Intermolecular Forces & Properties: every key term you need (+ practice quiz)
59 flashcard terms for AP Chemistry Unit 3, written to match the course framework. Study them here, then drill them as interactive flashcards, or test yourself with the 20-question quiz — free, no account needed.
Attractive forces BETWEEN molecules (not the covalent bonds within them). They are much weaker than bonds but determine physical properties like boiling point.
IMFs vs. Bonds
Breaking a covalent bond (intramolecular) requires far more energy than overcoming intermolecular forces — melting and boiling only disrupt IMFs, not bonds.
London Dispersion Forces (LDFs)
Weak, temporary attractions from momentary uneven electron distribution creating instantaneous dipoles. Present in ALL molecules, polar or not.
Instantaneous Dipole
A brief, random shift of electron density in a molecule that creates a temporary partial charge, inducing a dipole in a neighboring molecule.
Polarizability
How easily an atom's or molecule's electron cloud is distorted into a dipole. Larger, more electron-rich (heavier) species are more polarizable.
LDF Strength Trend
Dispersion forces increase with molecular size/mass (more electrons, more polarizable) and with more surface area for contact (e.g., straight-chain vs. branched isomers).
Dipole-Dipole Forces
Attractions between the positive end of one polar molecule and the negative end of another. Present only in polar molecules, in addition to LDFs.
Dipole-Dipole vs. LDF Strength
For molecules of similar size, dipole-dipole forces are generally stronger than LDFs alone, giving polar molecules higher boiling points than similarly sized nonpolar ones.
Hydrogen Bonding
An especially strong dipole-dipole attraction when H is bonded directly to N, O, or F — small, highly electronegative atoms that create a very concentrated partial charge.
Why Hydrogen Bonds Are Strong
N, O, and F are small and highly electronegative, concentrating a large partial negative charge that strongly attracts the exposed, nearly bare H nucleus of another molecule.
Hydrogen Bonding Requirement
Requires H bonded directly to N, O, or F in one molecule, attracted to a lone pair on N, O, or F in another — not just any molecule containing these atoms.
Evidence for Hydrogen Bonding
Anomalously high boiling points of NH₃, H₂O, and HF compared to other hydrides in their periodic groups reveal hydrogen bonding's extra strength.
Ion-Dipole Forces
An attraction between an ion and the oppositely charged end of a polar molecule — the strongest of the intermolecular-type forces, responsible for ionic compounds dissolving in water.
Hydration Shell
The shell of oriented water molecules (negative O toward cations, positive H toward anions) surrounding a dissolved ion via ion-dipole forces.
IMF Strength Ranking
From weakest to strongest (typical case): London dispersion < dipole-dipole < hydrogen bonding < ion-dipole — though strong LDFs in very large molecules can exceed weak dipole-dipole forces.
Boiling Point and IMF Strength
Stronger IMFs require more energy to separate molecules into the gas phase, so substances with stronger IMFs have higher boiling and melting points.
Vapor Pressure
The pressure exerted by a substance's vapor in equilibrium with its liquid. Weaker IMFs mean molecules escape the liquid more easily, giving higher vapor pressure.
Volatility
A substance's tendency to vaporize. Weak IMFs → high volatility and high vapor pressure; strong IMFs → low volatility and low vapor pressure.
Viscosity
A liquid's resistance to flow. Stronger IMFs (and larger, more entangled molecules) increase viscosity by resisting molecules sliding past one another.
Surface Tension
The energy required to increase a liquid's surface area, caused by unbalanced IMF attraction pulling surface molecules inward. Stronger IMFs mean higher surface tension.
Solid
Particles are packed in fixed positions with strong enough attractions to resist motion; definite shape and volume.
Liquid
Particles are close together but free to move past one another; definite volume, but takes the shape of its container.
Gas
Particles are far apart with negligible attraction, moving independently; no definite shape or volume, fills its container.
Phase Changes & Energy
Melting, vaporizing, and sublimating require energy input to overcome IMFs; freezing, condensing, and depositing release energy as IMFs re-form.
Heating Curve
A graph of temperature vs. heat added, showing temperature rising within a phase, then plateauing during a phase change (energy overcomes IMFs without raising temperature).
Heat of Fusion
The energy required to melt one mole (or gram) of a solid into a liquid at its melting point, without a temperature change.
Heat of Vaporization
The energy required to vaporize one mole (or gram) of a liquid into gas at its boiling point — always larger than heat of fusion since vaporization fully separates particles.
Solid Types — Ionic
A lattice of cations and anions held by strong electrostatic (Coulombic) attractions; hard, brittle, high melting point, conducts only when molten/dissolved.
