Unit 5: Kinetics
AP Chemistry: 86 practice questions with detailed explanations.
Unit Study Guide
Executive Summary
Kinetics asks how fast reactions go and which pathway they take. Rate laws come from experiment, not from balanced coefficients.
Reaction rates
Rate = Δconcentration/Δtime. Rates decrease as reactants are consumed; measure initial rates to avoid product complications. Rates are always positive.
Rate laws
For aA + bB → products, rate = k[A]ᵐ[B]ⁿ, where m and n (orders) come from experiment. Compare two trials where one reactant doubles while others stay constant to extract each order. Units of k depend on overall order: M^(1−order)·s⁻¹.
Integrated rate laws
Zero order: [A] = [A]₀ − kt (rate independent of concentration; plot [A] vs. t linear). First order: ln[A] = ln[A]₀ − kt (plot ln[A] vs. t linear; half-life t½ = 0.693/k, independent of [A]₀). Second order: 1/[A] = 1/[A]₀ + kt (plot 1/[A] vs. t linear; t½ = 1/(k[A]₀)).
Elementary reactions and mechanisms
An elementary step's rate law comes directly from its molecularity: unimolecular → first order, bimolecular → second order. A mechanism's slowest (rate-determining) step controls the observed rate law; intermediates appear in mechanisms but not in the overall equation or the final rate law.
Collision model and energy profiles
Reactants must collide with enough energy (≥ activation energy Ea) and proper orientation. The Arrhenius equation relates k to Ea and temperature: higher T or lower Ea (catalyst) raises k. An energy profile plots potential energy along the reaction coordinate; Ea is the hill height from reactants to the transition state.
Catalysis
A catalyst provides a lower-Ea pathway, speeds both directions equally, and is not consumed. It does not change ΔH, K, or equilibrium position — only the rate of approach.
Quantitative skill-set
Method of initial rates; integrated-rate-law line fitting; half-life calculations; interpreting multistep energy profiles (intermediates in valleys, transition states at peaks).
Exam traps
Orders never come from stoichiometric coefficients (unless the step is elementary). Doubling a zero-order reactant does not change the rate. Catalysts lower Ea, not ΔH. A fast equilibrium step before the slow step can substitute intermediates out of the rate law.