Intermolecular Forces Explained: London Dispersion, Dipole-Dipole, and Hydrogen Bonding

Intermolecular forces (IMFs) are the attractions between separate molecules — not the bonds within a molecule. There are three main types, ranked from weakest to strongest: London dispersion forces (present in ALL molecules, caused by temporary electron fluctuations), dipole-dipole forces (present only in polar molecules, caused by permanent partial charges), and hydrogen bonding (a special strong case of dipole-dipole, occurring when H is bonded to N, O, or F). The type and strength of IMFs a substance experiences directly determine its boiling point, melting point, viscosity, surface tension, and solubility. Stronger IMFs = more energy needed to separate molecules = higher boiling point.

Every molecule experiences London dispersion forces (also called van der Waals forces or induced dipole forces). They are the only intermolecular force that nonpolar molecules experience, but polar molecules experience them too — on top of their other forces.

Here is how they work. Electrons are constantly moving around a molecule. At any given instant, the electron distribution may be slightly uneven — a bit more electron density on one side than the other. This creates a temporary, instantaneous dipole. That temporary dipole induces a complementary dipole in a neighboring molecule (the electrons in the neighbor shift in response). For a fraction of a second, the two molecules attract each other. Then the electrons rearrange and it happens again, billions of times per second.

The strength of London forces depends on two factors: the number of electrons (more electrons = more polarizable = stronger London forces) and the surface area of the molecule (more contact area = more opportunities for temporary dipoles to form). This is why boiling point increases with molar mass in a homologous series — pentane (C5H12, BP 36°C) boils higher than methane (CH4, BP -162°C) because pentane has more electrons and more surface area.

Shape matters too. Neopentane (compact, spherical) boils at 10°C, while n-pentane (long, linear) boils at 36°C — same molecular formula (C5H12), same number of electrons, but the linear shape allows more surface contact between molecules, strengthening the London forces. This shape effect is a favorite exam question.

Polar molecules — those with a permanent dipole moment because of asymmetric charge distribution — experience dipole-dipole attractions in addition to London forces. The positive end of one molecule attracts the negative end of a neighbor.

The key to predicting dipole-dipole forces is knowing whether the molecule is polar. Two conditions must be met: the molecule must have polar bonds (significant electronegativity difference between bonded atoms), AND the molecular geometry must be asymmetric so the bond dipoles do not cancel. Water (bent geometry, bond dipoles do not cancel) is polar. Carbon dioxide (linear, bond dipoles cancel perfectly) is nonpolar despite having polar C=O bonds.

Dipole-dipole forces are stronger than London forces for molecules of similar size. This is why acetone (polar, MW 58, BP 56°C) boils much higher than propane (nonpolar, MW 44, BP -42°C) despite being similar in size. The permanent dipole adds an additional attractive force that holds acetone molecules together more tightly.

One subtlety students miss: polar molecules still have London forces. A polar molecule experiences both dipole-dipole AND London forces. The total IMF strength is the sum. So when comparing boiling points, you need to consider all forces acting on each molecule, not just the dominant one.

Hydrogen bonding is not a separate category of force — it is an exceptionally strong dipole-dipole interaction that occurs when hydrogen is covalently bonded to nitrogen, oxygen, or fluorine (the three most electronegative elements after noble gases). The small size of hydrogen, combined with the extreme polarity of the N-H, O-H, or F-H bond, creates an unusually strong attraction to lone pairs on a nearby N, O, or F atom.

Water is the textbook example. Each water molecule can form up to four hydrogen bonds — two through its two O-H bonds (as a donor) and two through the two lone pairs on oxygen (as an acceptor). This extensive hydrogen bonding network explains virtually all of water's anomalous properties: its unusually high boiling point (100°C, compared to H2S at -60°C — both are Group 16 hydrides, but only water has H-bonding), its high specific heat capacity, its high surface tension, and the fact that ice floats (the H-bond network in ice creates an open hexagonal structure that is less dense than liquid water).

Hydrogen bonding is also critical in biology. DNA's double helix is held together by hydrogen bonds between base pairs (A-T has two, G-C has three). Protein secondary structure (alpha helices and beta sheets) is stabilized by hydrogen bonds between backbone amino and carbonyl groups. Without hydrogen bonding, life as we know it could not exist.

A common mistake: thinking any molecule with hydrogen has hydrogen bonds. It does not. Methane (CH4) has four C-H bonds, but carbon is not electronegative enough to create the extreme polarity needed for hydrogen bonding. Only N-H, O-H, and F-H qualify. ChemistryIQ has practice problems that test your ability to distinguish hydrogen-bonding molecules from non-hydrogen-bonding ones in tricky cases.

Once you can identify which forces a molecule experiences, predicting relative physical properties becomes systematic.

Boiling point comparison: identify the IMFs for each molecule, then rank. Stronger total IMFs = higher boiling point. Example: rank the boiling points of CH4, CH3Cl, and CH3OH. CH4 has only London forces (nonpolar, no permanent dipole, no H-bonding). CH3Cl has London + dipole-dipole (polar, but no H-bonding because H is bonded to C, not N/O/F). CH3OH has London + dipole-dipole + hydrogen bonding (polar, with an O-H bond). So BP order: CH4 < CH3Cl < CH3OH.

Solubility follows the rule like dissolves like. Polar solvents (water) dissolve polar solutes and ionic compounds. Nonpolar solvents (hexane) dissolve nonpolar solutes. Ethanol is miscible with both water (because of H-bonding through its O-H group) and with some nonpolar solvents (because of its hydrocarbon chain). This dual character is why ethanol is such a useful solvent.

Viscosity and surface tension both increase with stronger IMFs. Glycerol (three O-H groups, extensive H-bonding) is viscous and syrupy. Hexane (London forces only) pours like water. Motor oil is viscous because of London forces from its very long hydrocarbon chains — massive surface area means strong cumulative London attraction even without polarity.

Vapor pressure is inversely related to boiling point (and therefore inversely related to IMF strength). Molecules with weak IMFs escape into the gas phase easily, producing high vapor pressure at room temperature. Diethyl ether (London + weak dipole) has a high vapor pressure and evaporates quickly. Water (strong H-bonding) has much lower vapor pressure and evaporates slowly.