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⏱Unit 5 · Kinetics Unit Hub Flashcards Cheat Sheet Essentials Visual Review MC Practice SAQ Practice

AP Chemistry Unit 5 Visual Review

A topic-by-topic visual walkthrough of Kinetics — reaction rates, rate laws, reaction mechanisms, the collision model, and catalysis.

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TOPIC 5.1 Reaction Rates product ↑ reactant ↓ concentration time → Rate = how fast concentration changes Rate = change in concentration ÷ change in time (M/s). Reactants DISAPPEAR (negative) and products APPEAR (positive). Rate is always reported as a positive number. Rate slows over time As reactants are used up, their concentration drops, so collisions become less frequent and the rate decreases — the curve flattens out. Factors that increase reaction rate Higher concentration · higher temperature · larger surface area · presence of a catalyst. All work by increasing the frequency or effectiveness of collisions between particles (Topic 5.5). ⏱Rate measures how fast concentrations change — it slows as reactants run out. The Review Hub · AP Chemistry Unit 5 TOPIC 5.2 Introduction to Rate Law Rate = k [A]ᵐ [B]ⁿ What each part means k = the rate constant (depends on temperature) [A], [B] = reactant concentrations m, n = the reaction orders (found by EXPERIMENT) Orders are NOT the balanced-equation coefficients! Reaction order tells the effect Zero order in A: changing [A] has NO effect on rate First order in A: double [A] → rate DOUBLES (×2) Second order in A: double [A] → rate QUADRUPLES (×4) Overall order = m + n (the sum of all orders). Finding order from experimental data (method of initial rates) Compare two trials where only ONE reactant's concentration changes. See how the rate responds: If doubling [A] doubles the rate → first order in A. If doubling [A] gives 4× the rate → second order in A. Once orders are known, plug one trial's values back in to solve for k (with its units, which depend on overall order). The rate law can only be determined experimentally — never just read off the balanced equation. Rate = k[A]ᵐ[B]ⁿ — the orders come from experiment, not the coefficients. The Review Hub · AP Chemistry Unit 5 TOPIC 5.3 Concentration Changes Over Time Integrated rate laws tell you [reactant] at any TIME — and each gives a straight-line plot Zero order Linear plot: [A] vs. time [A]ₜ = [A]₀ − kt slope = −k [A] drops at a constant rate — a straight line down. First order Linear plot: ln[A] vs. time ln[A]ₜ = ln[A]₀ − kt slope = −k Constant half-life; t½ = 0.693 ÷ k. Second order Linear plot: 1/[A] vs. time 1/[A]ₜ = 1/[A]₀ + kt slope = +k Half-life keeps getting longer as [A] falls. Determining order from a graph The order is whichever plot gives a STRAIGHT LINE: • [A] vs. t linear → zero order • ln[A] vs. t linear → first order • 1/[A] vs. t linear → second order Half-life (t½) is the time for [A] to fall to half. For FIRST order it's constant (t½ = 0.693/k), independent of concentration — the hallmark of radioactive decay and many first-order reactions. The linear plot reveals the order: [A], ln[A], or 1/[A] vs. time (zero, first, second). The Review Hub · AP Chemistry Unit 5 TOPIC 5.4 Elementary Reactions An elementary step happens in a single collision Most reactions occur as a SERIES of simple steps. Each single step (one collision or one molecule breaking apart) is an elementary reaction. Together, the steps make up the reaction mechanism (Topic 5.7). Key rule: for an ELEMENTARY step ONLY, the rate law comes directly from the coefficients. Molecularity The number of particles colliding in an elementary step: Unimolecular — one particle (A → products) Bimolecular — two particles collide (A + B → products) Termolecular (3 particles) is rare — 3-way collisions are unlikely. Rate law from an elementary step Only for a single elementary step, the order = the number of each reactant in that step. A + B → C has rate = k[A][B] This is why the slow step's molecularity determines the rate law. Remember: this "coefficients = order" rule ONLY applies to elementary steps — never to the overall reaction. For an elementary step only, the rate law comes straight from its coefficients (molecularity). The Review Hub · AP Chemistry Unit 5 TOPIC 5.5 The Collision Model Reactions require effective collisions Particles must COLLIDE to react — but not every collision works. An "effective" (successful) collision must have: 1 · enough energy (≥ the activation energy, Eₐ) AND 2 · the correct orientation Most collisions fail one of these two requirements and bounce apart unchanged. Activation energy (Eₐ) The minimum energy needed to reach the transition state and break existing bonds — the "energy hill" the reaction must climb. Lower Eₐ = faster reaction (more collisions succeed). Why temperature speeds reactions up Higher temperature → particles move faster and collide more often, AND a larger FRACTION of collisions now have enough energy to exceed Eₐ. The energy effect is the bigger one — small T rises can double rate. Concentration & surface area raise collision FREQUENCY; temperature raises both frequency AND collision energy. A reaction needs collisions with enough energy (≥ Eₐ) and correct orientation. The Review Hub · AP Chemistry Unit 5 TOPIC 5.6 