to the activation energy, which represents the minimum energy required for the transformation to occur. Theoretical approaches to understanding reaction rates focus either on the highest point of this energy barrier or on the dynamics of the system's motion across the entire surface.123