Chemical Kinetics Cheat Sheet

The core ideas of Chemical Kinetics distilled into a single, scannable reference — perfect for review or quick lookup.

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Quick Reference

Change in concentration per unit time: Rate = -d[A]/dt.

Rate = k[A]^m[B]^n where exponents are experimental orders.

Minimum energy (Ea) for reaction; barrier to transition state.

k=Ae^(-Ea/RT) relates rate constant to T and Ea.

Sequence of elementary steps; must match observed rate law.

Lowers Ea via alternative pathway; not consumed; no equilibrium change.

Time for [reactant] to halve. First-order: t1/2=0.693/k.

Reactions need collisions with energy >= Ea and proper orientation.

Activation Energy (Ea):Minimum energy for a successful collision to form products.

Arrhenius Equation:k = Ae^(-Ea/RT); relates rate constant to temperature and Ea.

Catalyst:Increases rate by lowering Ea without being consumed.

Collision Theory:Reactions need collisions with sufficient energy and proper orientation.

Elementary Step:A single molecular-level event in a reaction mechanism.

Half-Life:Time for reactant concentration to halve.

Integrated Rate Law:Equation relating concentration to time for a given order.

Intermediate:Species produced in one step and consumed in a later step.

Method of Initial Rates:Technique using initial rate data to determine reaction orders.

Rate Constant (k):Proportionality constant in a rate law; depends on T and Ea.

Rate-Determining Step:Slowest elementary step; limits overall reaction rate.

Rate Law:Rate as a function of concentrations: Rate = k[A]^m[B]^n.

Reaction Mechanism:Sequence of elementary steps converting reactants to products.

Reaction Order:Exponent on concentration in rate law; determined experimentally.

Transition State:Highest-energy configuration along reaction coordinate; activated complex.

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