AP Chemistry Unit 3 ReviewProperties of Substances and Mixtures

AP Chemistry Unit 3 is about one big idea, that the strength of attractions between particles explains almost everything you can observe about a substance, from its boiling point to whether it dissolves in water. The unit covers intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding), the particle-level picture of solids, liquids, and gases, the ideal gas law and kinetic molecular theory, solutions and molarity, separation techniques like chromatography and distillation, and spectroscopy with the Beer-Lambert law. At 18-22% of the AP exam, it ties Units 1 and 2 (structure) to everything that comes after (behavior), and it carries the second-highest weight of any unit.

What this unit covers

Intermolecular forces and what they predict

• London dispersion forces (LDFs) come from temporary, fluctuating dipoles. Every molecule has them, and they get stronger with more electrons, larger electron clouds, and more contact area between molecules. For big molecules, LDFs are often the strongest net force, which is why I2 is a solid while Cl2 is a gas.
• Dipole-dipole forces exist between polar molecules, where the partially positive end of one molecule attracts the partially negative end of another.
• Hydrogen bonding is an especially strong dipole-dipole interaction that requires H bonded directly to N, O, or F. It explains why water boils at 100 °C while the similar-sized CH4 boils at -161 °C.
• Ion-dipole forces form between ions and polar molecules, like Na+ surrounded by water molecules. These are central to why ionic compounds dissolve in water.
• Stronger IMFs mean lower vapor pressure and higher boiling point, because vaporizing a liquid means completely overcoming those attractions. Melting point trends are similar but subtler, since melting only rearranges interactions rather than breaking them all.

States of matter at the particle level

• Solids can be crystalline (particles in a regular, repeating 3D arrangement) or amorphous (no orderly pattern). Either way, particles vibrate in place and don't translate past each other.
• Liquid particles are in close contact but can slide past one another, so liquids flow and take the shape of their container while keeping a fixed volume.
• Gas particles are far apart, moving freely and randomly, with attractions weak enough to mostly ignore.
• The macroscopic property always traces back to two things, how the particles are arranged and how strongly they attract each other. That's the lens for the whole unit.

Gases: the ideal model and where it breaks

• The ideal gas law, PV = nRT, links pressure, volume, moles, and temperature for a gas sample.
• In a mixture, each gas exerts its partial pressure independently. Partial pressure is proportional to mole fraction (P_A = P_total × X_A), and the partial pressures add up to the total.
• Kinetic molecular theory (KMT) explains gas behavior through particle motion. Particles move continuously and randomly, and average kinetic energy depends only on temperature. The Maxwell-Boltzmann distribution shows the spread of particle speeds at a given temperature. Heat the sample and the curve flattens and shifts toward higher speeds.
• At the same temperature, lighter gases move faster on average than heavier ones, since KE = (1/2)mv² and the average KE is equal.
• Real gases deviate from ideal behavior in two situations. Near condensation conditions (high pressure, low temperature), attractions between particles make the measured pressure lower than predicted. At extremely high pressures, the particles' own volume matters and the gas takes up more space than predicted.

Solutions, mixtures, and separations

• A solution is a homogeneous mixture, meaning its properties are the same throughout. A heterogeneous mixture varies depending on where you sample it. Solutions can be solid, liquid, or gas.
• Molarity, M = moles of solute / liters of solution, is the standard lab measure of concentration. You'll use it constantly for dilution and stoichiometry problems.
• Particulate diagrams of solutions show both relative concentrations and the interactions between components, like water molecules oriented with their oxygen ends toward a Na+ ion.
• "Like dissolves like" is the solubility rule with a real mechanism behind it. Substances with similar intermolecular interactions are miscible because the new solute-solvent attractions can replace the old solute-solute and solvent-solvent attractions.
• Filtration can't separate a true solution. Instead, you exploit differences in IMFs. Chromatography (paper, thin-layer, column) separates components by how strongly they stick to the stationary phase versus the mobile phase. Distillation separates by differences in volatility, which trace back to IMF strength.
• Colligative properties, molality, and percent composition calculations are not tested on the AP exam.

Light, spectroscopy, and Beer-Lambert

• Different regions of the electromagnetic spectrum match different molecular changes. Microwave radiation causes rotational transitions, infrared causes vibrational transitions, and UV/visible light causes electronic transitions (electrons jumping energy levels).
• A photon's energy is quantized. When absorbed, it raises the atom or molecule's energy by exactly E = hv, and c = λv connects wavelength and frequency. Short wavelength means high frequency means high energy.
• The Beer-Lambert law, A = εbc, says absorbance depends on molar absorptivity (how intensely the species absorbs at that wavelength), path length, and concentration. With path length and wavelength fixed, absorbance is directly proportional to concentration, which is why a calibration curve of A versus c is a straight line through the origin.

Unit 3, Properties of Substances and Mixtures at a glance

Why Unit 3, Properties of Substances and Mixtures matters in AP Chem

This unit is where the course's central habit of mind gets locked in, connecting what you can see and measure (boiling points, pressures, colors of solutions) to what particles are doing at a level you can't see. It's the payoff for the structure work in Units 1 and 2, and it builds the explanatory toolkit you'll use for the rest of the year.

• IMF reasoning is the single most reusable skill in AP Chem. It shows up again in dissolution, evaporation, reaction energetics, and lab-based questions about separations.
• Particulate models are a tested skill in their own right. Drawing or interpreting particle diagrams of solutions, phases, and gas mixtures appears across both multiple choice and free response.
• Beer-Lambert and molarity are the workhorse quantitative tools of the AP Chem lab program, from titrations to kinetics experiments that track concentration with a spectrophotometer.

