What it covers: How fast — reaction rates, rate laws, collision theory, energy profiles, reaction mechanisms, and catalysis.
Exam weight: About 7–9% of the AP Chemistry exam.
The big question: What determines how fast a reaction proceeds, and how can a multistep mechanism explain the overall observed rate law?
Big Ideas covered: Chemical Effects (CE), Transformations (TRA).
Key topics at a glance
Reaction Rates & Rate Law
Rate = change in [concentration] / time. Rate law orders are found experimentally, not from the balanced equation.
Concentration vs. Time
Integrated rate laws give [A] as a function of time. Half-life is constant only for first-order reactions.
Elementary Reactions
Single collision events. Rate law CAN be written directly from coefficients — the one exception to "rate laws are experimental."
Collision Model
Reactions need sufficient energy + correct orientation. Activation energy (Ea) is the minimum energy threshold.
Reaction Energy Profile
Potential energy vs. progress. Peak = transition state. Height of peak above reactants = activation energy.
Reaction Mechanisms
Sum of elementary steps = overall reaction. Intermediates cancel out; they don't appear in the overall equation.
Rate-Determining Step
The slowest step controls the overall rate — it's the bottleneck of the whole mechanism.
Catalysis
Catalysts lower activation energy via a new pathway. They are NOT consumed and don't appear in the overall equation.
The key terms you must know
Rate law / rate constant (k) / reaction order — experimentally determined relationship between rate and concentration.
Integrated rate law / half-life — concentration as a function of time; half-life is constant only for first-order reactions.
Elementary reaction / molecularity — single-step collision events where the rate law matches the coefficients.
Activation energy (Ea) — minimum energy needed for a successful reaction-producing collision.
Transition state — the highest-energy point along the reaction pathway.
Reaction mechanism / intermediate — the sequence of elementary steps; intermediates are made and consumed, never in the final equation.
Rate-determining step — the slowest step; it determines the overall observed rate law.
Catalyst — speeds up a reaction by lowering Ea via a new pathway, without being consumed.
Key themes to remember
Rate laws are experimental, except for elementary steps. Never assume the order matches the balanced overall equation's coefficients.
The slowest step is the bottleneck. No matter how fast other steps are, the rate-determining step caps the overall reaction speed.
Higher temperature speeds up reactions because more molecules have enough energy to overcome Ea — not because of more frequent collisions alone.
Catalysts change the pathway, not the products. They speed up both forward and reverse reactions equally and don't shift equilibrium position (more in Unit 7).
Intermediates and the rate-determining step are connected. If the mechanism's first step is slow, the rate law uses only reactants from that step. If a later step is slow, the steady-state approximation handles intermediate concentrations.
Common exam traps
Don't use stoichiometric coefficients to predict reaction order. Only the rate-determining elementary step's molecularity gives you the rate law directly.
A catalyst lowers Ea — it does not change the overall energy change (ΔE) of the reaction. The reactants and products' energies stay the same; only the pathway between them changes.
An intermediate is not the same as a catalyst. A catalyst is present at the start and regenerated at the end; an intermediate is created partway through and consumed before the end.
Half-life is constant only for first-order reactions. Zero-order and second-order half-lives depend on the starting concentration.
Doubling a reactant's concentration doesn't always double the rate. It depends on that reactant's order — check the exponent in the rate law before assuming a direct proportional relationship.