AP Chemistry Unit 3 Learning Notes: How Forces Between Particles Control Phases and Material Properties

Intermolecular forces (IMFs)

Attractive forces between separate particles (molecules, atoms, or ions) that largely determine bulk properties like boiling point, melting point, viscosity, vapor pressure, and surface tension.

Intramolecular bonds

Bonds within a particle (e.g., covalent or ionic bonds) that hold atoms together inside a molecule or formula unit; not the same as IMFs.

London dispersion forces (LDF)

Attractions caused by temporary, fluctuating electron distributions that create instantaneous dipoles and induced dipoles; present between all particles and often dominate in nonpolar substances.

Polarizability

How easily an electron cloud can be distorted to form an induced dipole; generally increases with more electrons, larger size, and greater surface area of contact (less branching).

Instantaneous dipole

A momentary separation of charge due to random electron motion, which can induce a dipole in a neighboring particle and produce dispersion attractions.

Dipole–dipole forces

Attractions between permanent dipoles of polar molecules, where the partially positive end of one molecule aligns near the partially negative end of another.

Hydrogen bonding

A strong type of dipole–dipole attraction occurring when H is covalently bonded to N, O, or F and is attracted to a lone pair on N, O, or F of a nearby particle.

Ion–dipole forces

Attractions between an ion (full charge) and the partial charges on a polar molecule; important in dissolving ionic compounds in polar solvents like water.

Vapor pressure

The pressure of vapor above a liquid at a given temperature, reflecting how many particles escape from the liquid to the gas phase; stronger IMFs lower vapor pressure.

Volatility

A measure of how readily a substance vaporizes; higher volatility corresponds to higher vapor pressure and typically weaker IMFs.

Boiling point

The temperature at which a liquid’s vapor pressure equals the external pressure; stronger IMFs generally require a higher temperature to boil.

Viscosity

A liquid’s resistance to flow; tends to increase with stronger attractions and with long molecules that can tangle, and usually decreases as temperature increases.

Surface tension

The tendency of a liquid surface to resist being stretched, caused by a net inward attraction on surface molecules; stronger IMFs increase surface tension.

Crystalline solid

A solid with particles arranged in a repeating, ordered pattern (long-range order), typically showing a sharp melting point.

Amorphous solid

A solid lacking long-range order; often softens over a range of temperatures (e.g., glass, many plastics) rather than melting sharply.

Molecular solid

A crystalline solid made of neutral molecules at lattice points held together by IMFs (LDF, dipole–dipole, hydrogen bonding); typically low melting point and poor conductivity.

Ionic solid

A lattice of cations and anions held by electrostatic attraction; typically high melting point, brittle, and nonconducting as a solid but conducting when molten or dissolved.

Metallic solid

A solid of metal atoms/cations with delocalized, mobile valence electrons (“sea of electrons”); good conductor, malleable, and ductile with variable melting points.

Covalent network solid

A solid where atoms are connected in a continuous covalent-bond network; very high melting point and hard/rigid, usually poor conductors (graphite is an exception).

Heating curve

A graph of temperature vs heat added; sloped segments show temperature (kinetic energy) increasing within a phase, and flat segments show phase changes at constant temperature.

Phase change

A physical change between states (solid, liquid, gas) where energy is used mainly to change potential energy by overcoming attractions; temperature can remain constant during the change.

Triple point

The unique temperature and pressure at which solid, liquid, and gas phases coexist in equilibrium on a phase diagram.

Critical point

The temperature and pressure above which liquid and gas become indistinguishable (supercritical fluid region exists beyond this point).

Kinetic molecular theory (KMT)

A model for gases stating particles are far apart, move randomly with elastic collisions, and have average kinetic energy proportional to absolute temperature.

Ideal gas law

The relationship PV = nRT (with T in kelvins) describing ideal gas behavior, which assumes negligible IMFs and negligible particle volume.