CHEM 212 General Chemistry II – Intermolecular Forces and the Physical Properties of Liquids and Solids

Vocabulary-style flashcards covering key concepts from intermolecular forces, liquid/solid properties, and related topics from the CHEM 212 notes.

Intermolecular forces

Attractive forces between molecules in the condensed phases (liquids and solids) that determine state and properties; include van der Waals forces, dipole–dipole interactions, hydrogen bonding, and ion–dipole interactions.

Van der Waals forces

A group of weak intermolecular forces that includes dipole–dipole interactions, London dispersion forces, and ion–dipole interactions.

Dipole moment

A measure of molecular polarity arising from unequal charge distribution; related to electronegativity differences and often expressed in Debye units.

Dipole–dipole interactions

Attractive forces between polar molecules due to the alignment of permanent dipoles.

Hydrogen bonding

A strong type of dipole–dipole interaction where hydrogen is bonded to N, O, or F and interacts with lone pairs on a electronegative atom; donor and acceptor concepts apply.

London dispersion forces (dispersion forces)

Intermolecular attractions arising from instantaneous dipoles; present in all molecules and increase with molecular size/polarizability.

Ion–dipole interactions

Coulombic attractions between ions and polar molecules; strength depends on ion charge/size and dipole moment/size of the polar molecule.

Polar molecule

A molecule with a net dipole moment due to uneven distribution of electron density, often from polar bonds or asymmetric geometry.

Nonpolar molecule

A molecule with zero net dipole moment; usually symmetric; may exhibit dispersion forces as the primary intermolecular force.

Polarizability

Ease with which a molecule’s electron cloud can be distorted; larger or more easily distorted clouds lead to stronger dispersion forces.

Hydrogen-bond donor

A molecule or group that donates a hydrogen atom in a hydrogen bond (e.g., N–H, O–H, F–H bonds).

Hydrogen-bond acceptor

An atom with lone pairs (such as N, O, or F) that can accept a hydrogen bond.

Water and hydrogen bonding

Water’s properties (e.g., high boiling point, surface tension, ice structure) arise from extensive hydrogen bonding.

Phase changes

Transitions between solid, liquid, and gas (solid–liquid, liquid–gas, solid–gas); described by phase diagrams.

Crystal types

Ionic crystals (ionic bonds), covalent (network) crystals, molecular crystals (molecules held by intermolecular forces), and metallic crystals (metallic bonds).

Unit cell

The smallest repeating unit of a crystal lattice that defines the structure and symmetry of the crystal.

Closest packing

Efficient arrangement of spheres in a crystal (e.g., FCC or HCP) giving maximum packing density.

Amorphous solids

Solids lacking long-range order in their atomic arrangement (e.g., glass) compared with crystalline solids.

Surface tension

Energy required to increase the surface area of a liquid; arises from cohesive forces at the surface.

Viscosity

Resistance of a liquid to flow; higher viscosity means thicker, slower-flowing liquids.

Vapor pressure

Pressure exerted by a vapor in equilibrium with its condensed phase at a given temperature; relates to volatility.

Boiling point vs. dipole moment

Boiling points tend to be higher for molecules with larger dipole moments (for substances with similar molar mass), reflecting stronger intermolecular forces.

CO2 polarity

CO2 has polar C=O bonds, but the molecule is linear and nonpolar overall because bond dipoles cancel.

CH4 vs CH2Cl2 vs CH3Cl

CH4 is nonpolar with a tetrahedral geometry; CH2Cl2 and CH3Cl are polar due to substituent differences causing a net dipole moment.

Solvation around ions (water)

Water molecules orient around ions to stabilize them via ion–dipole interactions, forming hydration shells.

Cations vs anions in solvation

Cations generally interact more strongly with polar solvents than anions, due to favorable orientation and charge distribution.