Reaction rate

The reaction rate is defined as the change in concentration of a reactant or product per unit time, typically expressed in moles per liter per second (mol L⁻¹ s⁻¹). For a reaction aA + bB → cC + dD, the rate can be written as rate = −(1/a) d[A]/dt = (1/c) d[C]/dt, ensuring a positive value regardless of whether a reactant or product is monitored.¹ Rates are measured experimentally by techniques such as spectrophotometry, conductometry, or manometry, which track concentration changes over time.

Several factors influence reaction rate: concentration of reactants (higher concentration increases collision frequency), temperature (higher temperature increases kinetic energy and the fraction of molecules exceeding activation energy), pressure (for gases, increased pressure raises concentration), surface area (for solids, greater area exposes more particles), and the presence of a catalyst (which lowers activation energy without being consumed). The Arrhenius equation quantitatively relates rate constant to temperature: k = A e^(−Ea/RT), where Ea is activation energy and A is the pre-exponential factor.²

The rate law expresses the rate as a function of reactant concentrations: rate = k[A]^m[B]^n, where m and n are reaction orders determined experimentally, not from stoichiometry. The overall order is the sum m+n. For elementary reactions, orders match molecularity, but for complex reactions they reflect the rate-determining step.¹ Integrated rate laws allow determination of concentration over time for zero-, first-, and second-order reactions, and half-life calculations are particularly useful for first-order kinetics.

Reaction rates are critical in industrial chemistry for optimizing yield and selectivity, such as in the Haber process for ammonia synthesis, where catalysts and high pressures are used to increase rate. In environmental science, rates govern pollutant degradation and atmospheric ozone depletion. In biology, enzyme kinetics (Michaelis–Menten) describe how reaction rates depend on substrate concentration, essential for drug design and metabolic understanding.³

One subtlety is that reaction rates can be affected by isotopic substitution (kinetic isotope effects), where replacing hydrogen with deuterium slows a reaction due to changes in zero-point energy. Another is that some reactions exhibit negative activation energies, often for barrierless radical recombination, where the rate decreases with temperature.⁴ Additionally, the concept of 'rate' in oscillating reactions, such as the Belousov–Zhabotinsky reaction, shows periodic changes in concentration, challenging simple rate definitions. Historically, the first quantitative rate study was by Ludwig Wilhelmy in 1850 on sucrose inversion, laying groundwork for chemical kinetics.⁵

Activation energy

The minimum energy that colliding molecules must possess for a reaction to occur.

Rate constant

A proportionality constant (k) in the rate law, dependent on temperature and catalyst presence.

Reaction order

The exponent of a concentration term in the rate law, indicating how rate depends on that reactant.

Elementary reaction

A single-step reaction with a molecularity equal to its reaction order.