AP Chemistry Unit 3 Review: Properties of Substances and Mixtures

Every Unit 3 topic, what's actually tested, and what the exam explicitly excludes -- the biggest unit on the exam, verified against the current CED.

Unit 3 is the heavyweight of AP Chemistry: 18–22% of the multiple-choice section, roughly double the weight of most other units, spread across 13 topics and 14–15 class periods. It's also where the course shifts from individual atoms and bonds (Units 1–2) to how collections of particles behave — intermolecular forces, gas laws, solutions, and light. Here's everything in it, topic by topic, verified against the current CED — including three exclusion statements that cut real content most courses still teach.

Unit 3 at a Glance

Unit 3 (Properties of Substances and Mixtures) is nearly 20% of the exam by itself. See the full AP Chemistry units breakdown for how it fits alongside the other 8 units, and the Course and Exam Description for the full framework.

Intermolecular and Interparticle Forces

Four forces hold separate molecules or ions together, in increasing order of typical strength:

• London dispersion forces (LDF) arise from temporary, fluctuating dipoles and exist between all molecules. They're often the strongest net force between large, highly polarizable molecules -- dispersion forces increase with more electrons, a larger electron cloud, and the presence of pi bonding.
• Dipole-dipole forces exist between polar molecules and are generally stronger than LDF between molecules of comparable size, since they add on top of the dispersion force already present.
• Hydrogen bonding is a strong dipole-dipole interaction that occurs when H is covalently bonded to N, O, or F and attracted to a lone pair on N, O, or F in a nearby molecule.
• Ion-dipole forces occur between an ion and a polar molecule and are typically the strongest of the four -- this is what makes ionic compounds dissolve in water.

Precise vocabulary matters: the CED explicitly states that the term "London dispersion forces" should not be used synonymously with "van der Waals forces." Van der Waals forces is a broader, older umbrella term the exam avoids -- always name the specific force (LDF, dipole-dipole, hydrogen bonding, or ion-dipole).

Properties of Solids

Four categories of solid, distinguished by the type of particle interaction holding them together:

• Ionic solids (e.g., NaCl) have high melting points, low vapor pressure, are brittle, and conduct electricity only when molten or dissolved -- never as a solid.
• Covalent network solids (e.g., diamond, silicon dioxide) are held together by a continuous 3-D (or layered 2-D, like graphite) network of covalent bonds. Very high melting points, rigid and hard -- except graphite, whose layers slide past each other easily.
• Molecular solids (e.g., ice, dry ice, sugar) are individual covalently bonded molecules held together by comparatively weak intermolecular forces. Low melting points; don't conduct electricity.
• Metallic solids conduct electricity and heat via a "sea" of delocalized valence electrons, and are malleable and ductile because metal cores can rearrange without breaking bonds.

Solids, Liquids, and Gases

Solids can be crystalline (regular 3-D arrangement) or amorphous (no regular arrangement), but in both cases particle motion is limited. Liquid particles stay in close contact while continually moving and colliding -- which is why solids and liquids of the same substance have similar molar volumes. Gas particles are in constant, mostly independent motion, so a gas has neither a definite volume nor a definite shape.

Not tested: understanding or interpreting phase diagrams (the P-vs-T graphs showing solid, liquid, and gas regions) is explicitly excluded from the AP Exam -- don't spend study time memorizing triple points and critical points.

Ideal Gas Law

The macroscopic properties of an ideal gas are related by PV = nRT. In a mixture of gases, each component's partial pressure is proportional to its mole fraction, and the total pressure is the sum of the partial pressures: P_total = P_A + P_B + …

Worked example: A 2.00 L container holds 0.0400 mol N₂ and 0.0600 mol O₂ at 300 K. Total pressure: P = nRT/V = (0.100 mol)(0.08206 L·atm/mol·K)(300 K) / 2.00 L = 1.23 atm. N₂'s mole fraction is 0.0400/0.100 = 0.400, so its partial pressure is 0.400 × 1.23 atm = 0.492 atm.

Kinetic Molecular Theory

KMT links macroscopic gas behavior to particle motion. Average kinetic energy relates to velocity by KE = ½mv², and the Kelvin temperature of a sample is proportional to the average kinetic energy of its particles. The Maxwell-Boltzmann distribution graphs the spread of particle energies at a given temperature -- as temperature rises, the distribution broadens and shifts toward higher energy.

Deviation from Ideal Gas Law

Real gases deviate from PV = nRT for two reasons: interparticle attractions (most significant near condensation conditions) and the particles' own volume (most significant at extremely high pressure). Both effects are ignored by the ideal gas law's assumptions.

Solutions and Mixtures

A solution (homogeneous mixture) has uniform macroscopic properties throughout; a heterogeneous mixture's properties depend on location. Molarity, M = n_solute / L_solution, is the only concentration calculation required on the exam.

