Chapter 10: Intermolecular Forces, Liquids, and Solids

30 vocabulary flashcards covering key terms from Chapter 10 on intermolecular forces, properties of liquids, phase transitions, and solid-state concepts.

Intermolecular Forces (IMFs)

Non-covalent electrostatic attractions between separate molecules or atoms that govern many physical properties of liquids and solids.

Intramolecular Forces

Chemical bonds (covalent, ionic, metallic) that hold atoms together within a single molecule or formula unit.

Dispersion Forces (London Dispersion)

Weak, temporary dipole-induced dipole attractions present in all substances, arising from momentary electron fluctuations.

Dipole–Dipole Attractions

Forces between permanent molecular dipoles; occur only in polar molecules and are stronger than dispersion forces.

Hydrogen Bonding

The strongest van der Waals force; a special dipole–dipole attraction involving H covalently bonded to N, O, or F and a nearby lone pair on N, O, or F.

Van der Waals Forces

Collective term for dispersion forces, dipole–dipole attractions, and hydrogen bonding.

Polarizability

A measure of how easily a molecule’s electron cloud can be distorted to produce an induced dipole.

Kinetic Energy (KE) vs. IMFs

Thermal energy that can overcome intermolecular attractions; higher KE (temperature) promotes phase changes.

Adhesive Forces

Attractions between a liquid’s molecules and a different surface (e.g., water to glass).

Cohesive Forces

Attractions between identical molecules within a liquid (e.g., water to water).

Viscosity

A liquid’s resistance to flow; increases with stronger IMFs and decreases with higher temperature.

Surface Tension

Energy required to increase a liquid’s surface area; results in minimized surface shapes such as spherical drops.

Capillary Action

Rise or fall of a liquid in a narrow tube due to the competition between adhesive and cohesive forces.

Vaporization

Phase transition in which molecules escape from the liquid to the gas phase; also called evaporation.

Condensation

Phase transition in which gas molecules return to the liquid phase.

Dynamic Equilibrium (Liquid–Vapor)

State in a closed system where the rates of vaporization and condensation are equal.

Vapor Pressure

Pressure exerted by a vapor in equilibrium with its liquid (or solid) at a given temperature.

Boiling Point

Temperature at which a liquid’s vapor pressure equals the surrounding external pressure.

Normal Boiling Point

Boiling temperature of a liquid at 1 atm (101.3 kPa) external pressure.

Clausius–Clapeyron Equation

Mathematical relation that links a substance’s vapor pressure to temperature and enthalpy of vaporization.

Sublimation

Direct phase transition from solid to gas without passing through the liquid phase.

Deposition

Direct phase transition from gas to solid, opposite of sublimation.

Triple Point

Unique temperature and pressure at which solid, liquid, and gas phases coexist in equilibrium.

Critical Point

Temperature and pressure above which the distinction between liquid and gas disappears.

Supercritical Fluid

Single phase existing above the critical point that exhibits properties of both liquids and gases.

Phase Diagram

Graph of pressure vs. temperature showing the stability regions of solid, liquid, gas, and supercritical phases.

Enthalpy of Fusion (ΔHfus)

Heat required to melt one mole of a substance at its melting point.

Enthalpy of Vaporization (ΔHvap)

Heat required to vaporize one mole of a substance at its boiling point.

Enthalpy of Sublimation (ΔHsub)

Heat required to convert one mole of solid directly to gas; approximately equal to ΔHfus + ΔHvap.

Polar Molecule

Molecule with an uneven charge distribution resulting in a permanent dipole moment.