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Intermolecular forces (IMFs) are attractive forces between molecules that determine physical properties like boiling point, melting point, viscosity, and surface tension. London dispersion forces (temporary induced dipoles) act on all molecules and increase with molecular size and polarizability. Polar molecules additionally experience dipole-dipole forces. Hydrogen bonding — a strong dipole-dipole interaction — occurs when H is bonded directly to N, O, or F and is responsible for water's anomalously high boiling point and many biological phenomena.
Rank compounds by expected boiling point by identifying their dominant IMF type and relative strengths. Compare isomers like n-pentane vs. neopentane (both dispersion only, but different surface areas) and ethanol vs. dimethyl ether (H-bonding vs. dipole-dipole).
Covalent bonds hold atoms together within a molecule. But what holds molecules close to each other — as a liquid or solid — rather than flying apart as a gas? The answer is intermolecular forces (IMFs): attractive interactions between molecules. These forces are electrostatic in origin (opposite charges attract), but they arise from the distribution of electrons rather than from full ionic charges. Understanding IMFs explains a huge range of physical properties: why water is liquid at room temperature, why oils don't mix with water, why large alkanes are waxes while small ones are gases.
The weakest IMFs are London dispersion forces, which act on every molecule, polar or nonpolar. They arise from instantaneous fluctuations in electron distribution: at any given moment, the electron cloud of a molecule might be shifted slightly to one side, creating a temporary dipole. This temporary dipole induces a complementary dipole in a neighboring molecule, and the two are momentarily attracted. The key variable is polarizability — how easily the electron cloud can be distorted. Large molecules with many electrons are more polarizable and therefore have stronger dispersion forces. This is why boiling points of nonpolar molecules (like the alkane series) increase steadily with molecular size: more carbons mean more electrons, more polarizability, and stronger dispersion forces.
Polar molecules experience dipole-dipole forces in addition to dispersion. When you learned about molecular polarity, you found that molecules like HCl and SO₂ have permanent dipole moments — one end is persistently δ+ and the other δ−. Adjacent polar molecules orient themselves so that opposite partial charges align, creating a net attraction. These dipole-dipole interactions are stronger than dispersion forces for molecules of similar size.
Hydrogen bonding is a special, strong form of dipole-dipole interaction that occurs only when H is covalently bonded to N, O, or F — the three most electronegative elements. Because these elements are so electronegative, they pull the shared electron pair far from the hydrogen, leaving it nearly bare (a proton with very little electron shielding). This δ+ hydrogen can then strongly attract a lone pair on a neighboring N, O, or F atom. The resulting hydrogen bond (typically 15–30 kJ/mol) is much stronger than ordinary dipole-dipole forces, though still far weaker than a covalent bond. Water's unusually high boiling point, surface tension, and its expansion upon freezing all trace back to its extensive hydrogen bonding network. In biology, hydrogen bonds are essential to DNA base pairing and protein secondary structure — they are strong enough to maintain structure but weak enough to be broken and reformed dynamically.