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Chemical kinetics studies how fast reactions proceed and what factors control reaction rates. The rate law — rate = k[A]ⁿ[B]ᵐ — expresses rate as a function of reactant concentrations with experimentally determined orders (n, m) and a temperature-dependent rate constant k. Reaction order must be determined from experimental data, not inferred from stoichiometric coefficients (except for elementary steps). The half-life of a first-order reaction, t₁/₂ = ln(2)/k, is constant and independent of initial concentration — a key diagnostic.
Determine rate laws from initial rate experiments by comparing pairs of experiments where one concentration is varied while others are held constant. Practice integrated rate laws by plotting concentration data three ways (ln[A] vs. t, 1/[A] vs. t, [A] vs. t) and identifying which is linear to determine reaction order.
Chemical equilibrium tells you where a reaction ends up; chemical kinetics tells you how fast it gets there. These are completely independent. A reaction can have a large negative ΔG — strongly thermodynamically favored — yet proceed so slowly at room temperature that it is practically inert. Diamond converting to graphite is the paradigmatic example: thermodynamically spontaneous, but it essentially never happens because the activation energy barrier is enormous. Kinetics and thermodynamics must both be understood to predict what will actually happen in a chemical system.
The rate law, rate = k[A]ⁿ[B]ᵐ, summarizes how reaction rate depends on concentrations. The exponents n and m are the reaction orders with respect to each reactant; the sum n + m is the overall order. These exponents must be determined experimentally by measuring how rate changes when concentrations are varied one at a time. They cannot be read from the balanced equation — stoichiometric coefficients and rate law exponents are unrelated for multistep mechanisms. The rate constant k encodes temperature dependence and has units that depend on the overall order.
The integrated rate laws link the rate law to observable concentration changes over time. For a first-order reaction, [A] decays exponentially: [A] = [A]₀ e−kt, and a plot of ln[A] versus t is linear with slope −k. For a second-order reaction, 1/[A] versus t is linear. For zero-order, [A] versus t is linear. This graphical diagnostic — testing all three plots to see which one is linear — is how reaction order is determined from concentration-time data in practice.
The half-life is a useful shorthand for how quickly a reaction proceeds. For first-order reactions, t₁/₂ = ln(2)/k is constant regardless of the starting concentration — each successive half-life removes exactly half of what remains. This constant half-life is the diagnostic signature of first-order kinetics and underlies all of radioactive decay. For second-order reactions, t₁/₂ = 1/(k[A]₀) depends on the initial concentration, so half-lives lengthen as the reaction proceeds.
Temperature has a profound effect on rate through its influence on k. The Arrhenius equation — the next topic — quantifies this: k = A e−Ea/RT. A catalyst works by providing an alternative reaction pathway with lower activation energy Ea, which exponentially increases k at a given temperature. Crucially, the catalyst lowers the energy barrier for both the forward and reverse reactions equally, so the equilibrium constant (and ΔG) is unchanged — the reaction reaches the same equilibrium, just faster.