Chemical Kinetics Study Pack
Kibin's free study pack on Chemical Kinetics includes a 6-section study guide, 25 quiz questions, 30 flashcards, and 5 open-ended Explain review questions. Sign up free to track your progress toward mastery, plus upload your own notes and recordings to create personalized study packs organized by course.
Last updated May 27, 2026
Chemical Kinetics Study Guide
Master reaction rates, rate laws, and the Arrhenius equation as you work through reaction orders, rate-determining steps, and catalysis. This pack covers the core concepts college chemistry students need to understand how and why reactions speed up or slow down.
Key Takeaways
- •Chemical kinetics is the study of reaction rates — how fast reactants are converted to products — and the molecular-level factors that control those rates.
- •The rate of a reaction is expressed as the change in concentration of a reactant or product per unit time, and it varies continuously as concentrations change throughout the reaction.
- •A rate law relates reaction rate mathematically to reactant concentrations raised to experimentally determined exponents called reaction orders; the overall order is the sum of all individual exponents.
- •The rate constant k is temperature-dependent and increases exponentially with temperature according to the Arrhenius equation, which connects k to activation energy and absolute temperature.
- •Catalysts accelerate reactions by providing an alternative pathway with a lower activation energy, without being consumed in the overall process.
- •Reaction mechanisms describe the sequence of elementary steps — each with its own molecularity and rate law — through which reactants are converted to products, and the slowest step, called the rate-determining step, controls the overall rate.
Defining and Measuring Reaction Rates
The foundation of chemical kinetics is a precise, quantitative definition of how fast a chemical reaction proceeds, which requires careful attention to how concentration changes are tracked over time.
Average vs. Instantaneous Reaction Rate
- •The average rate is calculated as the magnitude of the change in molar concentration of a species divided by the elapsed time: rate = |Δ[X]| / Δt.
- •The instantaneous rate at any specific moment is found by taking the derivative of the concentration–time curve at that point, which corresponds graphically to the slope of a tangent line.
- •Because rates are always expressed as positive quantities, the change in reactant concentration is multiplied by −1 to remove the negative sign arising from consumption.
Stoichiometric Correction of Rates
- •When stoichiometric coefficients are not all equal to 1, the rate expression for each species must be divided by that species' coefficient so all species yield the same unique rate for the reaction.
- •For the reaction 2 NO₂(g) → 2 NO(g) + O₂(g), the rate is expressed as −(1/2)(Δ[NO₂]/Δt) = (1/2)(Δ[NO]/Δt) = Δ[O₂]/Δt.
- •This normalization ensures that one unambiguous reaction rate describes the overall process regardless of which species is monitored.
Experimental Methods for Tracking Concentration
- •Spectrophotometry measures absorbance at a wavelength specific to a colored reactant or product, allowing continuous monitoring without disturbing the reaction mixture.
- •Gas-phase reactions can be tracked by monitoring total pressure changes when the number of moles of gas changes during reaction.
- •Titration, pH measurement, and conductivity monitoring are common methods for reactions in aqueous solution.
Rate Laws and Reaction Order
A rate law is a mathematical expression that shows precisely how the rate of a reaction depends on the concentrations of reactants, and its form must be determined by experiment rather than derived from the balanced equation alone.
Structure of a Rate Law
- •For a general reaction involving reactants A and B, the rate law takes the form: rate = k[A]^m[B]^n, where k is the rate constant and m and n are the reaction orders with respect to each reactant.
- •The overall reaction order is m + n, and the units of k depend on the overall order so that the rate always carries units of mol L⁻¹ s⁻¹ (or equivalent).
- •The exponents m and n are typically small integers or zero, but they can also be fractions; they reflect the mechanism of the reaction, not the stoichiometry.
Common Reaction Orders and Their Concentration Dependence
- •A zero-order reaction has a rate that is completely independent of reactant concentration; the integrated rate law is [A]t = −kt + [A]₀, producing a straight line when [A] is plotted against time.
- •A first-order reaction has a rate directly proportional to one reactant's concentration; its integrated form is ln[A]t = −kt + ln[A]₀, linear when ln[A] is plotted against time, and it produces a constant half-life of t₁/₂ = 0.693/k.
- •A second-order reaction has a rate proportional to the square of one reactant's concentration or to the product of two reactants' concentrations; the integrated form 1/[A]t = kt + 1/[A]₀ is linear in a plot of 1/[A] vs. time, and its half-life increases as concentration decreases.
Determining Rate Laws from Initial Rate Data
- •The method of initial rates compares reaction rates measured at the very start of experiments conducted with different initial concentrations while holding all other variables constant.
- •Doubling [A] while holding [B] fixed and observing that the rate doubles indicates first-order behavior in A; a fourfold rate increase indicates second-order behavior in A.
- •The rate constant k is calculated by substituting experimentally determined concentrations and the corresponding measured rate into the rate law expression.
Molecular Collisions and Activation Energy
For a chemical reaction to occur, reactant molecules must physically collide with sufficient energy and with the correct spatial orientation — a framework developed in collision theory and refined by transition state theory.
Collision Theory Requirements
- •Not every collision between reactant molecules produces products; only collisions that meet two simultaneous criteria are effective: the collision must supply energy at least equal to the activation energy, and the molecules must be oriented correctly relative to each other.
- •The activation energy (Ea) is the minimum kinetic energy that colliding particles must possess for the collision to break existing bonds and begin forming new ones.
- •At higher temperatures, a larger fraction of molecules in the Maxwell–Boltzmann distribution have energies exceeding Ea, so the fraction of effective collisions — and therefore the rate — increases dramatically.
Transition State and Reaction Energy Profiles
- •The transition state (or activated complex) is the highest-energy, unstable arrangement of atoms at the peak of the potential energy diagram, positioned between reactants and products.
- •The energy difference between reactants and the transition state is Ea; the energy difference between reactants and products determines whether the overall reaction is exothermic or endothermic.
- •A catalyst lowers the activation energy by stabilizing the transition state, which shifts the peak of the energy profile downward without changing the energies of reactants or products.
The Arrhenius Equation and Temperature Dependence
- •The Arrhenius equation, k = Ae^(−Ea/RT), quantifies how the rate constant k depends on absolute temperature T, the activation energy Ea, and the universal gas constant R.
- •The pre-exponential factor A (the frequency factor) accounts for the collision frequency and the fraction of collisions with correct orientation.
- •A useful two-temperature form of the Arrhenius equation, ln(k₂/k₁) = (Ea/R)(1/T₁ − 1/T₂), allows calculation of Ea when k is known at two temperatures, or prediction of k at a new temperature when Ea is known.
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About this Study Pack
Created by Kibin to help students review key concepts, prepare for exams, and study more effectively. This Study Pack was checked for accuracy and curriculum alignment using authoritative educational sources. See sources below.
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Question 1 of 25
Your progress is saved after each question and counts toward mastery.
For the reaction 2 NO₂(g) → 2 NO(g) + O₂(g), which expression correctly represents the single unique reaction rate?
Card 1 of 30
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Concept 1 of 5
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Rate Laws and Reaction Order
Explain what a rate law is and what reaction order means in your own words. How is a rate law determined, and why can't you simply read it off from the balanced chemical equation?
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