Review AP Chem Unit 3 to understand how intermolecular forces, solid types, gas behavior, solution chemistry, and spectroscopy connect molecular structure to observable properties. This unit carries 18-22% of the exam and spans topics from London dispersion forces through Beer-Lambert law.
Fiveable's AP Chemistry Unit 3 review covers 13 topics, with guided practice questions for the unit.
Use this hub to review all 13 topics, practice with available questions and FRQs, and estimate your score with the AP score calculator.
Unit 3 asks you to explain properties of matter by reasoning from the particle level upward. Every macroscopic observation, whether a high boiling point, a gas that deviates from ideal behavior, or a colored solution that absorbs light, has a particle-level explanation rooted in forces, motion, or energy.
Unit 3 covers how intermolecular forces determine solid and liquid properties, how the ideal gas law and kinetic molecular theory describe gas behavior, how solutions are prepared and represented, how separation techniques exploit IMF differences, and how spectroscopy and Beer-Lambert law connect light absorption to molecular structure and concentration.
London dispersion, dipole-dipole, hydrogen bonding, and ion-dipole forces determine boiling point, melting point, vapor pressure, and solubility. Stronger forces mean higher boiling points and lower vapor pressures. Polarizability, molecular size, and the presence of N-H, O-H, or F-H bonds determine which forces are present and how strong they are.
PV = nRT describes ideal gas behavior. Dalton's law adds partial pressures for mixtures. Kinetic molecular theory explains pressure and temperature at the particle level using KE = 1/2 mv² and the Maxwell-Boltzmann distribution. Real gases deviate from ideal behavior at low temperature and high pressure due to intermolecular attractions and finite particle volume.
Molarity M = n/L and dilution M1V1 = M2V2 are the core solution calculations. Solubility follows the like-dissolves-like principle. Chromatography and distillation separate mixtures using IMF differences. Beer-Lambert law A = εbc links absorbance to concentration, and photon energy E = hν links light frequency to electronic, vibrational, or rotational transitions.
Whether you are comparing boiling points of alkanes, explaining why a real gas compresses less than predicted, or reading a calibration curve, the reasoning chain is the same: identify the particles, identify the forces or energy relationships between them, and connect those to the observable property. Unit 3 builds this reasoning pattern across solids, liquids, gases, solutions, and light-matter interactions.
Identify and rank London dispersion, dipole-dipole, hydrogen bonding, and ion-dipole forces from molecular structure. Connect IMF strength to boiling point, melting point, and vapor pressure.
Classify solids as ionic, covalent network, metallic, or molecular. Explain melting point, hardness, conductivity, and brittleness from particle-level structure and force type.
Use particulate models to represent particle arrangement, motion, and spacing in each phase. Connect phase properties like compressibility and fluidity to particle behavior.
Apply PV = nRT to solve for any gas variable. Use Dalton's law and mole fractions to find partial pressures in gas mixtures. Interpret P-V-T-n graphs.
Explain gas pressure and temperature using particle motion and elastic collisions. Interpret Maxwell-Boltzmann distribution curves for different temperatures and molecular masses.
Explain why real gases deviate from ideal behavior at low temperature (intermolecular attractions) and high pressure (finite particle volume). Predict direction of deviation.
Distinguish homogeneous from heterogeneous mixtures. Calculate molarity using M = n/L and apply the dilution equation M1V1 = M2V2.
Draw and interpret particulate diagrams showing relative concentration and solute-solvent interactions, including correct ion orientation and dissociation for electrolytes.
Explain chromatography and distillation results using IMF differences between components and the stationary or mobile phase. Interpret Rf values for relative polarity.
Apply the like-dissolves-like principle to predict solubility of ionic and molecular compounds in polar and nonpolar solvents. Identify the specific IMFs that drive dissolution.
Match microwave, infrared, and UV-visible radiation to rotational, vibrational, and electronic transitions. Relate wavelength and frequency to photon energy.
Use c = λν and E = hν to calculate photon frequency, wavelength, and energy. Connect photon absorption or emission to quantized energy changes in atoms and molecules.
