Method of Initial Rates

Chemists measure only the very first instant of each trial because that's the one moment where concentrations are known exactly — they're whatever you poured in. A few seconds later, products have formed, reactants have been consumed, and the reverse reaction may even be creeping in. The initial rate is a clean snapshot, uncontaminated by any of that drift.

To expose one reactant's order, run trials where you change its concentration while holding every other concentration fixed. If doubling [A] doubles the rate, A is first order: rate depends on [A]^1. If doubling [A] quadruples the rate, A is second order: [A]^2. If doubling [A] does nothing, A is zero order and doesn't appear in the rate law at all. Repeat the trick for each reactant, holding the others constant, and you build the full rate law piece by piece: rate = k[A]^m[B]^n. Once you know the orders, plug any single trial's concentrations and measured rate back into that equation and solve for k, the rate constant, which carries its own units depending on the overall order.

This matters because a rate law cannot be guessed from the balanced equation's coefficients — those describe stoichiometry, not mechanism. Rate depends on the slow, rate-determining step of the mechanism, something invisible in the overall equation. That's why this topic sits right before mechanisms: initial rates are the experimental evidence that lets you deduce which step is slow and what the mechanism must look like.