Chemical Kinetics
Chemical kinetics studies how fast reactions proceed and why, connecting measured rates to the sequence of molecular steps and the energy barriers that control them.
Definition
Chemical kinetics is the branch of physical chemistry concerned with the rates of chemical reactions, the factors that influence them, and the molecular mechanisms by which reactants are converted to products.
Scope
This area covers the measurement and interpretation of reaction rates: empirical rate laws and reaction orders; integrated rate equations and half-lives; the elucidation of reaction mechanisms from elementary steps using the steady-state and pre-equilibrium approximations; the temperature dependence of rates through the Arrhenius equation; and the theoretical frameworks of collision theory and transition state theory. Catalysis and its effect on rates is included, while the detailed quantum dynamics of single collisions and the thermodynamic equilibrium position are treated in neighbouring areas.
Sub-topics
Core questions
- How are rate laws and reaction orders determined experimentally?
- How can a sequence of elementary steps be reconstructed from an observed rate law?
- Why does reaction rate depend so strongly on temperature?
- How do collision theory and transition state theory account for absolute rate constants?
Key concepts
- Rate laws and reaction order
- Integrated rate equations and half-life
- Reaction mechanism and rate-determining step
- Arrhenius equation and activation energy
- Catalysis
Key theories
- Transition state theory
- Reaction rate is calculated from the population of an activated complex in quasi-equilibrium with reactants at the top of the energy barrier, multiplied by the rate at which that complex breaks apart toward products.
- Arrhenius temperature dependence
- Rate constants increase with temperature according to an exponential dependence on an activation energy, reflecting the fraction of collisions energetic enough to surmount the reaction barrier.
Clinical relevance
Chemical kinetics underpins the design and control of industrial reactors, the formulation of catalysts, the stability and shelf life of pharmaceuticals and foods, the modelling of atmospheric and combustion chemistry, and the quantitative analysis of enzyme action in biochemistry.
History
Kinetics emerged from the nineteenth-century rate studies of Wilhelmy, Guldberg, and Waage; Arrhenius proposed his activation-energy equation in 1889, and the 1930s brought the absolute rate theory of Eyring and of Evans and Polanyi, recasting kinetics in terms of statistical mechanics and potential energy surfaces.
Key figures
- Svante Arrhenius
- Henry Eyring
- Cyril Norman Hinshelwood
Related topics
Seminal works
- eyring1935
- atkins2018
- laidler1987
Frequently asked questions
- Why doesn't a balanced chemical equation tell you the rate law?
- Stoichiometry describes the overall conversion but not the molecular pathway; the rate law reflects the slowest elementary step and the species involved up to it, so it must be determined experimentally rather than read from the overall equation.
- If a reaction is thermodynamically favourable, why might it still not occur?
- Thermodynamics sets only whether a reaction can happen, not how fast; a large activation barrier can make a favourable reaction immeasurably slow, which is why diamond persists despite graphite being the more stable form of carbon.