Vaporization, Condensation & Vapor Pressure

• Vaporization is the change of state from liquid to gas, occurring via evaporation (below boiling point) or boiling (at boiling point).
• Cooling a vapor results in condensation, where gas changes to liquid due to decreasing kinetic energy and increasing attractive forces.
• Changes in state are reversible and involve opposite energy transfers (e.g., boiling requires energy input, condensation releases energy).
• In a closed container, evaporating liquid reaches equilibrium where the rate of molecules escaping into vapor equals the rate of vapor condensing back into liquid

Vapour Pressure

• Vapour pressure is the pressure exerted by a vapor in equilibrium with its liquid.
• Vapour pressure increases with temperature due to higher kinetic energy of gas particles, allowing them to overcome intermolecular forces more easily.
• Boiling point is when the vapor pressure of a liquid equals the external pressure (1 atmosphere or 101,325 Pa).

Solid State: Crystalline Structures

• Crystalline compounds (ionic, metallic, covalent) have a regular arrangement of particles called a crystal lattice.

Ionic Lattices

• Ionic lattices consist of alternating positive and negative ions in a three-dimensional arrangement.
• The type of lattice depends on the relative ion sizes.
• Compounds with ionic lattices are sometimes called giant ionic structures.
• Example: Sodium chloride (NaCl) and magnesium oxide (MgO) have cubic lattices.

Properties of Ionic Compounds

• Hardness: Requires significant energy to scratch due to strong attractive forces.
• Brittleness: Crystals split when a force displaces ion layers, causing like-charged ions to align and repel.
• High melting and boiling points result from strong electrostatic forces between oppositely charged ions acting in all directions.
• High melting point and boiling points increase with the charge density on the ions.
• Solubility: Many are soluble in water due to the formation of ion-dipole bonds.
• Electrical Conductivity: Conduct electricity only when molten or in solution due to mobile ions.

Metallic Lattices

• Metals conduct electricity when solid or liquid due to mobile delocalized electrons; ionic structures only conduct when molten or in aqueous solution due to mobile ions.
• Consist of ions surrounded by a "sea" of delocalized electrons, often packed in hexagonal layers or cubic arrangements.
• Metals are malleable and ductile because layers of ions can slide over each other, with new metallic bonds reforming easily due to delocalized electrons.
• High tensile strength and hardness are attributed to strong attractive forces between ions and delocalized electrons.

Alloys

• Alloys are mixtures of two or more metals, or a metal with a non-metal, where the added metal becomes part of the host metal's crystal lattice.
• Alloys are stronger than pure metals because different-sized ions disrupt the regularity of the lattice, hindering the sliding of layers.
• Example: Brass (70% copper, 30% zinc) is stronger and more malleable than pure copper.
• Aluminum alloys (with copper, magnesium, silicon, manganese) are lightweight, strong, and corrosion-resistant, suitable for aircraft bodies and engine blocks.
• Bronze is an alloy of copper and tin.

Simple Molecular Lattices

• Substances with simple molecular structures (e.g., iodine) can form crystals with molecules arranged in a lattice structure.
• Weak intermolecular forces between molecules lead to low melting points.
• The distance between nuclei of neighboring iodine molecules is greater than within the iodine molecule due to weak intermolecular forces versus strong covalent bonds.

Giant Molecular Structures

• Giant molecular structures (or giant covalent structures) feature a three-dimensional network of covalent bonds.
• High melting and boiling points due to the large number of strong covalent bonds.
• Examples: Carbon allotropes (diamond and graphite) and silicon dioxide (SiO2).
• Allotropes are different crystalline/molecular forms of the same element.

Graphite

• Carbon atoms arranged in planar layers, each bonded to three others in hexagons
• Each carbon atom has one electron in a p orbital that overlaps sideways, forming delocalized electron clouds above and below the plane.
• Layers held together by weak instantaneous dipole-induced dipole forces.
• High melting/boiling points due to strong covalent bonds within layers.
• Softness: Weak forces between layers allow them to slide easily, making graphite easily scratched and 'flaky'.
• Good electrical conductivity due to mobile delocalized electrons that can move along the layers.

Diamond

• Each carbon atom forms four covalent bonds in a tetrahedral arrangement.
• Atoms are arranged in a continuous network
• Regular arrangement gives diamond a crystalline structure.
• High melting and boiling points: strong covalent bonding throughout the structure.
• Hardness: difficult to scratch due to strong covalent bonds
• Does not conduct electricity because all four outer electrons are involved in covalent bonds.

Silicon(IV) Oxide (SiO2)

• Structure similar to diamond, each silicon atom bonded to four oxygen atoms, and each oxygen atom bonded to two silicon atoms.
• Forms hard, colorless crystals with high melting and boiling points.
• Does not conduct electricity.
• Sand is largely silicon(IV) oxide.

Key Properties of Giant Structures

• Most giant structures have networks of either covalent bonds (molecular giant structure), metallic bonds (metallic giant structure) or ionic bonds (ionic giant structure).
• The network of strong bonds is hard to break and so the melting and boiling points of these structures are generally very high

Fullerenes

• Fullerenes are allotropes of carbon, including hollow spheres or tubes (nanotubes).
• Individual particles have dimensions between 0.1 and 100 nanometers (nanoparticles).
• Structure is based on rings of carbon atoms (like graphite).
• Each carbon atom is bonded to three others.
• Rings of carbon atoms arranged in hexagons and sometimes pentagons.

Buckminsterfullerene (C60)

• A simple molecular structure with the shape of a football (soccer ball).
• Carbon atoms arranged at the corners of 20 hexagons and 12 pentagons.
• Some electrons are delocalized, but to a lesser extent than in graphite.
• Low sublimation point (600 °C) due to weak intermolecular forces.
• Relatively soft due to weak intermolecular forces.
• Poor conductor of electricity compared to graphite due to lower electron delocalization.
• Slightly soluble in solvents; more reactive than graphite or diamond.

Nanotubes

• Fullerenes of hexagonally arranged carbon atoms in a cylindrical form.
• High electrical conductivity along the long axis due to delocalized electrons.
• Very high tensile strength when force is applied along the long axis.
• Very high melting points (typically 3500 °C) due to strong covalent bonding
• These can be used as tiny wires in electrical circuits, to provide added strength to clothing and sports equipment and in cancer treatment

Graphene

• A single isolated layer of graphite in a hexagonally arranged sheet.
• Is not completely rigid and can be distorted.
• Most chemically reactive form of carbon
• Burns at very low temperatures and are much more reactive than graphite
• Extremely strong for its mass.
• For a given amount of material, graphene conducts electricity and heat much better than graphite.
• Applications include use in tiny electrical circuits and for tiny transistors, touchscreens, solar cells and other energy storage devices.