AP Chemistry
In short
AP Chemistry has 9 units: Atomic Structure and Properties, Compound Structure and Properties, Properties of Substances and Mixtures and 6 more. Each unit lists its topics in course framework order, with a self-check list and how each is tested on the exam. Start with Atomic Structure and Properties or jump to the unit you are revising.
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What units are in AP Chemistry?
1.1 Moles and Molar Mass
1.2 Mass Spectra of Elements
- I can use the mass spectrum of a single element to determine the identity and relative abundance of its isotopes.
- I can estimate the average atomic mass of an element from the weighted average of its isotopic masses and relative abundances.
1.3 Elemental Composition of Pure Substances
1.4 Composition of Mixtures
1.5 Atomic Structure and Electron Configuration
1.6 Photoelectron Spectroscopy
1.7 Periodic Trends
- I can qualitatively explain periodic trends using Coulomb's law, the shell model, shielding, and effective nuclear charge.
1.8 Valence Electrons and Ionic Compounds
2.1 Types of Chemical Bonds
2.2 Intramolecular Force and Potential Energy
2.3 Structure of Ionic Solids
- I can describe how cations and anions in an ionic crystal are arranged in a systematic, periodic 3-D array to maximize attractive forces and minimize repulsive forces.
2.4 Structure of Metals and Alloys
2.5 Lewis Diagrams
2.6 Resonance and Formal Charge
2.7 VSEPR and Hybridization
3.1 Intermolecular and Interparticle Forces
3.2 Properties of Solids
- I can relate the vapor pressure, boiling point, and melting point of a substance to the strengths and types of intermolecular forces present.
- I can use particulate-level representations to communicate how intermolecular interactions establish macroscopic properties.
- I can explain the properties of ionic solids, such as low vapor pressure, high melting/boiling points, brittleness, and conditional electrical conductivity, based on strong interactions between ions.
- I can describe the structure and properties of covalent network solids (e.g., diamond, graphite, silicon dioxide) and explain their high melting points and rigidity based on strong covalent interactions.
- I can explain why molecular solids generally have low melting points and do not conduct electricity, based on their weak intermolecular forces and tightly held valence electrons.
- I can explain the properties of metallic solids (conductivity, malleability, ductility) and alloys based on the presence of free valence electrons and the arrangement of metal cores.
- I can describe how noncovalent interactions dictate the shape, functionality, and properties of large biomolecules or polymers.
3.3 Solids, Liquids, and Gases
3.4 Ideal Gas Law
3.5 Kinetic Molecular Theory
- I can describe the distribution of kinetic energies of particles at a given temperature using a Maxwell-Boltzmann distribution.
- I can relate the average kinetic energy of a particle to its average velocity using the equation KE = ½ mv².
3.6 Deviation from Ideal Gas Law
3.7 Solutions and Mixtures
3.8 Representations of Solutions
- I can use particulate models for mixtures to represent interactions between components and represent the relative concentrations of components.
3.9 Separation of Solutions and Mixtures
- I can use a chromatogram to infer the relative polarities of components in a mixture.
3.10 Solubility
- I can state that substances with similar intermolecular interactions tend to be miscible or soluble in one another.
3.11 Spectroscopy and the Electromagnetic Spectrum
3.12 Properties of Photons
3.13 Beer-Lambert Law
4.1 Introduction for Reactions
4.2 Net Ionic Equations
4.3 Representations of Reactions
- I can translate balanced chemical equations in their various forms into symbolic particulate representations.
4.4 Physical and Chemical Changes
4.5 Stoichiometry
4.6 Introduction to Titration
4.7 Types of Chemical Reactions
4.8 Introduction to Acid-Base Reactions
4.9 Oxidation-Reduction (Redox) Reactions
5.1 Reaction Rates
5.2 Introduction to Rate Law
5.3 Concentration Changes Over Time
- I can identify the rate law expression of a chemical reaction using concentration-time data.
- I can infer the order of a reaction from a graph of reactant concentration versus time.
