Liquids, Solids, and Phase Changes: Structure, Properties, and Energetics

Liquids, Solids, and Phase Changes

Intermolecular Forces

Intermolecular forces are the forces of attraction that give liquids and solids their fixed volumes and determine many of their physical properties. These forces act between molecules or between molecules and ions.

• Ion-dipole forces: Attraction between an ion and a polar molecule.

• Dipole-dipole forces: Attraction between polar molecules.

• Hydrogen bonds: A special, strong type of dipole-dipole interaction involving H bonded to N, O, or F.

• London dispersion forces: Weak attractions due to temporary dipoles in all molecules, especially significant in nonpolar substances.

Relative strengths: Ion-dipole > H-bond > dipole-dipole > London dispersion.

Boiling point is a measure of the strength of a substance's intermolecular forces.

Viscosity

Viscosity is a measure of a liquid’s resistance to flow. It is influenced by the strength of intermolecular forces and particle size.

Name	Formula	Viscosity (N·s/m2)
Pentane	C5H12	2.4 × 10-4
Benzene	C6H6	6.5 × 10-4
Water	H2O	1.00 × 10-3
Ethanol	C2H5OH	1.20 × 10-3
Glycerol	C3H5(OH)3	1.49

• Viscosity increases with van der Waals force strength: glycerol > water > ethanol > benzene > pentane.

• Viscosity is also affected by particle size (larger particles, higher viscosity).

Types of Solids

Solids can be classified as crystalline or amorphous:

• Crystalline solids: Atoms, ions, or molecules are arranged in an ordered, repeating pattern (e.g., iron, table salt, ice, diamond).

• Amorphous solids: Particles are randomly arranged with no long-range order (e.g., rubber).

Types of Crystalline Solids

Crystalline solids are further subclassified based on the nature of their constituent particles and bonding:

• Metallic

• Ionic

• Molecular

• Covalent network

Each subclass has different forces of attraction and thus different physical properties.

Metallic Solids

Metallic solids consist of atoms arranged in regular patterns. The most common packing arrangements are:

• Simple cubic packing

• Body-centered cubic packing

• Hexagonal closest packing

• Cubic closest packing

Unit Cells

The unit cell is the smallest repeating unit in a crystal lattice. There are three cubic unit cells:

• Primitive cubic

• Body-centered cubic

• Face-centered cubic

Metallic Solids Summary Table

Packing Type	Unit Cell	Number of Nearest Neighbors	Packing Efficiency	Number of Metals with Packing Type
Simple cubic	Primitive cubic	6	52%	1
Body-centered cubic	Body-centered cubic	8	68%	16
Hexagonal closest	Hexagonal closest	12	74%	21
Cubic closest	Face-centered cubic	12	74%	18

Conclusion: Particles in metallic solids pack as closely as possible to maximize interparticle attractions.

Properties of Metallic Solids

• Malleable (can be shaped)

• Conduct electricity

• Bonding involves delocalized electrons (metallic bonding)

Metallic bonds are not ionic or localized covalent bonds; instead, electrons are shared over many atoms, forming a 'sea of electrons' that explains malleability and conductivity.

Band Theory: Conductors, Insulators, and Semiconductors

• In metals (conductors), the conduction and valence bands overlap, allowing electrons to flow freely.

• In insulators, a large band gap prevents electron flow.

• In semiconductors, a small band gap allows limited conductivity under certain conditions.

Ionic Solids

Ionic solids are composed of cations and anions arranged in a regular lattice. The classic example is sodium chloride (NaCl), where chloride ions form a face-centered cubic lattice and sodium ions occupy the holes between them.

• High melting points (strong ionic bonds)

• Electrical nonconductors in solid state (ions fixed in place)

• Hard and brittle (distortion causes repulsion between like charges)

Molecular Solids

Molecular solids are held together by van der Waals forces (dipole-dipole, hydrogen bonds, London dispersion).

