So far you've studied intramolecular forces — the bonds within a molecule (ionic, covalent) that hold its own atoms together. But when a substance melts or boils, those bonds don't break; what breaks are the much weaker forces between separate molecules — intermolecular forces (IMFs). Understanding IMFs explains why some molecular substances are gases at room temperature and others are liquids or solids. Every molecular substance is a crowd of molecules that are constantly moving and bumping into each other. The IMFs are the "stickiness" that holds the crowd together. Heating a substance makes its molecules move faster; it boils when they move fast enough to tear free of their neighbours. The stickier the molecules, the hotter you must go, so the whole trick to predicting melting and boiling points is to figure out which IMFs a substance has and how strong they are.
- Covalent bonds holding a molecule's own atoms together
- Very strong: breaking them is a chemical change
- Example: the two O–H bonds inside one H₂O molecule
- Broken in reactions such as electrolysis of water
- Attractions between separate, neighbouring molecules
- Much weaker: overcoming them is a physical change
- Example: the attraction between two separate H₂O molecules
- Overcome when a substance melts, boils or evaporates
London dispersion forces (LDFs)
Present in every molecule (polar or nonpolar), London dispersion forces arise from momentary, random shifts in electron distribution that create a brief, temporary dipole — which then induces a matching temporary dipole in a neighbouring molecule. Individually weak, but the only IMF available to nonpolar molecules. LDF strength increases with molar mass and the number of electrons, since a larger, more spread-out electron cloud is easier to temporarily distort. The halogens show this beautifully: fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid, even though all four are nonpolar diatomic molecules. Shape matters too. Long, straight molecules touch their neighbours along their whole length, while compact, ball-shaped molecules touch only at a few points. Pentane and its isomer 2,2-dimethylpropane have the same formula, , but the long chain boils at 36 °C and the ball-shaped one at about 10 °C.
| Halogen | Electrons per molecule | Boiling point | State at 25 °C |
|---|---|---|---|
| 18 | −188 °C | Gas | |
| 34 | −34 °C | Gas | |
| 70 | 59 °C | Liquid | |
| 106 | 184 °C | Solid |
Dipole-dipole forces
Between polar molecules, the permanent end of one molecule is attracted to the permanent end of a neighbouring molecule. This is stronger than LDFs between similarly sized molecules, because the attraction is constant rather than momentary. Compare two molecules with almost the same number of electrons: nonpolar (18 electrons) boils at −188 °C, while polar HCl (18 electrons) boils at −85 °C. Remember that polar molecules have dispersion forces as well; dipole-dipole forces are added on top.
Hydrogen bonding
A special, extra-strong case of dipole-dipole attraction that occurs only when hydrogen is bonded directly to nitrogen (N), oxygen (O), or fluorine (F) — the three smallest, most electronegative elements. The N–H, O–H, or F–H bond is intensely polar, and the tiny, exposed hydrogen nucleus can get unusually close to a lone pair on a neighbouring N, O, or F atom, producing an attraction far stronger than a typical dipole-dipole force. Despite the name, a hydrogen bond is not a covalent bond: it is an attraction between two different molecules. The clearest proof of its strength compares ethanol and dimethyl ether, which share the formula . Ethanol has an O–H group and boils at 78 °C; dimethyl ether has its oxygen between two carbons, no O–H, and boils at −24 °C.
| IMF | Present in | Relative strength | Requires |
|---|---|---|---|
| London dispersion | All molecules | Weakest (but adds up with size) | Any molecule — temporary induced dipoles |
| Dipole-dipole | Polar molecules | Moderate | A permanent bond dipole / net molecular dipole |
| Hydrogen bonding | Molecules with N–H, O–H, or F–H | Strongest of the three | H directly bonded to N, O, or F |
Predicting and comparing boiling points
- Decide whether each substance is molecular (IMFs apply) or ionic (much stronger ionic bonds hold the lattice, so melting and boiling points are far higher).
- Every molecule has London dispersion forces; note the number of electrons (or molar mass) to judge their size.
- Use shape and ΔEN to decide whether the molecule is polar. If it is, add dipole-dipole forces.
- Look for H bonded directly to N, O or F. If present, add hydrogen bonding.
- Compare: for molecules of similar size, the one with the stronger type of IMF has the higher boiling point. If sizes differ a lot, dispersion forces may decide it.