Reaction Kinetics and Reaction Mechanisms
Rate laws
Rate = k[A]^m[B]^n, where m and n are experimental orders, not the coefficients. Orders are found from initial-rate data: if doubling [A] doubles the rate, the order in A is first; if it quadruples, second; if nothing happens, zero. Units of k follow from the overall order: zero order mol L-1 s-1, first s-1, second L mol-1 s-1. For a first-order reaction, half-life = 0.693/k, independent of concentration.
Collision theory and Arrhenius
Molecules react only if they collide with sufficient energy (at least the activation energy Ea) and correct orientation. Raising temperature increases both collision frequency and the fraction of molecules above Ea, so the rate constant rises steeply - the Arrhenius equation k = Ae^(-Ea/RT). A catalyst provides an alternative pathway with lower Ea; it speeds the forward and reverse reactions equally and does not change ΔH.
Mechanisms
A mechanism is a sequence of elementary steps; the slowest step is rate-determining. Intermediates are formed and consumed during the reaction; the transition state is the energy maximum of a single step. The rate law of an elementary step follows its stoichiometry, which is why a proposed mechanism must be consistent with the measured order.
Worked example
From initial rates: exp 1 gives [A] = 0.10, [B] = 0.10, rate 2.0x10^-3; exp 2 gives [A] = 0.20, [B] = 0.10, rate 8.0x10^-3; exp 3 gives [A] = 0.10, [B] = 0.20, rate 4.0x10^-3.
Doubling A multiplies the rate by four, so the order in A is second. Doubling B doubles the rate, so the order in B is first. k = rate/([A]^2[B]) = 2.0x10^-3/(0.010 x 0.10) = 2.0 L^2 mol-2 s-1.
Traps
Reading orders from the balanced equation. Confusing an intermediate with a catalyst. Thinking a catalyst increases the yield. Using concentration instead of initial rate when determining order.