intermolecular forces

Water boils at 100 degrees but butane is a gas at room temperature, even though both are small molecules. The difference is not the bonds inside each molecule but the attractions between separate molecules — the forces that decide whether a substance is a gas, a liquid, or a solid, how high it boils, and what it dissolves in. These between-molecule attractions are intermolecular forces, and although each is far weaker than a real chemical bond, together they govern nearly all the physical properties you can see and feel.

There are three main kinds, in rising strength. London dispersion forces are the weakest and most universal: even nonpolar molecules have electron clouds that flicker, creating fleeting, momentary dipoles that attract neighbours, and these grow stronger as molecules get larger and have more surface to touch, which is why bigger alkanes are oilier and boil higher. Dipole-dipole forces act between permanently polar molecules, lining up positive end to negative end. Hydrogen bonding is the strongest of the three, a special, extra-strong dipole attraction that occurs only when a hydrogen bonded to N, O or F is drawn to a lone pair on another N, O or F — it is what holds water together, pairs the strands of DNA, and folds proteins.

Intermolecular forces matter because they, not covalent bonds, set the physical behaviour of organic compounds: boiling and melting points, viscosity, surface tension, and especially solubility, summed up by 'like dissolves like' (polar dissolves polar, nonpolar dissolves nonpolar). One distinction worth nailing down: a hydrogen bond is an intermolecular attraction, not a real covalent bond, and it is some ten to a hundred times weaker — strong enough to organize liquids and biomolecules, weak enough to break and reform constantly. Boiling a liquid overcomes these forces; it does not break the molecules' internal bonds.