2 States of Matter - Intermolecular Forces and Phase Behavior

Vocabulary flashcards covering key concepts in intermolecular forces, states of matter, phase transitions, and solid-state chemistry from the lecture notes.

Intermolecular forces

Forces that stabilize aggregates of molecules in gases, liquids, and solids; include van der Waals forces, hydrogen bonds, and ion–dipole interactions; balance of attraction and repulsion governs structure.

Cohesion

Attraction of like molecules.

Adhesion

Attraction of unlike molecules.

Van der Waals forces

Weak intermolecular attractions arising from dipole–dipole, dipole–induced dipole, and dispersion (London) forces; include Keesom, Debye, and London components.

Dipole–dipole forces (Keesom)

Attraction between permanent dipoles; orientation-dependent energy.

Dipole–induced dipole forces (Debye)

Permanent dipole induces a dipole in a neighboring molecule, creating attraction.

London dispersion forces

Attraction due to induced dipoles in nonpolar molecules; present in all molecules and increase with molecular size.

Ion–dipole interactions

Attraction between an ion and a polar molecule; contributes to solvation and miscibility.

Ion–Ion interactions

Electrostatic attraction between oppositely charged ions; can be very strong and influence salts and solid-state properties.

Hydrogen bonds

Electrostatic attraction involving a hydrogen donor (e.g., O–H, N–H) and a highly electronegative acceptor; relatively weak (roughly 2–8 kcal/mol) but crucial in water and biopolymers.

Ideal gas law

PV = nRT; assumes no intermolecular interactions and elastic collisions; used for calculating state relationships of gases.

Molar gas constant (R)

Constant in PV = nRT; value depends on units (e.g., 0.0821 L·atm·mol⁻¹·K⁻¹ when using liter-atm units).

Van der Waals equation

Real-gas equation that corrects the ideal gas law with constants a (attraction) and b (molar volume); P = RT/(V−nb) − a(n/V)² for one mole.

Critical temperature

Temperature above which a gas cannot be liquefied by pressure alone; Tc indicates strength of intermolecular attractions.

Critical pressure

Pressure required to liquefy a gas at its critical temperature.

Sublimation

Phase transition directly from solid to gas (deposition is the reverse process from gas to solid).

Freeze-drying (lyophilization)

Removal of solvent (usually water) by sublimation from frozen solids to yield dry powders; used for heat-sensitive pharmaceuticals.

Mesophase / Liquid crystal

Fourth state between liquid and solid; intermediate mobility; includes smectic and nematic phases.

Smectic phase

Liquid crystalline phase where molecules are mobile in two directions and rotate about one axis in layered (sheet-like) arrangements.

Nematic phase

Liquid crystal phase where molecules rotate about one axis and are mobile in three dimensions.

Supercritical fluid

Fluid above a substance’s critical temperature and pressure, exhibiting both gas-like and liquid-like properties; used for extraction and processing (e.g., CO₂).

Critical point

Point at which the properties of gas and liquid become indistinguishable; defined by Tc and Pc.

Phase rule (Gibbs)

F = C − P + 2; F is degrees of freedom, C = number of components, P = number of phases; governs how many intensive variables must be fixed.

Phase

A homogeneous, physically distinct portion of a system in contact with bounding surfaces; e.g., ice, water, and vapor can coexist as phases.

Triple point

Conditions where solid, liquid, and vapor of a substance are in equilibrium; for water, approximately 0.01°C and 0.0061 atm.

Tie line

In a two-phase region, a line on a phase diagram connecting the compositions of conjugate phases at a given temperature.

Binodal (miscibility gap)

Boundary in a phase diagram separating single-phase from two-phase regions; indicates immiscibility limits in binary mixtures.

Eutectic point

Composition at which a mixture has the lowest melting point and solidifies/melts at a single temperature into two solids and a liquid; involves invariant two-phase or three-phase equilibria.

Polymorphism

Ability of a solid to exist in more than one crystal form with different packing, melting points, solubilities, and diffraction patterns.

Crystallinity

Degree of structural order in a solid; crystalline vs amorphous states; affects melting point and dissolution.

Amorphous solids

Disordered solids lacking long-range order; may be glassy and sometimes crystallize over time; often exhibit glass transition temperature (Tg).

Solvate / solvates

Crystals that incorporate solvent molecules into the lattice; can be pseudopolymorphs with unique diffraction patterns.

X-ray diffraction (XRD)

Technique to determine crystal structures by analyzing diffraction of X-rays by a crystal lattice; includes powder XRD as a fingerprint for phases.

Differential Scanning Calorimetry (DSC)

Thermal analysis that measures heat flow to study thermal transitions (melting, crystallization, glass transitions) and determine ΔHf.

Thermogravimetric analysis (TGA)

Measures weight change of a sample as a function of temperature; used to study desolvation, decomposition, and moisture content.

Karl Fischer titration

Analytical method to quantify water content in solids via a redox-based titration involving iodine and sulfur reagents.

Vapor pressure

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

Boiling point

Temperature at which a liquid’s vapor pressure equals the surrounding pressure; latent heat of vaporization is released/absorbed during phase change.

Latent heat of vaporization (ΔHvap)

Energy required to vaporize one mole of liquid at its boiling point; varies with temperature (e.g., water ≈ 9720 cal/mol at 100°C).

Heats of fusion (ΔHf)

Energy required to melt one mole of solid; example: water ~1436 cal/mol at 0°C.

Tie-line distribution (phase diagrams)

Concept in ternary and binary diagrams describing how compositions of coexisting phases relate along a tie line.