Methane (, 16 g/mol) boils at about −162 °C, while water (, 18 g/mol) boils at 100 °C. The molecules are almost the same mass, so why is one a gas on the coldest Canadian winter day and the other a liquid in a hot kitchen? Boiling does not break any bonds inside the molecules; it pulls whole molecules away from their neighbours. What differs is how strongly the molecules attract each other. In this topic you will learn the attractions between particles, then use them, together with the bonds inside substances, to explain and predict the properties of every kind of solid.
- Act inside a molecule or throughout a crystal: covalent, ionic and metallic bonds
- Strong: breaking them is a chemical change
- Decide chemical properties and the melting points of ionic, network and metallic solids
- Act between separate molecules: dispersion, dipole-dipole, hydrogen bonding
- Much weaker: overcoming them is a physical change such as melting or boiling
- Decide the melting point, boiling point and solubility of molecular substances
- London dispersion forces act between all molecules. Electrons are always moving, so at any instant the electron cloud can be lopsided, creating a temporary dipole that induces a dipole in a neighbouring molecule. The more electrons a molecule has, the more easily its cloud distorts, so dispersion forces grow with molar mass. They also grow with surface contact: long, straight molecules touch their neighbours along more of their length than compact, branched ones.
- Dipole-dipole forces act between polar molecules. The partially positive end of one molecule attracts the partially negative end of the next. Among molecules of similar size, polar ones have higher boiling points than nonpolar ones.
- Hydrogen bonding is an especially strong dipole-dipole attraction. It forms when a hydrogen atom bonded directly to nitrogen, oxygen or fluorine is attracted to a lone pair on an N, O or F atom of a neighbouring molecule. The N–H, O–H and F–H bonds are so polar, and hydrogen is so small, that the attraction is unusually strong.
Hydrogen bonding explains many of water's unusual properties. Water boils at 100 °C while hydrogen sulfide, , a heavier molecule that cannot hydrogen bond, boils at about −60 °C. Hydrogen bonds give water a high surface tension, which lets some insects walk on a pond. When water freezes, hydrogen bonds lock the molecules into an open hexagonal arrangement with more empty space than the liquid, so ice is less dense than liquid water and floats, insulating the water below so lakes do not freeze solid.
| Force | Found between | Relative strength | Example |
|---|---|---|---|
| London dispersion | all molecules (the only IMF for nonpolar molecules) | weak per contact, but grows with molar mass and surface area | , , |
| Dipole-dipole | polar molecules | moderate | , |
| Hydrogen bonding | molecules with H bonded to N, O or F | strongest IMF | , , |
Four Types of Solids
| Property | Ionic | Molecular | Covalent network | Metallic |
|---|---|---|---|---|
| Particles | positive and negative ions | molecules | atoms | positive metal ions in a sea of delocalized electrons |
| Forces holding the solid together | ionic bonds throughout a lattice | intermolecular forces | covalent bonds throughout | metallic bonding |
| Melting point | high | low | very high | varies widely |
| Hardness | hard but brittle | soft | very hard (graphite is an exception) | malleable and ductile |
| Electrical conductivity | only when molten or dissolved | does not conduct | does not conduct (graphite does) | conducts as a solid and a liquid |
| Examples | NaCl, MgO, | ice, sucrose, iodine, dry ice () | diamond, graphite, quartz (), silicon carbide | copper, iron, sodium, mercury |
Each property follows from the particles and forces. Ionic solids are brittle because a sharp blow shifts one layer of ions so that like charges line up and repel, splitting the crystal; they conduct only when melted or dissolved, because only then are the ions free to move. Molecular solids melt easily because only weak IMFs must be overcome, and they never conduct because molecules are neutral. Covalent network solids are one giant molecule, so melting means breaking strong covalent bonds, which takes enormous energy. Graphite is the exception that proves the rule: its carbon atoms form flat sheets held together weakly, so the sheets slide (graphite is soft and slippery), and one electron per carbon is delocalized along each sheet, so it conducts. Metals conduct because their delocalized electrons move freely, and they are malleable because layers of metal ions can slide past each other while the electron sea keeps holding them together. Melting points of metals range from mercury (about −39 °C) to tungsten (about 3422 °C). Dissolving a solute in water also lowers the water's freezing point, which is why salt spread on winter roads melts ice.
Materials Designed from Structure
- Aramid fibres (Kevlar): long, rigid polymer chains line up side by side and are held together by hydrogen bonds between neighbouring chains, producing fibres strong enough for protective vests and cut-resistant gloves.
- Instant adhesives: cyanoacrylate glues are liquid monomers that polymerize very quickly when they touch traces of water on a surface, locking the surfaces together.
- Graphene and nanomaterials: graphene is a single layer of carbon atoms in a hexagonal network, a sheet of graphite one atom thick, and it is very strong and conducts electricity. Nanoparticles (roughly 1 to 100 nm across) have a huge surface area for their mass, which changes their properties; silver nanoparticles, for example, are added to some fabrics and wound dressings for their antibacterial effect.
- Superconductors: below a critical temperature these materials have zero electrical resistance, so current flows with no energy lost as heat. The strong magnets in MRI scanners use superconducting wire cooled by liquid helium, and some copper oxide ceramics, such as , superconduct above the boiling point of liquid nitrogen (77 K), which is much cheaper to use as a coolant.