Summary

This chapter developed chemical kinetics from first principles. It began by defining

average and instantaneous rate, and the single unambiguous "rate of reaction" obtained by dividing each

species' rate of change by its own stoichiometric coefficient. It then established that a reaction's

rate law and order must always be determined experimentally -- never assumed from the balanced

equation -- and carefully distinguished this empirical order from molecularity, the theoretical count

of colliding species in a single elementary mechanistic step. The integrated rate equations for zero

order ($[R] = [R]_0 - kt$) and first order ($k = (2.303/t)\log([R]_0/[R])$) reactions were derived, along

with their half-lives, revealing the key diagnostic difference that a zero order half-life depends on

initial concentration while a first order half-life, $t_{1/2} = 0.693/k$, does not. Collision theory

and the concept of activation energy then supplied the molecular-level explanation for why reactions

proceed at the rates they do, and why those rates rise so sharply with temperature -- a dependence

captured precisely, and made experimentally measurable, by the Arrhenius equation, $k = Ae^{-E_a/RT}$,

in both its graphical (single-line, multi-temperature) and two-point calculation forms. The chapter …