Reaction Energy Profile

Even a reaction that releases energy overall has to climb before it can fall. Molecules collide with random amounts of energy, and most collisions bounce off doing nothing. Only collisions energetic enough to reach the transition state — a strained, half-broken, half-formed arrangement sitting at the top of the hill — actually lead to products. That minimum energy requirement is the activation energy, Ea. It exists because bonds in the reactants have to start breaking before new bonds can finish forming, and that mid-process arrangement costs energy no matter which direction the reaction is headed.

On the diagram, reaction progress runs along the x-axis and potential energy runs up the y-axis. Reactants sit at one energy level, products at another, connected by a curve that rises to a single peak. Forward activation energy is measured from the reactants' energy up to that peak. Enthalpy change, deltaH, is measured from reactants straight across to products, ignoring the hill entirely. In an exothermic profile, products end up lower than reactants and deltaH is negative, even though the reaction still climbed a hill to get there. In an endothermic profile, products sit higher than reactants and deltaH is positive. Running the reaction in reverse just means reading the same curve backward: the reverse activation energy is measured from the products' energy up to that same peak, so deltaH = Ea,forward - Ea,reverse always connects the two paths on one diagram.

This connects directly to the idea of a rate-determining step: whichever elementary step in a mechanism has the tallest hill is the slow one that controls the overall rate, no matter how favorable the overall deltaH turns out to be.