General Chemistry 2 Notes: Intermolecular Forces and Phase Changes

Vocabulary flashcards covering key terms from the notes on states of matter, intermolecular and intramolecular forces, phase changes, and related properties.

Kinetic Molecular Theory (KMT)

Explains the behavior of matter based on the energy and movement of its particles; stronger intermolecular forces limit particle movement.

States of Matter

Gas, liquid, and solid; distinguished by particle arrangement, movement, and density.

Intermolecular Forces (IMF)

Attractions between molecules; weaker than chemical bonds; govern properties like boiling and melting points.

Intramolecular Forces

Forces within a molecule that hold atoms together (chemical bonds); typically stronger than IMF.

Ion-Dipole Forces

Attractions between an ion and a polar molecule; strongest IMF listed; e.g., Na+ with water.

Dipole-Dipole Forces

Attractions between two polar molecules with permanent dipoles; alignment of opposite poles.

London Dispersion Forces

Weakest IMF present in all molecules; arise from momentary dipoles; stronger with higher molar mass; important in nonpolar substances.

Hydrogen Bond

Strong dipole-dipole interaction when hydrogen is bonded to N, O, or F; contributes to water’s properties like high boiling point.

Ionic Bonding

Electrostatic attraction between oppositely charged ions formed by electron transfer.

Covalent Bonding

Bond formed by the sharing of electrons between atoms.

Metallic Bonding

Bonding in metals with a sea of delocalized electrons surrounding cations; explains conductivity and malleability.

Viscosity

Measure of a liquid’s resistance to flow; higher with stronger IMF.

Surface Tension

Tension at a liquid’s surface due to cohesive forces; water has high surface tension because of hydrogen bonding.

Capillarity

Rise of a liquid in a narrow tube due to cohesive and adhesive forces.

Evaporation

Process where surface molecules gain enough energy to escape into the gas phase; occurs below boiling point.

Vapor Pressure

Pressure exerted by the vapor in equilibrium with a liquid; higher with weaker IMF or higher temperature.

Boiling Point

Temperature at which vapor pressure equals atmospheric pressure; stronger IMF → higher boiling point.

Critical Temperature (Tc)

Highest temperature at which a substance can exist as a liquid; above Tc, only gas or supercritical fluid.

Critical Pressure (Pc)

Minimum pressure needed to liquefy a gas at its critical temperature.

Triple Point

Temperature and pressure where solid, liquid, and gas phases coexist.

Phase Diagram

Graphical representation of the phases of a substance as a function of temperature and pressure.

Heat of Fusion (ΔHf)

Heat required to melt 1 g or 1 mole of a solid at its melting point.

Heat of Vaporization (ΔHv)

Heat required to vaporize 1 g or 1 mole of a liquid at its boiling point.

Specific Heat (c)

Amount of heat needed to raise the temperature of 1 g of a substance by 1°C; Q = m c ΔT.

Fusion

Phase change from solid to liquid (melting).

Solidification

Phase change from liquid to solid (freezing).

Crystalline Solids

Solids with orderly, well-defined arrangements; typically have flat faces; examples include ice, sugar, salt, diamonds.

Amorphous Solids

Solids lacking long-range order; soften gradually when heated; examples include glass, rubber, plastics.

Unit Cell

Smallest repeating unit in a crystal that shows the full symmetry of the structure.

Simple Cubic Unit Cell (SC)

Corner atoms shared by 8 cells; each contributes 1/8 of an atom to the cell.

Body-Centered Cubic (BCC)

Center atom not shared; atoms at corners plus one at the cube’s center.

Face-Centered Cubic (FCC)

Corner and face-centered atoms; each contributes to the cell’s atom count; efficient packing.

Crystal Lattice

Three-dimensional arrangement of atoms, ions, or molecules in a crystalline solid.