Kinetics — AP Chemistry Unit Overview

For: AP Chemistry candidates sitting AP Chemistry.

Covers: The full Kinetics unit (Unit 5 in the AP Chemistry CED), including all 9 core topics: reaction rate, rate law, integrated rate laws, elementary reactions, collision model, energy profiles, reaction mechanisms, multistep profiles, and catalysis.

You should already know:

• Balancing chemical equations and stoichiometric mole relationships
• Basic enthalpy and energy changes for chemical reactions
• Atomic and molecular bonding and intermolecular force fundamentals

A note on the practice questions: All worked questions in the "Practice Questions" section below are original problems written by us in the AP Chemistry style for educational use. They are not reproductions of past College Board / Cambridge / IB papers and may differ in wording, numerical values, or context. Use them to practise the technique; cross-check with official mark schemes for grading conventions.

1. Why This Unit Matters

Kinetics is the study of the rate (speed) of chemical reactions and the molecular-level processes that produce those reaction rates. Unlike thermodynamics, which answers whether a reaction can occur spontaneously, kinetics answers how quickly that reaction will actually proceed, a critical distinction for both lab and real-world chemistry. For example, the conversion of diamond to graphite is thermodynamically spontaneous, but kinetically so slow that diamonds persist unchanged for millions of years.

This unit makes up 7–9% of the total AP Chemistry exam score, and questions appear in both the multiple-choice (MCQ) and free-response (FRQ) sections. FRQ questions often combine multiple kinetics sub-topics into a multi-part problem that requires connecting macroscopic experimental data to particulate-level explanations. Kinetics is also a foundational prerequisite for understanding equilibrium, acid-base reaction dynamics, and electrochemistry, as all of these topics rely on understanding how reaction rates change with concentration and temperature conditions.

2. Unit Concept Map

The 9 subtopics of Kinetics build sequentially from macroscopic measurement to particulate explanation, ending with applications to modifying reaction rates, with every later topic dependent on mastery of earlier foundational concepts:

1. We start with Reaction rate: the foundational definition, stoichiometric relationships for rates of consumption/formation, and how to measure rate experimentally from concentration data.
2. Next, Introduction to rate law: uses measured initial rate data to relate reaction rate to reactant concentrations, introducing reaction order and the rate constant.
3. Concentration changes over time: extends differential rate law to the integrated form, letting you calculate concentration at any time and determine reaction order graphically.
4. Elementary reactions: shifts from macroscopic rate data to the molecular-level individual reaction steps that make up overall reactions.
5. Collision model: explains why elementary reactions proceed at their observed rate, linking molecular collision frequency, energy, and orientation to reaction rate.
6. Reaction energy profile: visualizes the energy changes along the reaction coordinate, formalizing the concept of activation energy.
7. Introduction to reaction mechanisms: combines multiple elementary steps into an overall reaction, teaching how to identify intermediates, the rate-determining step, and test if a mechanism matches an experimental rate law.
8. Multistep reaction energy profile: extends energy profile concepts to multistep mechanisms, connecting peaks and valleys to activation energy, intermediates, and transition states.
9. Catalysis: applies all prior concepts to explain how catalysts speed up reactions by altering reaction mechanism and lowering activation energy, covering homogeneous and heterogeneous catalysis.

This progression moves from what you measure experimentally to why reactions have the rate they do, ending with how to modify reaction rates for practical use.

3. A Guided Tour of a Typical Exam Problem

We’ll walk through a common multi-part AP-style problem to show how core sub-topics connect in sequence:

Problem context: For the overall reaction $2\text{NO}(g)+{\text{O}}_2(g)\to 2{\text{NO}}_2(g)$, initial rate experimental data is provided, and you are asked to: (1) determine the experimental rate law, (2) identify which of three proposed mechanisms is consistent with the rate law, (3) draw a reaction energy profile for the correct mechanism if the first step is rate-determining.

Step 1: First, you apply skills from Introduction to rate law: you compare how changing the concentration of each reactant changes the initial rate to find the reaction order for each, resulting in the experimental rate law $rate=k[\text{NO}{]}^2[{\text{O}}_2{]}^1$. If you get this step wrong, the rest of your answers will be incorrect, as every subsequent step relies on this experimental foundation.

