A rate law connects how fast a chemical reaction proceeds to the concentrations of its reactants. Unlike the balanced chemical equation, which is fixed by mass conservation, the rate law is an experimental result — it has to be measured, not derived on paper. This article covers what the rate law tells you, how chemists determine reaction order in the lab, why the rate constant's units shift with order, and where the simple power-law model breaks down.
What the rate law tells you
The rate law rate = k[A]^m[B]^n packages three separate pieces of information: the rate constant k (how fast the reaction runs, all else equal, at a given temperature), and the orders m and n (how sensitively the rate responds to each reactant's concentration). A reaction that is second order in a reactant is far more sensitive to that reactant running low than a zero-order reaction, which barely reacts to concentration changes at all. Knowing the rate law lets chemists predict how a reaction slows as it proceeds, design a process to run at a target rate, or diagnose which step in a multi-step mechanism is rate-limiting.
Determining order experimentally (the method of initial rates)
Because the rate law can't be read off the balanced equation for most real reactions, chemists measure the initial rate — the rate right at the start, before products build up and complicate things — across a handful of experiments that vary one or more starting concentrations. Comparing any two experiments where only one concentration changed isolates that reactant's order: if doubling [A] doubles the rate, the order in A is 1; if it quadruples the rate, the order is 2; if the rate doesn't change at all, the order is 0. The Reaction Order tab generalizes this by solving the full log-linear system across three experiments simultaneously, so the concentrations don't have to be held perfectly constant between trials — any three experiments that vary independently enough will work.
Why k's units change with order
Reaction rate is always reported in concentration per time (typically M/s, i.e. mol·L⁻¹·s⁻¹), regardless of the rate law's form. Since rate = k × [concentration]ⁿ for an overall order-n reaction, the concentration terms on the right must combine with k's own units to leave just M/s. That forces k's units to be M^(1−n)·s⁻¹ — s⁻¹ for a first-order reaction (n=1, since M⁰ = 1), M⁻¹·s⁻¹ for second order, and so on. This is also a useful sanity check: if a computed or looked-up rate constant's units don't match its stated order, something is wrong.
Elementary vs. overall (complex) reactions
For an elementary reaction — one that happens in a single molecular collision, exactly as written — the rate law's exponents do equal the stoichiometric coefficients, because the rate genuinely depends on how often the required molecules collide together. But most reactions people study, including nearly everything in an intro chemistry course past the simplest examples, proceed through multiple elementary steps whose overall rate law reflects the slowest (rate-determining) step plus any fast pre-equilibria. That's precisely why order must be measured rather than assumed: a reaction 2A + B → products could be first order in A, second order in A, or anything else, depending on its actual mechanism.
Common pitfalls
The single most common mistake is assuming reaction order equals the stoichiometric coefficient from the balanced equation — that shortcut only holds for genuinely elementary steps. Another is mixing up initial rate (used to determine the rate law) with instantaneous rate at some later time, which keeps changing as reactants are consumed. Finally, watch units carefully when reporting k: a rate constant reported without its (order-dependent) units is not fully specified, and comparing rate constants across different-order reactions numerically is not meaningful unless the units are also converted.