Liquids, Solids, and Intermolecular Forces; Crystalline Solids and Modern Materials

Liquids, Solids, and Intermolecular Forces

Structure Determines Properties

The physical properties of substances are determined by the types and strengths of intermolecular forces present between molecules and atoms. These forces influence whether a substance exists as a solid, liquid, or gas under given conditions.

• Intermolecular Forces: Attractive forces between molecules and atoms. Their strength can be determined by chemical composition and molecular structure.

• Phases of Matter: The state (solid, liquid, gas) depends on the magnitude of intermolecular forces and thermal energy.

• Thermal Energy: Energy associated with the random motion of atoms and molecules.

Properties of Phases of Matter

Phase	Shape	Volume	Compressibility	Density
Gas	Indefinite	Indefinite	High	Low
Liquid	Indefinite	Definite	Low	High
Solid	Definite	Definite	Low	High

• Most solids have greater density than their liquids; H2O is an exception (ice floats).

Kinetic Molecular Theory

This theory explains the behavior of particles in different states of matter based on their kinetic energy and the forces between them.

• Kinetic Energy (KE): Energy of motion, $KE = \frac{1}{2}mv^2$

• Thermal energy increases with temperature.

• In gases, KE overcomes attractions; in solids, KE cannot overcome attractions; in liquids, KE partially overcomes attractions.

Degrees of Freedom

• Translational: Movement from place to place.

• Rotational: Rotation about an axis.

• Vibrational: Periodic motion of atoms within molecules.

Phase Changes

• Changing state requires overcoming intermolecular forces.

• Melting, vaporization, and sublimation require energy input (endothermic).

• Freezing, condensation, and deposition release energy (exothermic).

Intermolecular Forces

Intermolecular forces are the attractions between molecules that determine the physical properties of substances.

• Dispersion (London) Forces: Present in all molecules and atoms; result from temporary shifts in electron density.

• Dipole-Dipole Forces: Occur between polar molecules; strength increases with polarity.

• Hydrogen Bonding: Special dipole-dipole interaction between H and N, O, or F; strongest type of dipole-dipole force.

• Ion-Dipole Forces: Occur between ionic compounds and polar molecules; important in solutions.

Effect of Molecular Shape on Dispersion Forces

• Longer, less branched chains have higher boiling points due to increased surface area for dispersion forces.

Summary Table: Types of Intermolecular Forces

Type	Occurs Between	Relative Strength
Dispersion	All molecules/atoms	Weakest
Dipole-Dipole	Polar molecules	Intermediate
Hydrogen Bonding	H with N, O, or F	Strongest (of dipole-dipole)
Ion-Dipole	Ions and polar molecules	Strongest overall

Surface Tension

• Energy required to increase the surface area of a liquid.

• Stronger intermolecular forces lead to higher surface tension.

Viscosity

• Resistance of a liquid to flow.

• Increases with stronger intermolecular forces and longer molecular chains.

Capillary Action

• Ability of a liquid to flow against gravity up a narrow tube due to adhesive and cohesive forces.

Vaporization and Condensation

• Vaporization: Molecules escape from liquid to gas phase; requires energy (endothermic).

• Condensation: Gas molecules lose energy and return to liquid phase (exothermic).

Dynamic Equilibrium and Vapor Pressure

• At equilibrium, the rate of vaporization equals the rate of condensation.

• Vapor Pressure: Pressure exerted by vapor in equilibrium with its liquid.

• Increases with temperature and weaker intermolecular forces.

Clausius-Clapeyron Equation

Describes the relationship between vapor pressure and temperature:

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

• Where $P$ is vapor pressure, $T$ is temperature, $\Delta H_{vap}$ is enthalpy of vaporization, and $R$ is the gas constant.

Phase Diagrams

• Show the state of a substance at various temperatures and pressures.

• Triple Point: All three phases coexist.

• Critical Point: End of the liquid-gas boundary; above this, the substance is a supercritical fluid.

Crystalline Solids and Modern Materials

Crystalline Solids: Their Structure

• Atoms, ions, or molecules are arranged in a regular, repeating pattern.

• Described by a crystal lattice and unit cell.

Diffraction from a Crystal

• When X-rays strike a crystal, constructive and destructive interference occurs, producing a diffraction pattern.

• Bragg's Law: $n\lambda = 2d\sin\theta$

• Used to determine distances between layers of atoms in a crystal.

Covalent Network Solids

• Atoms held together by covalent bonds in a continuous network.

• Examples: Diamond, graphite, quartz (SiO2).

Properties of Diamond and Graphite

• Diamond: Each carbon forms four covalent bonds; very hard, high melting point, non-conductor.

• Graphite: Layers of carbon atoms in hexagonal sheets; soft, conducts electricity due to delocalized electrons.

Band Theory

• Describes bonding in solids using molecular orbitals spread over the entire crystal.

• Explains conductivity in metals, semiconductors, and insulators.

• Smaller band gap = better conductivity.

Polymers

• Large molecules made of repeating units (monomers).

• Can be natural (proteins, DNA) or synthetic (plastics).

• Addition Polymerization: Monomers add together without loss of atoms.

• Condensation Polymerization: Monomers join with loss of small molecules (e.g., water).