 Transition state theory

Transition state theory relates an elementary reaction's rate constant to a Gibbs energy of activation. The Eyring expression used in the catalyst example assumes every passage counted at the transition state succeeds. More generally, a factor accounting for unsuccessful passage multiplies it. Comparing barriers alone gives the quoted speedup only when the remaining prefactors agree. Free-energy profiles can depend on temperature, so the example's values at $300\mathrm{K}$ do not determine a temperature curve. See the IUPAC Gold Book: transition state theory.

Transition state theory (TST), also known as activated complex theory, is a theoretical framework in chemical kinetics that describes the rates of chemical reactions. The main idea behind this theory is that during a reaction, reactants must pass through a high-energy state called the "transition state" or "activated complex" before transforming into products.