Solid Types — Metallic
Cations in a sea of delocalized electrons; conducts electricity as a solid, malleable and ductile, wide range of melting points.
Solid Types — Covalent Network
Atoms bonded covalently in a continuous 3-D lattice (diamond, quartz, SiO₂); extremely hard, very high melting point, generally poor conductors.
Solid Types — Molecular
Discrete molecules held together by IMFs (LDFs, dipole-dipole, hydrogen bonds); relatively low melting points, often soft, poor conductors.
A solid with a highly ordered, repeating particle arrangement (e.g., salt, diamond, most metals).
Amorphous Solid
A solid lacking long-range order in particle arrangement (e.g., glass, rubber); often softens gradually rather than melting sharply.
Solutions
Homogeneous mixtures where a solute is dispersed uniformly in a solvent, held together by intermolecular attractions between solute and solvent particles.
'Like Dissolves Like'
Substances with similar polarity and IMF types tend to be miscible/soluble: polar dissolves polar, nonpolar dissolves nonpolar.
Why Ionic Compounds Dissolve in Water
Water's polar molecules surround and stabilize ions via strong ion-dipole forces, releasing energy that offsets breaking the ionic lattice.
Why Oil and Water Don't Mix
Nonpolar oil molecules can only form weak LDFs, which cannot compete with water's strong hydrogen bonding network — water 'excludes' the oil.
Solubility and Temperature (Solids)
Most solid solutes become more soluble in water as temperature increases, since added energy helps overcome the solute's lattice/IMF forces.
Solubility and Temperature (Gases)
Gas solubility in water typically decreases as temperature increases, since more kinetic energy lets gas molecules escape the solvent more easily.
Henry's Law
The solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid.
Concentration — Molarity
Moles of solute per liter of solution (mol/L); the standard way to express solution concentration in chemistry.
Dilution Equation
M₁V₁ = M₂V₂ — relates the concentration and volume before and after adding solvent, since moles of solute stay constant.
Colligative Properties
Physical properties of a solution that depend on the NUMBER of dissolved particles, not their identity: boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure.
Boiling Point Elevation
Dissolved solute particles disrupt the solvent's surface, lowering vapor pressure and requiring a higher temperature to boil than the pure solvent.
Freezing Point Depression
Dissolved solute particles disrupt the solvent's ability to form an ordered solid lattice, so the solution freezes at a lower temperature than the pure solvent.
Van't Hoff Factor (i)
The number of particles a solute produces in solution (e.g., NaCl → i=2, glucose → i=1). Colligative effects scale with the TOTAL particle concentration (i × molality).
Why Salt Melts Ice
Dissolved salt ions depress the freezing point of water below 0°C, so ice melts even at temperatures where pure water would stay frozen.
Vapor Pressure Lowering
Nonvolatile solute particles occupy some of the liquid's surface, reducing the rate of solvent evaporation and thus the equilibrium vapor pressure.
Osmotic Pressure
The pressure needed to stop osmosis (net water movement) across a semipermeable membrane; increases with solute particle concentration.
Kinetic Molecular Theory (Gases)
Gas particles are in constant, random motion; have negligible volume compared to container size; and experience negligible IMFs between collisions (ideal gas assumptions).
Ideal Gas Law
PV = nRT — relates pressure, volume, moles, and temperature of a gas, assuming negligible particle volume and no intermolecular attractions.
Real Gases Deviate from Ideal
At high pressure (particles forced close) or low temperature (slower particles), real IMFs and particle volume become significant, causing deviation from ideal behavior.
Gas Behaves Most Ideally When...
Pressure is low and temperature is high — particles are far apart (volume negligible) and moving fast (IMFs negligible relative to kinetic energy).
Diffusion and Effusion
Diffusion is gas mixing through random motion; effusion is gas escaping through a tiny hole. Both depend on molar mass — lighter gases move/effuse faster (Graham's Law).
Graham's Law
The rate of effusion is inversely proportional to the square root of molar mass — lighter gas particles move faster and effuse more quickly.
Structure-Property Theme (IMFs)
A molecule's size, shape, and polarity determine which IMFs it experiences, which in turn determine its physical properties — the recurring 'structure determines function/properties' logic.
Comparing Boiling Points — Worked Logic
To compare boiling points: identify the strongest IMF each substance can have (H-bonding > dipole-dipole > LDF only), then break ties using size/mass for stronger LDFs.
Branching and Boiling Point
Among isomers, more branching decreases surface area contact between molecules, weakening LDFs and lowering the boiling point compared to straight-chain isomers.