Reaction Order & Rate Law Calculator

Reaction Kinetics Guide

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 ste

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.

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.

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.