Intermolecular Forces and Properties

What this unit covers

The topics below follow the published Chemistry course framework for Unit 3. This unit is worth 18–22% of the exam, so budget your time against that rather than against how long the unit takes to teach.

Lessons in this unit

• Intermolecular Forces13 min · 3 objectivesDistinguish intermolecular forces from the covalent bonds within a molecule · Rank London dispersion, dipole–dipole, hydrogen bonding, and ion–dipole by strength · Predict relative IMF strength from polarity, polarizability, and molecular size
• Properties of Liquids & Solids · Phase Changes13 min · 3 objectivesConnect IMF strength to boiling point, vapor pressure, viscosity, and surface tension · Classify phase changes as endothermic or exothermic and read a phase diagram · Explain how temperature and pressure determine the state of a substance
• Gases & the Ideal Gas Law14 min · 3 objectivesState the assumptions of kinetic molecular theory and how they define an ideal gas · Apply PV = nRT to solve for any one variable, converting units to K and atm · Use mole fractions to find partial pressures in a gas mixture
• Solutions & Mixtures12 min · 3 objectivesCalculate molarity and use it to relate moles, mass, and solution volume · Predict solubility with the "like dissolves like" rule and identify ion–dipole solvation · Explain how chromatography and distillation separate the components of a mixture
• Gas Laws Deep Dive14 min · 3 objectivesUse Dalton’s law and mole fractions to find partial pressures, including gas collected over water · Apply Graham’s law to compare effusion and diffusion rates from molar masses · Explain when and why real gases deviate from ideal behavior at high pressure or low temperature
• Solutions & Spectroscopy14 min · 3 objectivesPrepare solutions and perform dilutions using M₁V₁ = M₂V₂ and molarity · Carry out solution stoichiometry, converting between volume, molarity, and moles in a reaction · Apply the Beer–Lambert law, A = εbc, to find a concentration from a measured absorbance

Formulas in Unit 3

Every term in Unit 3

All 22 terms we publish for Intermolecular Forces and Properties, with definitions. Reading them through is the fastest way to find the ones you cannot define — then drill those in cram mode until you can produce them without the prompt.

London dispersion forces

Temporary induced dipoles present in all molecules. Strength rises with the number of electrons and with surface contact — the reason larger nonpolar molecules boil higher.

Dipole-dipole forces

Attraction between permanent dipoles of polar molecules; stronger than dispersion for molecules of similar size.

Hydrogen bonding

A strong dipole-dipole interaction where H is bonded to N, O or F. It explains water's anomalously high boiling point.

Ion-dipole forces

Between an ion and a polar molecule; the interaction that dissolves ionic solids in water.

Boiling point and IMF strength

Boiling separates molecules, so it depends on intermolecular forces, not on bond strength within molecules.

Vapor pressure

Pressure of vapor above a liquid at equilibrium. Weak intermolecular forces mean high vapor pressure and easy evaporation.

Ideal gas law

PV = nRT with R = 0.08206 L·atm·mol⁻¹·K⁻¹. Temperature must be in kelvin.

Kinetic molecular theory assumptions

Negligible particle volume, no intermolecular forces, elastic collisions, and average kinetic energy proportional to absolute temperature.

Deviations from ideality

Real gases deviate at high pressure (particle volume matters) and low temperature (attractions matter). Deviation is largest for large polar molecules.

Partial pressure and Dalton's law

Total pressure is the sum of partial pressures; each gas's partial pressure is its mole fraction times the total.

Maxwell-Boltzmann distribution

The spread of molecular speeds at a temperature. Higher temperature broadens and shifts it right, putting more molecules above the activation energy.

Molarity

Moles of solute per liter of solution. Dilution follows M₁V₁ = M₂V₂ because moles of solute are conserved.

Beer-Lambert law

A = εbc — absorbance is proportional to concentration, which lets a calibration curve turn a color measurement into a concentration.

Chromatography

Separates components by their relative attraction to a mobile and a stationary phase; a component more attracted to the mobile phase travels further.

Distillation

Separates liquids by boiling point, which is separation by intermolecular force strength.

Why intermolecular forces are weaker than bonds

They are attractions between whole molecules rather than shared or transferred electrons, so boiling requires far less energy than decomposing.

Comparing boiling points correctly

Identify the strongest force present in each substance, then compare within the same type using size and polarisability. A larger nonpolar molecule can outboil a small polar one.

Surface tension and viscosity

Both increase with stronger intermolecular forces, since molecules resist being separated or moved past one another.

Capillary action

Adhesion to the tube walls competing with cohesion within the liquid. Water climbs glass; mercury does not.

Solubility and "like dissolves like"

A solute dissolves when solute-solvent interactions are comparable to those it gives up. Polar dissolves polar because both form strong ion-dipole or dipole-dipole interactions.

Real gas deviation direction

Attractions make the measured pressure lower than ideal; finite particle volume makes the measured volume larger. Which dominates depends on conditions.

Effusion and Graham's law

Lighter gases effuse faster, with rate inversely proportional to the square root of molar mass, since equal temperature means equal average kinetic energy.

What examiners penalize here

• On free-response, always name the IMF *and* the reason. "Water has hydrogen bonding, which is stronger than the dipole–dipole forces in H₂S, so water has the higher boiling point" earns the point; just saying "water boils higher" does not.
• Reading a phase diagram: moving right (heating) at constant pressure crosses from solid to liquid to gas; moving up (compressing) at constant temperature can push a gas into a liquid or solid. Locate the triple point and critical point first — they anchor the whole graph.
• When a problem gives grams but asks about a reaction or concentration, convert to moles first — moles are the currency of both molarity (mol ÷ L) and stoichiometry. Watch that volume is in liters and that "solution" volume, not solvent volume, goes in the denominator.
• For real-gas questions, tie the deviation to a *cause*: high pressure → molecular volume no longer negligible (real V larger); low temperature → intermolecular attractions reduce wall impacts (real P smaller). Naming the mechanism, not just "it deviates," is what earns the point.
• For a Beer’s-law free-response, expect a calibration graph of absorbance vs concentration: its slope is εb, so a steeper line means a larger molar absorptivity. Read an unknown by finding its absorbance on the line and dropping to the concentration axis — and always confirm the path length is 1.00 cm before using ε directly.

Practice Chemistry

Our practice bank is drawn from across the whole course rather than filtered to one unit, which is closer to how the exam asks anyway — it will not tell you which unit a question is testing.

Questions about this unit

All 9 units of AP Chemistry

1. Unit 1 · Atomic Structure & Properties
2. Unit 2 · Molecular and Ionic Compound Structure and Properties
3. Unit 3 · Intermolecular Forces and Properties
4. Unit 4 · Chemical Reactions
5. Unit 5 · Kinetics
6. Unit 6 · Thermodynamics
7. Unit 7 · Equilibrium
8. Unit 8 · Acids & Bases
9. Unit 9 · Applications of Thermodynamics