How fast reactions go and what speeds them up
Chemical kinetics measures how quickly reactants turn into products. The rate of reaction is the change in concentration per unit time. The rate law, rate = k[A]ˣ[B]ʸ, is found by experiment, and the order of reaction is x + y, which can be zero, whole or fractional. Molecularity, by contrast, is the number of molecules in an elementary step and is always a whole number.
For a zero-order reaction, [A] = [A]₀ − kt and the half-life depends on initial concentration. For a first-order reaction, k = (2.303/t) log([A]₀/[A]) and the half-life is constant, t½ = 0.693/k, which is why radioactive decay has a fixed half-life. Rate increases with temperature because more molecules have enough energy to cross the activation energy barrier. The Arrhenius equation, k = Ae^(−Ea/RT), describes this, and for many reactions the rate roughly doubles for every 10 °C rise. Catalysts provide a lower-energy pathway without being consumed and do not change the equilibrium position.
Why it matters
Kinetics explains why food lasts longer in a refrigerator, how medicines are dosed by their half-life in the body, and why catalytic converters cut vehicle pollution. Industries choose temperatures and catalysts to make products quickly and cheaply. Carbon dating uses first-order decay to estimate the age of fossils.
Key terms
- Rate constant
- The proportionality constant k in a rate law, which depends on temperature.
- Order of reaction
- The sum of the powers of concentration terms in the experimentally determined rate law.
- Half-life
- The time taken for the concentration of a reactant to fall to half its initial value.
- Activation energy
- The minimum energy reacting molecules need for a reaction to occur.