The rate of a chemical reaction measures how rapidly concentrations of reactants decrease or products increase per unit time, always expressed as a positive value. The rate law, $\text{rate} = k[A]^m[B]^n$, is determined experimentally and the exponents are not necessarily equal to stoichiometric coefficients. For a zero-order reaction, concentration decreases linearly with time and the half-life is directly proportional to the initial concentration ($t_{1/2} = [A]_0/2k$). For a first-order reaction, the half-life $t_{1/2} = 0.693/k$ is independent of concentration, making it the most NEET-tested formula in this chapter. In a second-order reaction, $1/[A]$ increases linearly with time and the half-life is inversely proportional to the initial concentration. Order is an experimentally determined quantity that can be zero, fractional, or integer, while molecularity is a theoretical count of molecules in an elementary step and is always a positive integer. The Arrhenius equation $k = Ae^{-E_a/RT}$ quantifies how the rate constant increases with temperature as more molecules acquire sufficient energy to surpass the activation energy barrier. A catalyst provides an alternative pathway of lower activation energy, increasing $k$ without altering the thermodynamic quantities $\Delta H$ or $K$. The relationship $E_a(\text{forward}) - E_a(\text{backward}) = \Delta H$ connects kinetics to thermodynamics through the energy profile. Pseudo first-order reactions occur when one reactant is in such large excess that its concentration is effectively constant, simplifying the kinetics to apparent first-order behaviour.