Apply A = εbc to calculate absorbance, concentration, path length, or molar absorptivity. Use calibration curves to determine unknown concentrations from measured absorbance values.
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Review Properties of Photons with attention to how the concept appears in AP-style source and evidence questions.
Review Deviation from Ideal Gas Law with attention to how the concept appears in AP-style source and evidence questions.
Review Ideal Gas Law with attention to how the concept appears in AP-style source and evidence questions.
Review Representations of Solutions with attention to how the concept appears in AP-style source and evidence questions.
Intermolecular forces (IMFs) are attractions between separate particles. Their strength determines physical properties like boiling point, melting point, and vapor pressure. For any comparison question, identify which forces are present, then rank their strength based on molecular structure.
Given two molecules, identify all IMFs present in each and predict which has the higher boiling point, explaining your reasoning in terms of force type and strength.
| Force type | Molecules involved | Relative strength | Structural requirement |
|---|---|---|---|
| London dispersion | All molecules | Weakest to strongest (size-dependent) | All molecules; increases with electron count and surface area |
| Dipole-dipole | Polar molecules | Moderate | Permanent dipole moment (asymmetric polar bonds) |
| Hydrogen bonding | Molecules with N-H, O-H, or F-H | Strong | H bonded to N, O, or F; lone pair acceptor nearby |
| Ion-dipole | Ion + polar molecule | Strongest IMF | Ion present; polar solvent molecule |
Solids are classified by the particles present and the forces holding them together. Each type has a predictable set of macroscopic properties. The three phases differ in particle arrangement, freedom of motion, and compressibility.
Given a substance's melting point, conductivity, and hardness, identify its solid type and explain the particle-level basis for each property.
| Solid type | Particles | Forces | Melting point | Conductivity |
|---|---|---|---|---|
| Ionic | Cations and anions | Electrostatic (Coulombic) | High | Only when molten or dissolved |
| Covalent network | Atoms | Covalent bonds | Very high | Usually none (diamond); yes (graphite) |
| Metallic | Metal cations + electron sea | Metallic bonds | Variable (generally high) | Yes (solid and liquid) |
| Molecular | Molecules | IMFs (LDF, dipole, H-bond) | Low to moderate | No |
The ideal gas law PV = nRT relates the four macroscopic properties of a gas sample. For mixtures, Dalton's law of partial pressures applies: each gas contributes independently to the total pressure based on its mole fraction.
A container holds 0.50 mol N2 and 0.25 mol O2 at a total pressure of 3.0 atm. Calculate the partial pressure of each gas and the mole fraction of N2.
Kinetic molecular theory (KMT) explains ideal gas behavior at the particle level. Real gases deviate from ideal predictions when intermolecular attractions or finite particle volumes become significant.
Sketch two Maxwell-Boltzmann curves: one for N2 at 300 K and one at 600 K. Describe how the peak position and curve shape change and explain why.
| Condition | Dominant effect | Deviation from ideal | Explanation |
|---|---|---|---|
| Low temperature | Intermolecular attractions | Pressure lower than ideal | Particles slow down and attract each other before hitting walls |
| High pressure | Finite particle volume | Volume larger than ideal | Particle volume is no longer negligible relative to container volume |
| High temperature, low pressure | Neither effect significant | Behaves nearly ideally | Particles move fast and are far apart |
A solution is a homogeneous mixture with uniform macroscopic properties throughout. Molarity is the primary concentration unit. Particulate diagrams show both the relative number of solute particles (concentration) and how solute and solvent particles interact.
Draw a particulate diagram showing NaCl dissolved in water. Label the ion-dipole interactions and show the correct orientation of water molecules around Na+ and Cl-.
Separation of mixtures and solubility both depend on differences in intermolecular interactions. The like-dissolves-like principle governs which solutes dissolve in which solvents, and the same logic explains how chromatography and distillation separate components.