- I can recognize that a linear plot of ln[reactant] versus time indicates a first-order reaction.
- I can recognize that a linear plot of 1/[reactant] versus time indicates a second-order reaction.
- I can determine the rate constant for zeroth, first, and second-order reactions using the slopes of their respective linear concentration-time graphs.
- I can calculate the half-life or rate constant for a first-order reaction using the equation t1/2 = 0.693/k.
- I can apply first-order kinetics principles to radioactive decay processes.
5.4 Elementary Reactions
5.5 Collision Model
- I can explain that successful collisions require both sufficient energy to overcome activation energy and the correct orientation.
- I can use the Maxwell-Boltzmann distribution curve to qualitatively estimate the fraction of collisions with sufficient energy and describe how this depends on temperature.
5.6 Reaction Energy Profile
- I can state that elementary reactions typically involve breaking some bonds and forming new ones.
- I can describe the reaction coordinate as the axis plotting the motions involved in rearranging reactants to products.
- I can identify the transition state, reactants, products, and activation energy on a reaction energy profile.
- I can explain the temperature dependence of reaction rates using the proportion of energetic collisions and the Arrhenius equation conceptually.
5.7 Introduction to Reaction Mechanisms
- I can verify that the elementary steps of a mechanism combine to yield the overall balanced chemical equation.
- I can identify a reaction intermediate as a species produced in one elementary step and consumed in another.
- I can understand that detecting a reaction intermediate provides experimental evidence to support a proposed reaction mechanism.
5.8 Reaction Mechanism and Rate Law
5.9 Pre-Equilibrium Approximation
5.10 Multistep Reaction Energy Profile
5.11 Catalysis
6.1 Endothermic and Exothermic Processes
6.2 Energy Diagrams
- I can describe a physical or chemical process using an energy diagram that shows whether it is endothermic or exothermic.
6.3 Heat Transfer and Thermal Equilibrium
6.4 Heat Capacity and Calorimetry
6.5 Energy of Phase Changes
- I can state that energy must be transferred to a system to melt or boil, and is released when freezing or condensing, while the temperature of a pure substance remains constant.
- I can use the molar enthalpy of a phase change to calculate energy absorbed or released, and recognize that complementary phase changes have equal but opposite energy changes.
6.6 Introduction to Enthalpy of Reaction
6.7 Bond Enthalpies
6.8 Enthalpy of Formation
6.9 Hess’s Law
7.1 Introduction to Equilibrium
- I can explain the relationship between a reversible process and the establishment of equilibrium using experimental observations.
- I can identify examples of reversible physical and chemical processes, such as phase changes, dissolution, acid-base reactions, and redox reactions.
- I can recognize that at equilibrium, no observable changes occur, both reactants and products are present, and their concentrations or partial pressures remain constant.
- I can describe the dynamic nature of equilibrium, where forward and reverse processes occur at equal rates.
- I can interpret graphs of concentration, partial pressure, or rate of reaction versus time to determine when chemical equilibrium is established.
7.2 Direction of Reversible Reactions
7.3 Reaction Quotient and Equilibrium Constant
7.4 Calculating the Equilibrium Constant
7.5 Magnitude of the Equilibrium Constant
7.6 Properties of the Equilibrium Constant
- I can invert the value of K when a chemical reaction is reversed.
- I can calculate the K of an overall reaction by multiplying the K values of the individual reactions that were summed together.
- I can apply all valid algebraic manipulations of K to the reaction quotient Q.
7.7 Calculating Equilibrium Concentrations
- I can determine the direction a reaction will proceed by comparing Q to K, and recognize that when Q = K, the system is at dynamic equilibrium.
7.8 Representations of Equilibrium
- I can use particulate representations to describe the relative numbers of reactant and product particles present prior to and at equilibrium, and relate them to the value of the equilibrium constant.
7.9 Introduction to Le Châtelier’s Principle
- I can predict the response of a system to stresses such as addition or removal of a chemical species, change in temperature, change in volume/pressure, or dilution using Le Châtelier’s principle.
