Chapter 22: Reaction Kinetics

Reaction Kinetics thoroughly investigates the quantitative principles of reaction kinetics, focusing on how to define, determine, and utilize the rate equation, which expresses the reaction rate as directly proportional to the rate constant (k) and the concentrations of reactants raised to their respective orders (e.g., rate = k [A] m [B] n ). This essential equation must be determined exclusively through experimental data and cannot be deduced from the stoichiometric coefficients of the balanced chemical reaction. Core terminology includes the order of reaction (the exponent of a concentration term, typically 0, 1, or 2, but sometimes fractional), the overall order of reaction (the sum of individual orders), and the rate constant (k), whose units depend specifically on the reaction's overall order (e.g., mol dm-3 s-1 for rate units). The reaction order can be determined experimentally by analyzing graphs of reaction rate against concentration (where shape reveals the order), concentration against time graphs, or by calculating successive half-lives. The half-life (t 1/2​ ) is defined as the time taken for the concentration of a limiting reactant to decrease by half, and critically, for a first-order reaction, the half-life is constant and independent of the initial concentration, enabling the calculation of k using the relationship k=0.693/t 1/2​ . Kinetic data is instrumental in confirming proposed reaction mechanisms, particularly identifying the rate-determining step (RDS), which is the slowest elementary step and governs the overall rate; only reactants involved in or formed before the RDS appear in the experimental rate equation. Furthermore, the text addresses the effect of temperature, explaining that increasing temperature significantly raises the rate constant and reaction rate because a greater fraction of molecules possesses energy equal to or greater than the activation energy, as modeled by the Boltzmann distribution. Finally, the chapter details catalysis, noting that catalysts accelerate reactions by providing an alternative pathway with a lower activation energy. Catalysis is classified as either homogeneous (catalyst and reactants in the same phase, often involving redox cycles and changes in oxidation number, such as iron ions or atmospheric nitrogen oxides) or heterogeneous (catalyst in a different phase, like iron in the Haber process or platinum metals in catalytic converters). The mechanism of heterogeneous catalysis involves the sequential processes of adsorption of reactants, weakening of bonds and surface reaction, and subsequent desorption of the products.