Two compounds are separated by paper chromatography. Compound A has a higher Rf than compound B. Which is more polar? Explain using IMF interactions with the stationary phase.
| Technique | Basis of separation | IMF principle |
|---|---|---|
| Chromatography | Differential interaction with stationary vs. mobile phase | More polar compound interacts more strongly with polar stationary phase |
| Distillation | Differences in vapor pressure / boiling point | Weaker IMFs give higher vapor pressure; that component distills first |
| Filtration | Particle size | Does not separate true solutions; only removes suspended solids |
Different regions of the electromagnetic spectrum interact with matter in different ways depending on the energy of the photons. The two key equations are c = λν (connecting wavelength and frequency) and E = hν (connecting frequency and photon energy).
A photon has a wavelength of 500 nm. Calculate its frequency and energy. Identify which type of molecular transition this photon could cause.
| Spectral region | Transition type | Relative photon energy |
|---|---|---|
| Microwave | Molecular rotation | Lowest |
| Infrared | Molecular vibration | Moderate |
| Ultraviolet/Visible | Electronic transitions | Highest |
The Beer-Lambert law A = εbc quantitatively relates the absorbance of a solution to the concentration of the absorbing species, the path length of the light through the solution, and the molar absorptivity. It is the basis for using a spectrophotometer to determine unknown concentrations.
A solution has an absorbance of 0.45 at λmax. The molar absorptivity is 150 M⁻¹cm⁻¹ and the path length is 1.00 cm. Calculate the concentration of the absorbing species.
Try stimulus-based AP practice questions and written prompts after you review the notes.
A timed, AP-style test built only from unit 3 questions — multiple choice plus free response, scored with AI feedback. Use the standard format or build your own.
| Term | Definition |
|---|---|
| London dispersion forces | Temporary attractive forces between all molecules arising from instantaneous fluctuating dipoles. Strength increases with molecular size, electron count, and surface contact area. |
| Hydrogen bonding | A strong IMF requiring a hydrogen atom bonded directly to N, O, or F interacting with a lone pair on another N, O, or F atom. Responsible for the high boiling point of water and many biological properties. |
| Polarizability | The ease with which an electron cloud distorts in response to an electric field. Larger molecules with more electrons and pi bonds have higher polarizability and stronger London dispersion forces. |
| PV = nRT | The ideal gas law relating pressure (atm), volume (L), moles, the gas constant R (0.08206 L·atm·mol⁻¹·K⁻¹), and absolute temperature (K). |
| Dalton's Law of Partial Pressure | The total pressure of a gas mixture equals the sum of the partial pressures of each component. Each partial pressure equals the mole fraction of that gas times the total pressure. |
| Maxwell-Boltzmann distribution | A graph showing the distribution of kinetic energies or speeds of gas particles at a given temperature. Higher temperature shifts the curve right and broadens it; heavier molecules shift it left. |
| Molarity | Concentration expressed as moles of solute per liter of solution (M = n/L). The denominator is total solution volume, not solvent volume alone. |
| like dissolves like | Polar and ionic solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents. Dissolution is favorable when solute-solvent IMFs are similar in type and strength to the forces being disrupted. |
| Stationary Phase | The fixed medium in chromatography (e.g., silica gel, paper) with which solute components interact. More polar compounds interact more strongly with a polar stationary phase and travel more slowly. |
| Distillation | A separation technique that exploits differences in vapor pressure (IMF strength) between liquid components. The component with weaker IMFs has higher vapor pressure and vaporizes first. |
| E = hν | Planck's equation relating photon energy (J) to frequency (Hz) via Planck's constant h = 6.626 × 10⁻³⁴ J·s. Higher frequency means higher energy per photon. |
| Absorbance | A unitless measure of how much light a solution absorbs at a given wavelength. Directly proportional to concentration when path length and wavelength are held constant (Beer-Lambert law). |
| Calibration curve | A graph of absorbance vs. known concentration for standard solutions. Used to determine the concentration of an unknown solution by comparing its measured absorbance to the linear trend. |
| particulate-level model | A representation of matter showing individual particles, their arrangement, spacing, and interactions. Used in AP Chem to explain macroscopic properties from molecular structure. |
PV = nRT and KE = (3/2)kT both require absolute temperature. Always convert: K = °C + 273. Using 25 instead of 298 will give a wrong answer every time.