- I can predict the effect that a stress will have on experimentally measurable properties such as pH, temperature, and color of a solution.
7.10 Reaction Quotient and Le Châtelier’s Principle
- I can explain that a disturbance to a system at equilibrium causes Q to differ from K, and that the system responds by bringing Q back into agreement with K.
- I can identify that changes in concentration change Q only, while changes in temperature change K, and describe how species redistribute to equalize Q and K.
7.11 Introduction to Solubility Equilibria
- I can describe the dissolution of a salt as a reversible process using the solubility-product constant, Ksp.
- I can calculate the solubility of a substance from its Ksp and use this to predict the relative solubility of different substances.
- I can relate solubility rules to Ksp, recognizing that Ksp values >1 correspond to soluble salts.
7.12 Common-Ion Effect
8.1 Introduction to Acids and Bases
- I can calculate pH and pOH from hydronium and hydroxide ion concentrations using negative logarithms.
- I can use the autoionization constant of water (Kw = 1.0 × 10^-14 at 25°C) to relate hydronium and hydroxide concentrations.
- I can explain that a neutral solution is defined by pH = pOH, which is 7.0 at 25°C.
8.2 pH and pOH of Strong Acids and Bases
- I can calculate the pH and pOH of a strong acid or strong base solution based on its initial concentration.
8.3 Weak Acid and Base Equilibria
- I can explain that weak acids only partially ionize, leaving the vast majority of molecules un-ionized, and that [H3O+] is much less than the initial acid concentration.
- I can calculate the pH of a weak acid solution using the initial acid concentration and Ka or pKa.
- I can explain that weak bases only partially ionize, leaving the vast majority of molecules un-ionized, and that [OH-] does not equal the initial base concentration.
- I can calculate the pH of a weak base solution using the initial base concentration and Kb or pKb.
- I can calculate the percent ionization of a weak acid or base from its pKa/pKb and initial concentration, or from equilibrium concentrations.
- I can relate the acid ionization constant and base ionization constant for a conjugate acid-base pair using Kw.
8.4 Acid-Base Reactions and Buffers
8.5 Acid-Base Titrations
- I can use a titration curve plotting pH against the volume of titrant added to summarize results from a titration.
- I can use the relationship at the equivalence point (moles of titrant added equals moles of analyte originally present) to obtain the concentration of the analyte for monoprotic acids or bases.
- I can determine the pKa of a weak acid or base from the pH at the half-equivalence point in a titration, where the conjugate acid and base have equal concentrations.
- I can determine the pH at the equivalence point based on the major species in solution, recognizing that strong acid-strong base titrations are neutral, while weak acid or weak base titrations produce basic or acidic solutions respectively.
- I can use titration curves for polyprotic acids to determine the number of acidic protons, identify the major species present at any point, and find the pKa associated with each proton.
8.6 Molecular Structure of Acids and Bases
- I can infer the relative strength of protons in a molecule based on its molecular structure, including the effects of electronegativity, inductive effects, and resonance on conjugate base stabilization.
8.7 pH and pKa
- I can predict the protonation state of an acid or base by comparing the pH of a solution to the pKa of the acid.
- I can explain how acid-base indicators change properties based on their protonated versus deprotonated states in response to pH.
- I can select an appropriate acid-base indicator for a titration by choosing one with a pKa close to the pH at the equivalence point.
8.8 Properties of Buffers
- I can state that a buffer contains large concentrations of a conjugate acid-base pair and describe how each component reacts with added acid or base.
8.9 Henderson-Hasselbalch Equation
- I can explain how the pH of a buffer is related to the pKa and the concentration ratio of the conjugate acid-base pair using the Henderson-Hasselbalch equation, and understand why adding small amounts of acid or base does not significantly change the pH.
8.10 Buffer Capacity
- I can explain that increasing the concentration of buffer components (while keeping their ratio constant) keeps the pH the same but increases the buffer capacity.