Worked example: How many grams of NaOH (molar mass 40.0 g/mol) are needed to make 250. mL of a 0.500 M solution? n = M × V = 0.500 mol/L × 0.250 L = 0.125 mol; mass = 0.125 mol × 40.0 g/mol = 5.00 g.

Not tested (the big one): colligative properties, and calculations of molality, percent by mass, and percent by volume, are all explicitly excluded from the AP Exam. If a resource drills molality or boiling-point elevation for "AP Chem," it's teaching content the current exam doesn't assess.

Representations of Solutions

Particulate drawings of solutions communicate relative concentration and the interactions among components. You should be able to sketch a solution showing more solvent particles than solute particles at low concentration, and interpret drawings showing ion-dipole or hydrogen-bonding interactions between solute and solvent.

Separation of Solutions and Mixtures

Because dissolved components can't be separated by filtration, separation techniques exploit differences in intermolecular interaction strength instead:

• Chromatography (paper, thin-layer, column) separates species by their differing attraction to a mobile phase versus a stationary phase; the resulting chromatogram indicates relative polarity.
• Distillation separates species by differences in vapor pressure that result from differing intermolecular interaction strength.

Solubility

The guiding principle is simple: substances with similar intermolecular interactions tend to be miscible or soluble in one another ("like dissolves like"). A polar solvent like water dissolves ionic and polar solutes well; a nonpolar solvent dissolves nonpolar solutes well.

Spectroscopy and the Electromagnetic Spectrum

Different regions of the EM spectrum correspond to different types of molecular or electronic transitions:

• Microwave radiation → molecular rotational transitions
• Infrared radiation → molecular vibrational transitions
• UV/visible radiation → electronic energy-level transitions

Properties of Photons

A photon's energy relates to frequency by Planck's equation, E = hν, and frequency relates to wavelength by c = λν.

Worked example: What's the energy of a photon with wavelength 500 nm (5.00 × 10⁻⁷ m)? First, ν = c/λ = (3.00 × 10⁸⁸ m/s) / (5.00 × 10⁻⁷ m) = 6.00 × 10⁹⁴ Hz. Then E = hν = (6.626 × 10⁻₁₃⁴ J·s)(6.00 × 10⁹⁴ Hz) = 3.98 × 10⁻₁⁹ J per photon.

Beer-Lambert Law

The Beer-Lambert law, A = εbc, relates absorbance to molar absorptivity (ε), path length (b), and concentration (c). When path length and wavelength are held constant, absorbance is directly proportional to concentration -- the basis for every calibration curve you'll build in the Unit 3 lab.

Worked example: A solution has molar absorptivity ε = 5.60 × 10³ L/(mol·cm) at its wavelength of maximum absorbance and is measured in a 1.00 cm cuvette with absorbance A = 0.420. Solving for concentration: c = A / (εb) = 0.420 / (5.60 × 10³ × 1.00) = 7.50 × 10⁻⁵ M.

Common Mistakes in Unit 3

• Calling every intermolecular force "van der Waals." Name the specific force -- LDF, dipole-dipole, hydrogen bonding, or ion-dipole -- since the CED explicitly separates these terms.
• Forgetting that polar molecules also have LDF. Dipole-dipole forces add on top of London dispersion forces; they don't replace them.
• Using Celsius in PV = nRT. The ideal gas law and KMT both require absolute (Kelvin) temperature.
• Assuming higher absorbance always means higher concentration between different substances. The Beer-Lambert law only gives a direct concentration comparison when ε and b are held constant -- different molecules absorb light differently even at identical concentrations.
• Studying molality or colligative properties for the AP Exam. Both are explicitly excluded; molarity is the only required concentration calculation.

How Unit 3 Connects to the Rest of the Course

• The intermolecular forces here directly explain the boiling points, solubility, and reaction rates covered throughout every later unit.
• Molarity from this unit is the foundation for every equilibrium, acid-base, and electrochemistry calculation for the rest of the course.
• The Beer-Lambert law and spectroscopy resurface directly in kinetics experiments that track reaction progress by absorbance.

Related Resources

• AP Chemistry Units (All 9)
• AP Chemistry Unit 1 Review
• AP Chemistry Unit 2 Review
• AP Chemistry Unit 4 Review
• AP Chemistry Unit 5 Review
• AP Chemistry Unit 8 Review
• AP Chemistry Course and Exam Description
• AP Chemistry FRQ Calculator
• AP Chemistry Unit 9 Review

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Sourced from College Board's official AP Chemistry Course and Exam Description, Effective Fall 2024. This page describes the document's real content and current exclusion statements; it is not a copy of it and is not affiliated with or endorsed by College Board.