Chemical Kinetics
Chemical kinetics tests how rate depends on concentration and temperature. Expect order determination from data, integrated rate calculations and Arrhenius reasoning.
Rate law and order
Rate = k[A]^m[B]^n; m + n is the order, which is experimental and not read from the balanced equation. Molecularity is the number of species colliding in an elementary step and is a whole number. Units of k follow mol^(1−n) L^(n−1) s−1 for order n.
Integrated rate equations
Zero order: [A] = [A]0 − kt, t1/2 = [A]0/2k. First order: ln([A]0/[A]) = kt, k = (2.303/t) log([A]0/[A]), t1/2 = 0.693/k, independent of initial concentration. A plot of ln[A] against t is linear for first order; a plot of 1/[A] against t is linear for second order.
Arrhenius and collision theory
k = A e^(−Ea/RT), so ln k = ln A − Ea/RT. A plot of ln k against 1/T has slope −Ea/R. Collision theory adds the steric factor P: rate ∝ P Z e^(−Ea/RT). A catalyst lowers Ea and increases k but does not change ΔH or the equilibrium constant.
Worked example
A first-order reaction has k = 2.0 × 10−3 s−1 and [A]0 = 0.10 M. Time to reach 0.025 M: t = (2.303/k) log(0.10/0.025) = (2.303/2.0 × 10−3) × 0.602 = 693 s. Equivalently, 0.10 → 0.050 → 0.025 is two half-lives, each 0.693/2.0 × 10−3 = 347 s.
Mark-losing traps
Confusing order with molecularity, using concentration instead of time correctly in the log term, forgetting that zero-order t1/2 depends on [A]0, and stating that a catalyst shifts equilibrium are frequent errors. Always check units of k.