• Soft, low melting points

• Nonconductors (no charged particles)

Example: Ice is a molecular solid held together by hydrogen bonds, forming a hexagonal lattice.

Covalent Network Solids

Covalent network solids are held together by a continuous network of covalent bonds.

• Hard, high melting points

• Electrical nonconductors

Examples: Diamond (3D network of C atoms), graphite (2D sheets of C atoms).

Summary Table: Crystalline Solids and Their Properties

Type of Solid	Interparticle Forces	Properties
Metallic	Metallic bonds	Variable hardness and melting point, malleable, conducting
Ionic	Ion-ion forces (ionic bonds)	Brittle, hard, high-melting
Molecular	Dipole-dipole forces, hydrogen bonds, London dispersion forces	Soft, low-melting, nonconducting
Covalent network	Covalent bonds	Hard, high-melting

Phase Changes

Phase changes are physical changes in the state of matter that do not alter chemical identity.

• Fusion (melting): solid to liquid

• Freezing: liquid to solid

• Vaporization: liquid to gas

• Condensation: gas to liquid

• Sublimation: solid to gas

• Deposition: gas to solid

Enthalpy, Entropy, and Phase Changes

• Phase changes involve changes in enthalpy ($\Delta H$) and entropy ($\Delta S$).

• For example, vaporization: $\Delta H > 0$, $\Delta S > 0$ (energy and disorder increase).

Energetics of Phase Changes

Calculating the heat required for a phase change involves summing the energy for temperature changes and the energy for phase transitions:

• For temperature changes: $q = nC_{\text{(phase)}}\Delta T$

• For phase changes: $q = n\Delta H_{\text{(transition)}}$

Example: To convert 1.00 mol H2O from -25.0°C to 125.0°C, sum the energy for each step (heating ice, melting, heating liquid, boiling, heating steam).

Heating Curve for 1 mol of H2O

The heating curve shows temperature vs. heat added. Plateaus represent phase changes (melting, boiling), where temperature remains constant as energy is used to break intermolecular forces.

Molecular-Level Interpretation of Heating Curves

• During heating/cooling within a phase, temperature changes (kinetic energy changes).

• During a phase change, temperature is constant; potential energy (intermolecular forces) changes.

Evaporation and Vapor Pressure

Evaporation occurs because, at any temperature, some molecules have enough kinetic energy to escape the liquid phase. In a closed system, this leads to vapor pressure.

• Vapor pressure ($P_{\text{vap}}$): The pressure exerted by a vapor in equilibrium with its liquid.

• Vapor pressure increases with temperature and decreases with stronger intermolecular forces.

Example: Ether (dipole-dipole), ethanol (H-bonds), and water (two H-bonds per molecule) have different vapor pressures due to differences in intermolecular forces.

Clausius-Clapeyron Equation

The Clausius-Clapeyron equation relates vapor pressure and temperature:

$\ln P_{\text{vap}} = -\frac{\Delta H_{\text{vap}}}{R}\left(\frac{1}{T}\right) + C$

For two temperatures and pressures:

$\ln\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_{\text{vap}}}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)$

Where $R = 8.314$ J/(mol·K).

Vapor Pressure and Boiling Point

• The boiling point is the temperature at which vapor pressure equals external pressure.

• The normal boiling point is when vapor pressure equals 1 atm.

Summary of Key Points

• Viscosity increases with van der Waals force strength and particle size.

• Metals pack atoms closely to maximize attractions; metallic bonds involve delocalized electrons.

• Ionic solids are hard, brittle, and high-melting due to strong ionic bonds.

• Molecular solids are soft, low-melting, and nonconducting due to weak intermolecular forces.

• Covalent network solids are hard and high-melting due to strong covalent bonds.

• Phase changes alter physical state, not chemical identity.

• Vapor pressure is determined by temperature and intermolecular forces; described by the Clausius-Clapeyron equation.