- I can predict that a buffer with more conjugate acid than base has a greater capacity for added base, and a buffer with more conjugate base than acid has a greater capacity for added acid.
8.11 pH and Solubility
9.1 Introduction to Entropy
- I can explain how entropy increases when matter becomes more dispersed, such as during phase changes, volume increases, or reactions that produce a greater number of moles of gas.
- I can explain how entropy increases when energy is dispersed, using the kinetic molecular theory to relate an increase in temperature to a broader distribution of kinetic energy.
9.2 Absolute Entropy and Entropy Change
- I can use the equation ΔS°_reaction = ΣS°_products − ΣS°_reactants to calculate the entropy change for a process.
9.3 Gibbs Free Energy and Thermodynamic Favorability
- I can identify the standard state conditions for Gibbs free energy (pure substances, 1.0 M solutions, or gases at 1.0 atm/bar).
- I can use the term 'thermodynamically favored' for processes where ΔG° < 0 and avoid equating it with the rate of the reaction.
- I can calculate the standard Gibbs free energy change of a reaction using the standard Gibbs free energies of formation of the reactants and products.
- I can recognize situations where both enthalpy and entropy must be considered to determine if a process is thermodynamically favored.
- I can calculate ΔG° directly from the values of ΔH°, ΔS°, and temperature.
- I can predict the temperature conditions under which a process will be thermodynamically favored based on the signs of ΔH° and ΔS°, without needing to calculate ΔG° in certain cases.
9.4 Thermodynamic and Kinetic Control
9.5 Free Energy and Equilibrium
9.6 Free Energy of Dissolution
9.7 Coupled Reactions
9.8 Galvanic (Voltaic) and Electrolytic Cells
9.9 Cell Potential and Free Energy
- I can explain that thermodynamically favored redox reactions produce a positive voltage, while unfavored reactions produce a negative voltage and require an externally applied potential.
- I can calculate the standard cell potential of an electrochemical cell by identifying the oxidation and reduction half-reactions and their standard reduction potentials.
- I can use the equation ΔG° = −nFE° to relate standard cell potential to standard Gibbs free energy change, and use the sign of E° to determine if a reaction is thermodynamically favored.
9.10 Cell Potential Under Nonstandard Conditions
- I can explain that cell potential varies with the concentrations of active species and that the farther a reaction is from equilibrium, the greater the magnitude of the cell potential.
- I can state that Le Châtelier's principle does not apply to electrochemical systems because they are not in equilibrium.
- I can predict how deviations from standard conditions (Q = 1) change the magnitude of the cell potential relative to E°, and determine the direction of spontaneous electron flow in concentration cells.
- I can qualitatively use the Nernst equation to reason about the effects of concentration on cell potential without relying solely on algorithmic calculations.
9.11 Electrolysis and Faraday’s Law
- I can use Faraday’s laws to determine the stoichiometry of a redox reaction with respect to the number of electrons transferred, mass of material deposited/removed, current, time elapsed, and charge of ionic species.
How should you use these resources?
Revise
Work through your course unit by unit, in the order it is taught, so nothing gets missed.
Test yourself
Check each topic against what you should be able to do, and practise on exam-style questions as they are added.
Improve
Go back to what you could not do, or work with a tutor to close the gaps before your exams.
AP Chemistry: common questions
What units are in AP Chemistry?
AP Chemistry has 9 units: Unit 1: Atomic Structure and Properties; Unit 2: Compound Structure and Properties; Unit 3: Properties of Substances and Mixtures; Unit 4: Chemical Reactions; Unit 5: Kinetics; Unit 6: Thermochemistry; Unit 7: Equilibrium; Unit 8: Acids and Bases; Unit 9: Thermodynamics and Electrochemistry.
How should I revise AP Chemistry?
Work through one unit at a time in order, since later units build on earlier ones. For each topic, check you can do everything on its list before moving on, and go back to the ones you could not.
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