How this unit connects across the course

• Backward to Compound Structure and Properties (Unit 2): you can't predict IMFs without polarity, and you can't get polarity without Lewis structures, molecular geometry, and electronegativity. Unit 3 is Unit 2's structure work put to use.
• Backward to Atomic Structure and Properties (Unit 1): photon energy, electronic transitions, and the electromagnetic spectrum extend the photoelectron spectroscopy and electron-energy ideas from Unit 1 into molecules.
• Forward to Chemical Reactions (Unit 4) and Kinetics (Unit 5): molarity and solution chemistry feed directly into solution stoichiometry, net ionic equations, and rate experiments that use Beer-Lambert absorbance data to track concentration over time.
• Forward to Thermochemistry (Unit 6) and Thermodynamics (Unit 9): the energy needed to overcome IMFs explains enthalpies of vaporization and dissolution, and the particle-motion picture from KMT underlies entropy reasoning later.

Key equations and processes

• PV = nRT, the ideal gas law. Use it to relate pressure, volume, moles, and temperature for any gas sample (R = 0.08206 L·atm/mol·K with atm and liters).
• P_A = P_total × X_A, Dalton's law via mole fraction. Use it to find the partial pressure of one gas in a mixture.
• KE_avg depends only on temperature. At the same T, all gases have the same average kinetic energy, so lighter molecules move faster.
• M = n_solute / L_solution, molarity. The standard concentration measure for dilution and solution stoichiometry.
• c = λv, relating wavelength and frequency of light (c = 3.00 × 10⁸ m/s).
• E = hv, Planck's equation for photon energy (h = 6.626 × 10⁻³⁴ J·s). Combine with c = λv to get E = hc/λ.
• A = εbc, the Beer-Lambert law. With fixed wavelength and path length, absorbance is directly proportional to concentration, which is the basis of every calibration-curve problem.
• Chromatography and distillation as processes. Chromatography separates by differential attraction to a stationary phase; distillation separates by differences in volatility. Both work because the components have different IMFs.

Unit 3, Properties of Substances and Mixtures on the AP exam

Unit 3 is 18-22% of the exam, tied for the largest share of any unit, so expect it everywhere. On the multiple-choice section, you'll rank boiling points or vapor pressures from molecular structures, interpret Maxwell-Boltzmann distributions, identify which particulate diagram correctly shows a dissolved ionic compound, and solve ideal gas law and partial pressure problems. Free-response questions love this unit's lab side. A typical prompt gives you a chromatography or distillation setup and asks you to explain the separation in terms of intermolecular forces, or hands you spectrophotometry data and a calibration curve and asks you to determine a concentration with Beer-Lambert. Justification matters as much as the answer. "Stronger IMFs" earns nothing by itself; you need to name the specific force, tie it to the structure (more electrons, H bonded to O, larger dipole), and connect it to the property. Drawing particulate representations is also fair game, so practice sketching water molecules oriented correctly around ions.

Essential questions

• Why do substances with nearly identical molar masses have wildly different boiling points?
• How does the motion and arrangement of invisible particles produce the properties of matter you can actually measure?
• Why does "like dissolves like" work, and how can differences in attraction be used to pull a mixture apart?
• How can shining light through a solution tell you exactly how concentrated it is?

Key terms to know

• London dispersion forces: Attractions caused by temporary, fluctuating dipoles, present in all molecules and strongest in large, polarizable ones.
• Polarizability: How easily a molecule's electron cloud distorts to form a temporary dipole; it increases with more electrons and a larger cloud.
• Hydrogen bonding: A strong dipole-dipole attraction requiring hydrogen bonded directly to N, O, or F.
• Vapor pressure: The pressure of a vapor in equilibrium with its liquid; lower vapor pressure means stronger IMFs.
• Crystalline vs. amorphous solid: A crystalline solid has a regular, repeating particle arrangement; an amorphous solid does not.
• Kinetic molecular theory: The model linking gas properties to continuous, random particle motion, with average kinetic energy set by temperature alone.
• Maxwell-Boltzmann distribution: A graph of how particle kinetic energies (or speeds) are spread out at a given temperature.
• Partial pressure: The pressure one gas in a mixture exerts on its own, proportional to its mole fraction.
• Molarity: Moles of solute per liter of solution, the standard lab concentration unit.
• Chromatography: A separation method based on how strongly components stick to a stationary phase versus a mobile phase.
• Distillation: A separation method that exploits differences in volatility, which depend on IMF strength.
• Molar absorptivity (ε): A constant describing how intensely a species absorbs light at a specific wavelength in the Beer-Lambert law.
• Electronic transition: An electron jumping between energy levels, associated with UV/visible light absorption or emission.

Common mix-ups

• Intermolecular vs. intramolecular. Boiling water breaks hydrogen bonds between molecules, not the O-H covalent bonds inside them. Saying a substance boils because "bonds break" loses points.
• Hydrogen bonding requires H directly bonded to N, O, or F. CH4 has hydrogens but no hydrogen bonding, and HCl is polar but does not hydrogen bond.
• Temperature sets average kinetic energy, not speed. Two different gases at the same temperature have equal average KE, but the lighter gas moves faster.
• Real gases deviate most at low temperature and high pressure. Low T lets attractions matter (pressure drops below ideal), and very high P makes particle volume matter (volume exceeds ideal). High temperature and low pressure are the most ideal conditions, not the least.
• LDFs are not always the weakest force in practice. Between large molecules, total dispersion forces can outweigh dipole-dipole attractions, so compare actual structures rather than reciting the generic ranking.