Step 2: Next, you apply skills from Introduction to reaction mechanisms: you check each proposed mechanism to see if the rate law derived from the rate-determining step (RDS) matches your experimental rate law. For example, a mechanism with a slow first step: $2\text{NO}\rightleftharpoons {\text{N}}_2{\text{O}}_2$ (fast equilibrium), then ${\text{N}}_2{\text{O}}_2+{\text{O}}_2\to 2{\text{NO}}_2$ (slow, RDS) gives a rate law of $rate=k[\text{NO}{]}^2[{\text{O}}_2]$, which matches your experimental result, so you select this mechanism.

Step 3: Finally, you apply skills from Multistep reaction energy profile: you know the second step is rate-determining, so the second activation energy barrier must be the tallest, there is a valley between the two steps for the intermediate ${\text{N}}_2{\text{O}}_2$, and the overall exothermic reaction means products are lower energy than reactants.

This sequence demonstrates how the unit builds: you cannot solve higher-order parts of the problem without mastering the foundational earlier sub-topics.

4. Cross-Cutting Common Pitfalls (and how to avoid them)

• Wrong move: Confusing the stoichiometric coefficient of a reactant in the overall reaction with its reaction order in the rate law. Why: Students learn that elementary steps have rate orders equal to their stoichiometric coefficients, so they incorrectly extend this rule to overall reactions. Correct move: Always treat reaction order as an experimentally determined value; only use stoichiometry for rate orders when working with elementary reactions or elementary steps in a mechanism.
• Wrong move: Mixing up the axes and slope signs for zero-, first-, and second-order integrated rate law graphs. Why: The three different graphs have different axes and slope relationships to k, and students confuse them under time pressure. Correct move: Memorize the mnemonic "Z12" = Zero: [A] vs t, slope=-k; 1: ln[A] vs t, slope=-k; 2: 1/[A] vs t, slope=+k, and recall it before any graph problem.
• Wrong move: Drawing the activation energy for a catalyzed reaction with reactants/products at different energy levels than the uncatalyzed reaction. Why: Students confuse activation energy lowering with a change in overall reaction thermodynamics. Correct move: For catalyzed vs uncatalyzed energy profiles, always draw reactants and products at the exact same energy level, only adding lower activation barriers along the new reaction path.
• Wrong move: Leaving an intermediate in the final rate law for a mechanism with a fast initial equilibrium step. Why: Students forget that both catalysts and intermediates do not appear in the overall reaction, so they do not substitute out intermediate concentrations. Correct move: If the rate-determining step has an intermediate as a reactant, always use the equilibrium constant from the preceding fast step to substitute the intermediate concentration with reactant concentrations before writing the final rate law.
• Wrong move: Assuming a higher reaction order always means a faster reaction. Why: Students associate higher exponents in the rate law with faster rate, ignoring the magnitude of the rate constant and actual reactant concentrations. Correct move: Always compare reaction rates using the full rate law expression, accounting for reaction order, rate constant magnitude, and current reactant concentrations.

5. Quick Check: When To Use Which Sub-Topic

Test your understanding by matching each question to the correct sub-topic:

1. You need to find the concentration of a reactant after 50 seconds, given the initial concentration and rate constant for a first-order reaction.
2. You need to explain why increasing temperature increases reaction rate at the molecular level.
3. You need to determine if a proposed mechanism matches an experimentally determined rate law.
4. You need to compare the rate of consumption of a reactant to the rate of formation of product for a balanced overall reaction.
5. You need to explain how a catalyst increases reaction rate without being consumed.

6. Practice Questions (AP Chemistry Style)

Question 1 (Multiple Choice)

Which of the following correctly compares the rate of consumption of ${\text{N}}_2$ to the rate of formation of ${\text{NH}}_3$ for the reaction ${\text{N}}_2(g)+3{\text{H}}_2(g)\to 2{\text{NH}}_3(g)$? A) $-\frac{\Delta [{\text{N}}_2]}{\Delta t}=-\frac{\Delta [{\text{NH}}_3]}{\Delta t}$ B) $-\frac{\Delta [{\text{N}}_2]}{\Delta t}=\frac{1}{2}\frac{\Delta [{\text{NH}}_3]}{\Delta t}$ C) $\frac{\Delta [{\text{N}}_2]}{\Delta t}=2\frac{\Delta [{\text{NH}}_3]}{\Delta t}$ D) $-\frac{2\Delta [{\text{N}}_2]}{\Delta t}=\frac{\Delta [{\text{NH}}_3]}{\Delta t}$

Worked Solution: For any balanced reaction, the reaction rate is normalized by stoichiometric coefficients: $\text{rate}=-\frac{\Delta [{\text{N}}_2]}{\Delta t}=-\frac{1}{3}\frac{\Delta [{\text{H}}_2]}{\Delta t}=\frac{1}{2}\frac{\Delta [{\text{NH}}_3]}{\Delta t}$. ${\text{N}}_2$ is consumed, so its change in concentration is negative, while ${\text{NH}}_3$ is formed, so its change is positive. This matches exactly the relationship given in option B. Correct answer: B.

