Rucete ✏ AP Chemistry In a Nutshell
7. Liquids and Solids
This chapter explains the structure, properties, and intermolecular forces of liquids and solids. It describes physical behaviors such as boiling, vapor pressure, viscosity, and melting, using molecular structure and forces as a basis.
Comparison of Liquids and Solids to Gases
• Liquids and solids are much denser than gases and have stronger intermolecular forces.
• Gases expand to fill containers; liquids take the shape of containers but have definite volume; solids have both definite shape and volume.
• The differences in physical properties stem from the strength of molecular attractions.
Dipole-Dipole Forces
• Polar molecules align so the partially positive end (δ⁺) of one is near the partially negative end (δ⁻) of another.
• Stronger dipole-dipole forces result in higher boiling points and influence surface tension, viscosity, and solubility.
• These forces are significant in liquids and solids where molecules are closer together.
London Dispersion Forces
• Nonpolar molecules develop temporary (instantaneous) dipoles due to unequal electron distribution.
• Induced dipoles can attract each other, forming weak London forces.
• Larger atoms or molecules with more electrons are more polarizable and exhibit stronger London forces.
Hydrogen Bonding
• Special case of dipole-dipole force occurring when hydrogen is bonded to F, O, or N.
• Leads to abnormally high boiling points (e.g., water, ammonia, HF).
• Water forms a network of hydrogen bonds that affect properties like density and boiling point.
Summary of Attractive Forces
• London forces: weak, temporary attractions from shifting electrons.
• Dipole-dipole forces: permanent polarity between molecules.
• Hydrogen bonding: strongest, occurs only with H–F, H–O, or H–N bonds.
Surface Tension
• Results from greater cohesive forces at the surface of a liquid.
• Causes liquids to form droplets and resist external force (e.g., insects walking on water).
• Cohesive vs. adhesive forces determine if a liquid beads or spreads on a surface.
Viscosity
• Resistance to flow caused by molecular attraction.
• Decreases with increasing temperature due to higher molecular kinetic energy.
• Hydrogen bonding and large molecules increase viscosity.
Evaporation
• Molecules at the surface with enough kinetic energy escape into gas phase.
• Rate increases with temperature and surface area.
• Evaporation causes cooling because high-energy molecules leave the liquid.
Vapor Pressure
• Pressure exerted by vapor in equilibrium with its liquid.
• Increases with temperature; depends on strength of intermolecular forces.
• Dynamic equilibrium is reached when evaporation equals condensation.
Boiling Point
• The temperature at which the vapor pressure equals external pressure.
• Normal boiling point is measured at 1 atm.
• Lower external pressure (e.g., at high altitudes) → lower boiling point.
• Stronger intermolecular forces → higher boiling point.
Heat of Vaporization
• Energy required to convert 1 mol of liquid to gas at its boiling point.
• Higher for substances with stronger intermolecular forces.
• Related to cooling effects during evaporation.
Melting Point and Heat of Fusion
• Melting point: temperature at which solid and liquid phases are in equilibrium.
• Heat of fusion: energy needed to melt 1 mol of a solid at melting point.
• Ionic solids usually have higher melting points than molecular solids.
Types of Solids
• Crystalline solids: highly ordered structures (e.g., NaCl, quartz).
• Amorphous solids: lack long-range order (e.g., glass, plastics).
• Classification based on bonding: ionic, molecular, covalent network, and metallic solids.
Ionic Solids
• Made of cations and anions held together by strong electrostatic forces.
• High melting points, hard and brittle, poor conductors unless molten or dissolved.
Molecular Solids
• Held together by intermolecular forces (London, dipole, hydrogen bonding).
• Soft, low melting points, often volatile or poor conductors.
Covalent Network Solids
• Atoms covalently bonded in a continuous network (e.g., diamond, SiO₂).
• Very hard, very high melting points, poor conductors (except graphite).
Metallic Solids
• Metal cations surrounded by delocalized electrons ("sea of electrons").
• Malleable, ductile, good conductors, varying hardness and melting points.
Phase Changes
• Transitions between solid, liquid, and gas require or release energy.
• Endothermic: melting, vaporization, sublimation (absorb heat).
• Exothermic: condensation, freezing, deposition (release heat).
Heating Curves
• Show temperature changes during heating, with flat segments representing phase changes.
• Use specific heat (q = mcΔT) for sloped regions and enthalpy values (ΔH) for flat regions.
Phase Diagrams
• Graphs showing phase stability as a function of pressure and temperature.
• Triple point: all three phases coexist.
• Critical point: temperature/pressure above which liquid and gas become indistinguishable (supercritical fluid).
In a Nutshell
Liquids and solids exhibit behaviors based on intermolecular forces, which affect properties like boiling point, melting point, vapor pressure, and structure. Understanding how molecular interactions influence physical properties allows prediction of phase changes and classification of materials based on bonding and structure.