States of Matter Related to Pharmaceutical Formulations

Vocabulary-style flashcards covering states of matter, intermolecular forces, phase behavior, and related pharmaceutical concepts.

States of Matter

The distinct forms matter can take—solid, liquid, and gas—each with unique particle arrangement, kinetic energy, and volume.

Solid

A state with fixed shape and volume, strong intermolecular forces, and low kinetic energy; particles vibrate in fixed positions.

Liquid

A state with definite volume that takes the shape of its container; higher kinetic energy than a solid; flows and is moderately compressible.

Gas

A state with high kinetic energy, no fixed shape or volume, highly compressible, and capable of expanding to fill available space.

Vapor Pressure

The equilibrium pressure exerted by a vapor in equilibrium with its liquid; depends on temperature and not on container size.

Boiling Point

The temperature at which a liquid’s vapor pressure equals the surrounding atmospheric pressure.

Heat of Vaporization

Energy required to vaporize one mole of a liquid; reflects the strength of intermolecular forces.

Heat of Fusion

Energy required to melt one mole of a solid; reflects the strength of intermolecular forces in the solid.

Intermolecular Forces

Forces of attraction or repulsion between molecules, including cohesive and adhesive interactions.

Cohesive Forces

Attractive forces between like molecules.

Adhesive Forces

Attractive forces between unlike molecules.

Van der Waals Forces

General term for intermolecular forces including dipole–dipole, dipole–induced dipole, and induced–induced (London) interactions.

Dipole–Dipole Forces (Keesom)

Attraction between permanent dipoles in polar molecules; typically 1–7 kcal/mol; examples include water, alcohols, acetone.

Dipole–Induced Dipole Forces (Debye)

A polar molecule induces a temporary dipole in a nonpolar molecule; weaker than dipole–dipole; ~1–3 kcal/mol.

Induced Dipole–Induced Dipole Forces (London Dispersion)

Temporary attractions from instantaneous dipoles in nonpolar molecules; ~0.5–1 kcal/mol; important in gases and nonpolar substances.

Ion–Dipole Forces

Interactions between ions and polar molecules; typically 1–7 kcal/mol; important in salts and ion-containing species.

Hydrogen Bond

A strong dipole–dipole interaction between a hydrogen atom bonded to an electronegative atom (O, N, F) and another electronegative atom; highly directional and significant in water and biology.

Phase Diagram

A map of the phases of a substance as a function of temperature and pressure, showing regions for solid, liquid, and gas and their boundaries.

Eutectic Mixture

A mixture with a single melting point lower than the melting points of its components, shown as a eutectic point on a phase diagram.

Solids Properties (pharmaceutical context)

Fixed shape, high density, strong intermolecular forces, low kinetic energy; properties like surface energy, hardness, elasticity, porosity affect dosage forms.

Gaseous State and Ideal Gas Law

Gases under many conditions behave approximately as ideal gases; PV = nRT; Boyle’s law (P·V = constant at fixed T) and Charles’ law (V ∝ T).

Equilibrium Vaporization/Condensation

At equilibrium, rate of vaporization equals rate of condensation; defines equilibrium vapor pressure above the liquid.

Kinetic vs Potential Energy (States of Matter)

Kinetic energy increases with temperature and dominates in gases; potential energy relates to intermolecular forces and phase stability.