Reaction Order & Rate Law Calculator
Determine the order of a reaction and calculate the rate constant from experimental concentration-time data or initial rate experiments.
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Reaction Kinetics Guide
What should I know about Orders of Reaction?
The order of a reaction describes how the rate depends on reactant concentrations. Rate = k[A]ⁿ. Zero order (n=0): rate is constant, independent of concentration. First order (n=1): rate is directly proportional to [A] — doubling [A] doubles the rate. Second order (n=2): rate proportional to [A]² — doubling [A] quadruples the rate. The overall order = sum of all individual orders. Order must be determined experimentally — it cannot be deduced from the balanced equation (except for elementary steps, where the order does match the stoichiometric coefficients by definition). For any multi-step mechanism, the order must be found from experimental data — typically by the initial rates method, varying one reactant's concentration at a time and measuring how the initial rate responds.
What do I need to know about Initial Rates Method?
Compare two experiments where only [A] changes. Order = log(rate₂/rate₁) / log([A]₂/[A]₁). Example: [A] doubles (×2) and rate quadruples (×4). log(4)/log(2) = 2 → second order. If rate doubles when [A] doubles: log(2)/log(2) = 1 → first order. Rate constant from rate law: k = rate / [A]ⁿ. Units of k depend on order: zero order k in mol/L/s, first order in s⁻¹, second order in L/mol/s.
What's the key thing to understand about Integrated Rate Laws?
Zero order: [A] = [A]₀ − kt. Plot [A] vs t → straight line. t½ = [A]₀/2k. First order: ln[A] = ln[A]₀ − kt. Plot ln[A] vs t → straight line. t½ = 0.693/k (constant, independent of concentration). Second order: 1/[A] = 1/[A]₀ + kt. Plot 1/[A] vs t → straight line. t½ = 1/(k[A]₀). The linearisation technique: try each plot — the one that gives a straight line reveals the order. First order is most common for elementary reactions.
What's the key thing to understand about Mechanism and Rate-Determining Step?
The rate law reflects the slow (rate-determining) step of the mechanism. If mechanism is A → B (slow), then B → C (fast): rate = k[A] (first order in A). If mechanism is A + A → B (slow), then B → products (fast): rate = k[A]² (second order). Intermediates (species produced then consumed) do not appear in the overall rate law. This connection between mechanism and rate law is how chemists propose and test reaction mechanisms experimentally.