Chapter 10: Intermolecular Forces, Liquids, and Solids

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

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30 Terms

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Intermolecular Forces (IMFs)

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

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Intramolecular Forces

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

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Dispersion Forces (London Dispersion)

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

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Dipole–Dipole Attractions

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

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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.

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Van der Waals Forces

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

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Polarizability

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

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Kinetic Energy (KE) vs. IMFs

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

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Adhesive Forces

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

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Cohesive Forces

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

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Viscosity

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

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Surface Tension

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

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Capillary Action

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

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Vaporization

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

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Condensation

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

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Dynamic Equilibrium (Liquid–Vapor)

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

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Vapor Pressure

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

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Boiling Point

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

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Normal Boiling Point

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

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Clausius–Clapeyron Equation

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

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Sublimation

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

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Deposition

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

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Triple Point

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

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Critical Point

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

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Supercritical Fluid

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

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Phase Diagram

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

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Enthalpy of Fusion (ΔHfus)

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

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Enthalpy of Vaporization (ΔHvap)

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

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Enthalpy of Sublimation (ΔHsub)

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

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Polar Molecule

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