London dispersion forces are one specific type of intermolecular force. The term van der Waals forces is broader and includes dipole-dipole interactions as well. Do not use the two terms interchangeably on the exam.
Molarity is moles of solute per liter of solution, not per liter of solvent. When a solid is dissolved in water to make 250 mL of solution, the denominator is 0.250 L of solution, not the volume of water added.
At low temperature, attractions handle and pressure is lower than ideal. At very high pressure, particle volume dominates and the gas occupies more volume than ideal. The direction of deviation depends on which effect is larger under the given conditions.
A higher Rf means the compound traveled farther with the mobile phase, indicating weaker interaction with the polar stationary phase, which means the compound is less polar. Students often reverse this relationship.
AP Chem frequently asks you to draw or interpret particle-level diagrams for solutions, phases, or gas mixtures. For Unit 3, this means showing correct IMF interactions between solute and solvent, correct ion dissociation for electrolytes, correct relative particle counts for concentration comparisons, and correct phase representations with appropriate spacing and motion implied.
Many Unit 3 questions ask you to justify a ranking of boiling points, vapor pressures, or solubilities by naming the specific IMF, explaining why it is present from molecular structure, and connecting its strength to the observed property. A complete response names the force, explains its structural basis, and links it to the macroscopic trend.
Gas law and Beer-Lambert law problems often pair a calculation with a conceptual question. For example, after solving PV = nRT for moles, you may be asked to explain why a real gas deviates from that value. After using A = εbc to find concentration, you may be asked to explain what happens to absorbance if path length doubles. Practicing both the calculation and the explanation together prepares you for multi-part questions.
For any molecule, identify all IMFs present (LDF, dipole-dipole, H-bonding, ion-dipole), rank their strength, and explain how that ranking predicts boiling point, vapor pressure, or solubility.
Given melting point, conductivity, hardness, or brittleness data, identify the solid type (ionic, covalent network, metallic, molecular) and explain each property from particle-level structure.
Use PV = nRT with T in Kelvin. Calculate partial pressures using PA = XA × Ptotal and verify that partial pressures sum to total pressure.
Predict how the distribution curve shifts with increasing temperature or increasing molecular mass. Explain what the curve shape means for average kinetic energy and particle speed.
State whether a real gas will have higher or lower pressure than ideal at a given condition, and attribute the deviation to either intermolecular attractions (low T) or finite particle volume (high P).
Apply M = n/L using total solution volume. Use M1V1 = M2V2 for dilution problems. Draw or interpret particulate diagrams showing correct relative concentrations.
Use A = εbc to solve for any variable. Read an unknown concentration from a calibration curve by locating the measured absorbance on the y-axis and reading the corresponding concentration on the x-axis.
Read the topic guides for 3.1, 3.2, and 3.3. For each IMF type, write one example molecule and one property it explains. For each solid type, list the particles, forces, and two macroscopic properties. Practice drawing particulate models for each phase.
Review PV = nRT and Dalton's law with the topic guides for 3.4 and 3.5. Solve at least five ideal gas law problems, including mixture problems using mole fractions. Then read 3.6 and practice explaining real gas deviations in writing before checking your reasoning.
Use the 3.7 topic guide to practice molarity and dilution calculations. Then use the 3.8 guide to practice drawing particulate diagrams for strong electrolytes, weak electrolytes, and nonelectrolytes. Check that your diagrams show correct ion orientation around water molecules.
Review chromatography and distillation using the 3.9 topic guide. For each technique, write a one-sentence explanation of why a specific component moves faster or distills first, naming the IMF involved. Then review 3.10 and practice predicting solubility using like-dissolves-like.
Use the topic guides for 3.11 and 3.12 to memorize the three spectral region-to-transition pairings and practice E = hν and c = λν calculations with unit tracking. Finish with 3.13: practice A = εbc calculations and reading calibration curves. Use available practice questions and FRQs to test all three topics together.
Open the individual guides for Unit 3 when you want a closer review of one topic.
