FirstOrderKinetik

First-order kinetics refers to a class of processes in which the rate of change is proportional to the amount of a single species present. In chemistry and related fields, the rate law is typically written as d[A]/dt = -k[A], where [A] is the concentration of the reacting species and k is the first-order rate constant with units of time−1.

The integrated form of the first-order rate law is [A](t) = [A]0 e^(−kt), or equivalently ln([A](t)/[A]0) = −kt. A hallmark of first-order processes is that the half-life, t1/2 = ln 2 / k, is constant and independent of the initial concentration. On a semi-log plot of [A] versus time, the data appear as a straight line with slope −k.

Temperature and other conditions influence k through the Arrhenius relationship, k = A e^(−Ea/RT), so higher temperatures generally increase the rate. First-order kinetics is common in radioactive decay, many chemical decompositions, and in pharmacokinetics for the elimination of drugs under linear (non-saturating) conditions.

In practice, a reaction may exhibit pseudo-first-order kinetics when one reactant remains in large excess, making the overall rate effectively first order with respect to the limiting reactant. First-order behavior is often contrasted with zero-order kinetics, where the rate is independent of concentration (d[A]/dt = −k) and the concentration decreases linearly with time, and with second-order kinetics, where the rate depends on the square of the concentration or on the product of two concentrations.

Overall, first-order kinetics provides a simple, widely applicable framework for describing exponential decay processes and concentration changes over time.