Reaction Energy Profile Eₐ (activation energy) ΔH (exothermic here) reactants products transition state (peak) energy reaction progress → Reading the profile Peak = transition state (highest energy) Eₐ = reactants → peak height ΔH = products − reactants energy Eₐ affects rate; ΔH affects energy released. Exo vs. endothermic Products LOWER than reactants → exothermic (releases energy, −ΔH). Products HIGHER → endothermic (+ΔH). ΔH is independent of the path. A reverse reaction has the same transition state; its activation energy = Eₐ(forward) − ΔH. The energy profile shows Eₐ (the hill) and ΔH (net energy change) — Eₐ sets the rate. The Review Hub · AP Chemistry Unit 5 TOPIC 5.7 Introduction to Reaction Mechanisms A mechanism is the step-by-step path A reaction mechanism is the sequence of elementary steps that add up to the overall reaction. Adding all the steps together must give the overall balanced equation. A valid mechanism must (1) sum to the overall reaction and (2) be consistent with the experimental rate law. Intermediates A species that is PRODUCED in one step and CONSUMED in a later step. It appears in the middle but NOT in the overall equation. Contrast: a catalyst is consumed first, then regenerated. The rate-determining step The SLOWEST step is the bottleneck — it limits the overall reaction rate, just like the slowest worker on an assembly line. The overall rate law is based on this slow step (Topic 5.8). A mechanism is a proposed hypothesis — it can never be proven, only supported (or ruled out) by the rate law. A mechanism is a series of elementary steps; the slowest step controls the overall rate. The Review Hub · AP Chemistry Unit 5 TOPIC 5.8 Reaction Mechanism & Rate Law The slow step sets the rate law If the FIRST step is the slow (rate-determining) step, the rate law is written directly from its reactants — using its coefficients as the orders (because a single step's rate law follows its molecularity, Topic 5.4). A proposed mechanism is only valid if its predicted rate law MATCHES the experimental rate law. Worked example Step 1 (slow): NO₂ + NO₂ → NO₃ + NO Step 2 (fast): NO₃ + CO → NO₂ + CO₂ The slow step is bimolecular in NO₂, so the rate law is rate = k[NO₂]². Notice CO does NOT appear in the rate law — even though it's in the overall reaction — because it's only in the fast step. NO₃ is an intermediate (made in step 1, used in step 2). If a later step is slow, an intermediate may appear — then use the pre-equilibrium approximation (Topic 5.9). The rate-determining (slow) step gives the rate law — species only in fast steps don't appear. The Review Hub · AP Chemistry Unit 5 TOPIC 5.9 Pre-Equilibrium Approximation When the fast step comes first Sometimes a fast, reversible step happens BEFORE the slow step. The slow step's rate law then contains an INTERMEDIATE — but rate laws can't contain intermediates, so we must substitute it out. The pre-equilibrium approximation lets us rewrite the intermediate in terms of the actual reactants. How it works 1. The fast first step reaches equilibrium quickly, so its forward rate = its reverse rate. 2. Set those two rates equal and solve for the intermediate's concentration in terms of the reactants. 3. Substitute that expression into the slow step's rate law. The result is a rate law written only in terms of real reactants (and a combined constant) — matching experiment. Key idea: replace any intermediate in a rate law using the equilibrium of the fast step before it. A fast pre-equilibrium lets you replace an intermediate in the rate law with real reactants. The Review Hub · AP Chemistry Unit 5 TOPIC 5.10 Multistep Reaction Energy Profile TS 1 (higher) TS 2 intermediate (valley) reactants products energy reaction progress → One hump per step Each elementary step has its own peak (transition state) and Eₐ. Between peaks, a VALLEY is the intermediate — a real species. The highest peak = slow step The step with the LARGEST Eₐ (the tallest hump) is rate-determining. Here that's step 1 (TS 1 is highest). Overall ΔH still = products − reactants. Count the peaks to count the steps; count the valleys to count the intermediates. Each step is a peak; a valley is an intermediate; the tallest peak is the slow step. The Review Hub · AP Chemistry Unit 5 TOPIC 5.11 Catalysis uncatalyzed (high Eₐ) catalyzed (low Eₐ) energy reaction progress → A catalyst lowers activation energy It provides an ALTERNATIVE pathway with a lower Eₐ, so more collisions succeed and the reaction speeds up — both forward AND reverse rates increase equally. What a catalyst does NOT change • It is NOT consumed (regenerated by the end) • It does NOT change ΔH or the equilibrium position • It only helps you REACH equilibrium faster Types of catalysts Homogeneous — same phase as reactants (e.g., an aqueous catalyst in solution). Heterogeneous — different phase (e.g., a solid surface where gases react — like a catalytic converter). Enzymes are biological catalysts. A catalyst lowers Eₐ via a new pathway — speeding the reaction without being used up or changing ΔH. The Review Hub · AP Chemistry Unit 5
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This is great for review the night before the exam — fast, visual, and covers everything you need to remember about Unit 5's kinetics content.