Question 2 (Free Response)

For the reaction $A+B\to C$, the following initial rate data is collected:

(a) Determine the order of reaction with respect to A and with respect to B. Justify your answer. (b) Write the overall rate law for the reaction and calculate the value of the rate constant $k$, including units. (c) A proposed mechanism for this reaction is: Step 1 (fast equilibrium): $2B\rightleftharpoons X$ Step 2 (slow): $X+A\to Y$ Step 3 (fast): $Y+A\to C$ Show that this mechanism is consistent with your rate law from part (b), and identify any intermediates.

Worked Solution: (a) Comparing Trial 1 and 2: $[B]$ is constant, $[A]$ doubles, and rate doubles. This means the order in A is 1, since $2^1=2$. Comparing Trial 2 and 3: $[A]$ is constant, $[B]$ doubles, and rate quadruples ($\frac{1.6\times 10^{-2}}{4.0\times 10^{-3}}=4=2^2$), so order in B is 2. (b) The overall rate law is $rate=k[A][B{]}^2$. Substitute values from Trial 1: $2.0\times 10^{-3}=k(0.10)(0.10{)}^2=k(1.0\times 10^{-3})$, so $k=2.0$ M⁻²s⁻¹. (c) The rate is determined by the slow Step 2: $rate=k[X][A]$. From the fast equilibrium Step 1: $K=\frac{[X]}{[B{]}^2}$, so $[X]=K[B{]}^2$. Substituting gives $rate=kK[A][B{]}^2$, which matches the experimental rate law. Intermediates are $X$ and $Y$, as both are produced in an early step and consumed in a later step, and do not appear in the overall reaction.

Question 3 (Application / Real-World Style)

The decomposition of benzoic acid, a common food preservative, is a first-order reaction with a rate constant of $1.2\times 10^{-3}$ year⁻¹ at 20°C. If a food manufacturer adds benzoic acid at an initial concentration of 0.050 M, what is the concentration remaining after 100 years of storage? Explain what this means for the shelf life of the product.

Worked Solution: For a first-order reaction, the integrated rate law is $\ln [\text{A}{]}_t=-kt+\ln [\text{A}{]}_0$. Substitute values: $kt=(1.2\times 10^{-3}{\text{year}}^{-1})(100\text{years})=0.12$, and $\ln (0.050)=-2.996$. So $\ln [\text{A}{]}_t=-0.12-2.996=-3.116$. Exponentiating both sides gives $[\text{A}{]}_t=e^{-3.116}\approx 0.044$ M. After 100 years of storage, 88% of the original preservative remains, so benzoic acid remains effective far beyond the typical 2-3 year shelf life of most processed food products.

7. Quick Reference Cheatsheet

8. What's Next / See Also

After completing this unit overview, you will move through each of the in-depth sub-topic study guides listed below to master every skill required for the AP exam. Kinetics is the foundational prerequisite for Unit 6 Thermodynamics, where you will connect reaction rate and activation energy to reaction spontaneity and Gibbs free energy, and Unit 7 Equilibrium, where equilibrium is defined as the point where forward and reverse reaction rates are equal. Without mastering core kinetics concepts like rate law, reaction order, and reaction mechanisms, you will not be able to interpret dynamic equilibrium or solve problems involving changing concentrations over time for acid-base, precipitation, or electrochemistry reactions. On the exam, multi-part FRQ questions often combine kinetics concepts from multiple sub-topics, so it is critical to master each individual topic before moving to more complex problems.

All in-depth sub-topic study guides for this unit are linked below:

• Reaction rate
• Introduction to rate law
• Concentration changes over time
• Elementary reactions
• Collision model
• Reaction energy profile
• Introduction to reaction mechanisms
• Multistep reaction energy profile
• Catalysis