Chemical Reaction Rates
Reaction rate measures how fast reactants convert to products, expressed as change in concentration over time: rate = -?[reactant]/?t or ?[product]/?t. Units are typically M/s (mol/(L?s)). Rates depend on concentration, temperature, catalysts, and surface area. The rate law expresses rate as a function of concentration: rate = k[A]?[B]?, where k is the rate constant, and m and n are reaction orders determined experimentally, not from balanced equations.
Reaction Orders and Rate Constants
Zero-order reactions have rates independent of concentration: rate = k. First-order reactions: rate = k[A], giving exponential decay with constant half-life (t?/? = 0.693/k). Second-order reactions: rate = k[A]^2 or k[A][B], with half-life depending on initial concentration. The rate constant k is temperature-dependent, increasing exponentially with temperature (Arrhenius equation). Catalysts increase k by lowering activation energy, speeding reactions without being consumed.
Applications in Science and Industry
Understanding reaction rates is crucial in chemistry, biology, and engineering. Pharmaceutical scientists optimize drug synthesis rates. Chemical engineers design reactors with appropriate residence times. Environmental scientists model pollutant degradation rates. Biochemists study enzyme kinetics (Michaelis-Menten equations). Food scientists slow spoilage reactions by refrigeration or preservatives. Explosives specialists deal with very fast reactions. Polymer chemists control polymerization rates. Reaction kinetics informs everything from cooking times to industrial chemical production.
Quick Tips
- Always verify units are consistent
- Use scientific notation for very large/small numbers
- Results are approximations — real conditions may vary
Frequently Asked Questions
Rate changes during a reaction as concentrations change. Rate constant (k) is independent of concentration but depends on temperature and catalysts. Rate = k x (concentration terms) relates them.
Higher temperature means more molecular kinetic energy. More molecules have sufficient energy to overcome activation energy barriers, increasing collision frequency and success rate. The Arrhenius equation quantifies this relationship.
Half-life is the time required for reactant concentration to decrease by half. For first-order reactions, half-life is constant (t?/? = 0.693/k). For other orders, half-life depends on initial concentration.
Catalysts lower activation energy, allowing more molecules to react at a given temperature. This increases the rate constant k, speeding the reaction. Catalysts aren't consumed and don't affect equilibrium position, only how fast equilibrium is reached.
No, reaction order must be determined experimentally. The balanced equation shows stoichiometry but not mechanism. Elementary steps have rate laws matching stoichiometry, but overall reactions often don't because they proceed through multiple steps.