224 Chemical Kinetics
A study of how fast chemical reactions occur that is used to determine reaction pathways and control how quickly products form.
Elementary reactions and molecularity. Complex reactions proceed through a sequence of discrete molecular events. An elementary reaction is a single collision event that forms products without a lasting intermediate. Molecularity is the number of reactant molecules that must meet in that event. This principle is used to propose mechanisms and to derive overall rate laws that experiments can test.
The Arrhenius activation barrier. Colliding molecules react only when their approach energy exceeds a minimum threshold. Activation energy is that minimum energy needed to start bond breaking. Raising temperature increases the fraction of molecules that clear the barrier and speeds the reaction. This principle is used to explain and predict how strongly reaction speed depends on temperature.
Transition-state theory. During a reactive collision the potential energy rises to a peak, then the atoms form products; otherwise they fall back to reactants. The transition state is the geometry at that energy peak. This principle is used to calculate rate constants from the properties of the reactants and of the transition state.
The rate-determining-step approximation. In a multi-step mechanism the elementary step with the smallest rate constant limits how fast products can form. That slowest step is the rate-determining step. This principle is used to simplify the overall rate law by focusing on the bottleneck step.
The steady-state approximation. Highly reactive intermediates are destroyed as fast as they are created, so their amounts stay tiny and nearly constant. A reaction intermediate is a short-lived species made in one step and consumed in another that does not appear in the overall equation. This principle is used to solve multi-step rate equations when no single step is the clear bottleneck.
The Lindemann-Hinshelwood mechanism of unimolecular reactions. A single molecule does not decompose on its own; it must first gain vibrational energy in a collision. An energized molecule has that energy but has not yet broken the bond. At low pressure a later collision can remove the energy before the bond breaks. This principle is used to explain how gas-phase unimolecular rates depend on pressure.
224.1 References
- Levine, I. N. Physical Chemistry. Ch. 16 §16.5 — elementary reactions. §16.6 — rate-determining step and steady state. §16.8 — activation energy. §16.11 — unimolecular reactions. Ch. 22 §22.4 — transition-state theory.
- Atkins, P., de Paula, J., & Keeler, J. Atkins’ Physical Chemistry. Topic 17D — activation. Topic 17E — mechanisms. Topic 17F — unimolecular reactions. Topic 18C — transition-state theory.