AP Chemistry Unit 3 Cheat Sheet

The basics

What it covers: The largest unit on the exam — intermolecular forces, states of matter, gas laws, solutions, solubility, and spectroscopy.

Exam weight: About 18–22% of the AP Chemistry exam — nearly a quarter of all questions.

The big question: How does the strength and type of intermolecular force between particles determine the physical properties and behavior of a substance?

Big Ideas covered: Structure & Properties (SAP), Scale Proportion & Quantity (SPQ), Transformations (TRA), Chemical Effects (CE).

The key terms you must know

• London dispersion force — temporary dipole attraction present in ALL molecules; the only IMF in nonpolar substances.
• Dipole-dipole force / hydrogen bonding — stronger IMFs found only in polar molecules (hydrogen bonding requires N–H, O–H, or F–H).
• Vapor pressure, viscosity, surface tension — macroscopic properties directly explained by IMF strength.
• Ideal gas law (PV = nRT) — the core equation relating pressure, volume, moles, and temperature.
• Kinetic molecular theory — the model behind ideal gas behavior; average KE is proportional to temperature in Kelvin.
• Real gas deviation — occurs at high pressure (volume) and low temperature (IMFs).
• Molarity — mol solute per liter of solution; the standard concentration unit.
• "Like dissolves like" — polar dissolves polar, nonpolar dissolves nonpolar.
• Beer-Lambert law (A = εlc) — absorbance is proportional to concentration.
• Phase diagram — pressure vs. temperature graph showing solid/liquid/gas regions, triple point, and critical point.

Key themes to remember

• IMFs are everything in this unit. Nearly every property — boiling point, viscosity, solubility, vapor pressure — traces back to IMF type and strength.
• The structure-property chain runs from Unit 2 into Unit 3. Molecular shape and polarity (Unit 2) determine which IMFs apply (Unit 3), which determine physical properties.
• "Ideal" gas behavior is a useful fiction. Real gases only deviate when molecular volume or IMFs actually start to matter — high pressure and low temperature.
• Dissolving is a competition between IMFs. A solute dissolves only if new solute-solvent IMFs are strong enough to replace the ones being broken.
• Spectroscopy connects light to concentration. The Beer-Lambert law is the mathematical bridge between an absorbance reading and an actual concentration value.

Common exam traps

• Every molecule has London dispersion forces — even polar ones. Don't say a polar molecule has "no London dispersion forces" just because it also has dipole-dipole forces.
• Hydrogen bonding requires a specific setup. It's not "any molecule with hydrogen" — it requires H bonded directly to N, O, or F, attracted to a lone pair on a neighboring N, O, or F.
• Higher molar mass usually (not always) means stronger London dispersion. Molecular shape and surface area matter too — long, skinny molecules have more surface contact than compact, round ones of similar mass.
• Real gas deviation isn't random — name the cause. High pressure → molecular volume becomes significant. Low temperature → intermolecular attraction becomes significant. Always specify which.
• Don't confuse molarity with molality or mass percent. Molarity is specifically moles of solute per liter of solution — know the units cold.