Unit 3: Intermolecular Forces and Properties

Chemical bond

An intramolecular attraction that holds atoms together within a particle (e.g., within a molecule).

Intermolecular force (IMF)

An attraction between separate particles (e.g., molecule–molecule or ion–molecule) that largely controls properties like boiling point, vapor pressure, viscosity, and miscibility.

Polarity

The way charge is distributed within a bond or molecule; uneven charge distribution creates partial positive and partial negative regions.

Electron density

How electron “cloud” is distributed around atoms; uneven electron density can create partial charges and dipoles.

Electronegativity

A measure of how strongly an atom attracts shared electrons in a covalent bond; differences help predict bond polarity.

Polar covalent bond

A covalent bond with unequal electron sharing due to electronegativity differences, producing partial charges on the bonded atoms.

Partial charge (δ+ / δ−)

A small, non-integer charge on an atom in a polar bond; written using delta symbols (δ+ and δ−).

Dipole

Opposite charges separated by a distance (a positive end and a negative end) within a bond or molecule.

Dipole moment

A measure of the degree of charge separation in a bond or molecule (how strong the dipole is).

Nonpolar covalent bond

A covalent bond with essentially equal electron sharing (very similar electronegativities), producing little to no dipole.

Molecular polarity

Whether a whole molecule has a net dipole; depends on both bond polarities and molecular geometry (dipole cancellation or not).

Dipole cancellation

When individual bond dipoles in a molecule add to zero because of symmetry/geometry, making the molecule nonpolar overall despite polar bonds.

VSEPR (Valence Shell Electron Pair Repulsion)

A model used to predict molecular shape from electron group repulsions, helping determine whether bond dipoles cancel.

Symmetry shortcut (polarity)

Heuristic: if the central atom has no lone pairs and all surrounding atoms are identical, the molecule is often symmetric and nonpolar; lone pairs and/or different atoms often make it polar.

London dispersion forces (LDF)

Attractions caused by temporary fluctuations in electron density that create instantaneous dipoles and induce dipoles in nearby particles; present in all substances.

Instantaneous dipole

A momentary, uneven electron distribution in a particle that briefly creates a dipole, initiating dispersion attractions.

Induced dipole

A dipole created in a neighboring particle when an instantaneous dipole distorts its electron cloud.

Polarizability

How easily an electron cloud can be distorted; increases with more electrons/larger molar mass, strengthening dispersion forces.

Surface area of contact (IMFs)

The extent of particle-to-particle contact; greater surface area (often less branching/more linear shape) strengthens London dispersion forces.

Dipole–dipole forces

Attractions between polar molecules where the positive end of one molecule is attracted to the negative end of another.

Hydrogen bonding

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

Hydrogen-bond donor

A molecule/group that provides the H involved in hydrogen bonding (an H covalently bonded to N, O, or F).

Hydrogen-bond acceptor

An atom with a lone pair (typically N, O, or F) that attracts the H from a donor in hydrogen bonding.

Ion–dipole forces

Attractions between an ion and a polar molecule (e.g., Na+ with the oxygen end of water), important in dissolving ionic compounds in polar solvents.

IMF strength hierarchy (AP-style)

Common simplified ranking: ion–dipole (often strongest in solutions) > hydrogen bonding > dipole–dipole > London dispersion (always present; can be strong for large, polarizable molecules).

Vapor pressure

The pressure exerted by vapor above a liquid in a closed container at a given temperature; stronger IMFs give lower vapor pressure.

Boiling point

The temperature at which a liquid’s particles have enough energy to enter the gas phase (bubble formation throughout); stronger IMFs generally raise boiling point.

Melting point

The temperature at which a solid becomes a liquid; depends on both attraction strength and how efficiently particles pack in the solid lattice.

Enthalpy of vaporization (ΔHvap)

Energy required to vaporize a given amount of liquid at its boiling point; larger for substances with stronger IMFs.

Enthalpy of fusion (ΔHfus)

Energy required to melt a given amount of solid at its melting point; increases with stronger attractions holding the solid together.

Viscosity

A liquid’s resistance to flow; stronger intermolecular attractions generally increase viscosity.

Surface tension

The tendency of a liquid surface to minimize area due to cohesive forces pulling surface molecules inward; stronger IMFs increase surface tension.

Heating curve

A graph of temperature vs. heat added showing warming within phases (sloped regions) and phase changes (flat plateaus).

Phase-change plateau

The flat part of a heating curve where temperature stays constant because added energy is used to overcome intermolecular attractions during melting or boiling.

Phase diagram

A graph showing which phase (solid, liquid, gas) is stable at different temperatures and pressures, including boundaries where phases coexist.

Triple point

The point on a phase diagram where solid, liquid, and gas all coexist at equilibrium.

Critical point

The point on a phase diagram beyond which liquid and gas become indistinguishable (a supercritical fluid forms).

Molecular solid

A solid made of neutral molecules held together by IMFs (dispersion, dipole–dipole, and/or hydrogen bonding); typically low melting and nonconductive.

Ionic solid

A crystal lattice of cations and anions held by electrostatic attraction (ionic bonding); high melting, brittle, conducts only when molten or dissolved.

Metallic solid

A solid of metal atoms with delocalized valence electrons (“sea of electrons”); conducts well and is malleable/ductile.

Covalent network solid

A continuous network of atoms bonded by covalent bonds (not discrete molecules); very high melting points and often very hard (graphite is an exception for conductivity).

Kinetic molecular theory (KMT)

Ideal-gas assumptions: negligible particle volume, constant random motion, elastic collisions, and no intermolecular attractions (except during collisions).

Ideal gas law

The relationship among pressure, volume, moles, and temperature for an ideal gas: PV = nRT (T in kelvin).

Gas constant (R)

The proportionality constant in PV = nRT; commonly 0.0821 L·atm/(mol·K) when using liters and atmospheres.

Combined gas law

For constant moles: (P1V1)/T1 = (P2V2)/T2, relating two states of the same gas sample.

Dalton’s law of partial pressures

In a gas mixture, total pressure equals the sum of each gas’s partial pressure: Ptotal = P1 + P2 + …

Mole fraction (Xi)

The fraction of total moles contributed by component i: Xi = ni/ntotal; for ideal gases, Pi = XiPtotal.

Effusion

The escape of gas through tiny holes from higher to lower pressure; faster at higher temperature and for lower molar mass gases (which move faster at the same T).

Beer’s Law

A relationship in spectroscopy linking absorbance to concentration: A = abc, where a is molar absorptivity, b is path length, and c is concentration.

Retention factor (Rf)

In chromatography, Rf = (distance traveled by solute)/(distance traveled by solvent front); larger Rf indicates greater affinity for the mobile phase relative to the stationary phase.