Use AP-style practice after you review the notes so you can check what you understand.
Practice free-response reasoning and compare your answer with scoring guidance.
Watch past review streams filtered to Unit 3 when you want a video walkthrough.
Open the Fiveable one-page unit review, then explore visual cheatsheets for a quick refresher.
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get printablesAP Chem Unit 3 covers 13 topics across intermolecular forces, states of matter, and solution chemistry. Key topics include Intermolecular and Interparticle Forces, Properties of Solids, Ideal Gas Law, Kinetic Molecular Theory, Deviation from Ideal Gas Law, Solubility, Spectroscopy and the Electromagnetic Spectrum, Properties of Photons, and Beer-Lambert Law. Here's the full topic list:
- 3.1 Intermolecular and Interparticle Forces
- 3.2 Properties of Solids
- 3.3 Solids, Liquids, and Gases
- 3.4 Ideal Gas Law
- 3.5 Kinetic Molecular Theory
- 3.6 Deviation from Ideal Gas Law
- 3.7 Solutions and Mixtures
- 3.8 Representations of Solutions
- 3.9 Separation of Solutions and Mixtures
- 3.10 Solubility
- 3.11 Spectroscopy and the Electromagnetic Spectrum
- 3.12 Properties of Photons
- 3.13 Beer-Lambert Law See AP Chem Unit 3 for matched practice on all of these.
AP Chem Unit 3 makes up 18-22% of the AP exam, making it one of the heavier-weighted units. It covers solubility, intermolecular forces, gas laws, solutions and mixtures, and spectroscopy topics like the Beer-Lambert Law. Expect a solid chunk of both MCQ and FRQ questions drawn from this material.
The AP Chem Unit 3 progress check in AP Classroom includes both MCQ and FRQ parts that pull from all 13 topics in this unit. The MCQ section tests conceptual understanding of intermolecular forces, gas laws, solubility, and spectroscopy. The FRQ part typically asks you to interpret data, apply the Beer-Lambert Law, or explain solution behavior at the particle level. The progress check is designed to mirror real exam difficulty, so it's worth treating it like a mini exam. For extra practice on the same topics, check out AP Chem Unit 3.
AP Chem Unit 3 FRQs most often target solubility, intermolecular forces, gas law calculations, and spectroscopy concepts like the Beer-Lambert Law. These questions ask you to explain phenomena at the particle level, interpret graphs or spectra, or set up and solve quantitative problems. The best way to practice is to write out full explanations, not just answers, because graders reward specific reasoning. Focus on topics 3.1, 3.4, 3.10, and 3.13 first since they show up most in free-response contexts. You can find practice sets and worked examples at AP Chem Unit 3.
The best place to find AP Chem Unit 3 practice questions, including MCQ and full practice test sets, is AP Chem Unit 3. That page has targeted multiple-choice questions covering solubility, intermolecular forces, the Ideal Gas Law, Kinetic Molecular Theory, and spectroscopy. When you're doing MCQ practice, pay close attention to questions on Deviation from Ideal Gas Law and Beer-Lambert Law since those tend to trip students up. Mixing topic-specific drills with timed full-unit practice tests is the most effective approach.
Start AP Chem Unit 3 by building a strong foundation in intermolecular forces (3.1), because those concepts carry through solubility, states of matter, and solution behavior later in the unit. Then work through the gas law topics (3.4, 3.5, 3.6) together since Kinetic Molecular Theory explains why ideal gas assumptions break down. Here's a practical study sequence:
1. Learn intermolecular forces and how they predict physical properties.
2. Work through Ideal Gas Law problems, then study where and why gases deviate.
3. Study Solutions and Mixtures (3.7-3.10), focusing on solubility rules and what affects them.
4. Finish with spectroscopy (3.11-3.13), including the Beer-Lambert Law and how photon energy relates to the electromagnetic spectrum. For each topic, practice explaining concepts in words before doing calculations. Unit 3 is 18-22% of the exam, so it rewards consistent review. Head to AP Chem Unit 3 for practice questions organized by topic.