An intermolecular bond is an attractive force that acts between molecules, holding them near each other without actually joining their atoms together. These forces are what keep water liquid at room temperature, make ice solid, and allow geckos to climb walls. They’re much weaker than the bonds holding atoms together within a molecule, but they determine most of the physical properties you can observe: whether a substance is a solid, liquid, or gas, how easily it evaporates, and how it behaves in your body.
Intermolecular vs. Intramolecular Bonds
The distinction matters because these two types of forces operate on completely different scales. Intramolecular forces are the bonds between atoms inside a molecule. They’re what makes water H₂O instead of separate hydrogen and oxygen atoms. Intermolecular forces, by contrast, are the attractions between one molecule and another.
A useful comparison: it takes about 464 kJ/mol of energy to break the bonds between oxygen and hydrogen atoms within a water molecule. It takes only about 19 kJ/mol to pull water molecules apart from each other. That’s roughly 24 times less energy. When you boil water, you’re overcoming the intermolecular forces between water molecules, not breaking the molecules themselves. The steam rising from your pot is still H₂O.
London Dispersion Forces
Every molecule experiences London dispersion forces, making them the most universal type of intermolecular attraction. They arise from the constant motion of electrons around atoms. At any given instant, electrons may cluster slightly to one side of a molecule, creating a tiny, fleeting imbalance of charge: one end becomes slightly negative, the other slightly positive. This temporary dipole then influences nearby molecules, nudging their electrons in response and creating a chain of momentary attractions.
These forces are individually very weak. In helium, they amount to just 0.076 kJ/mol. But they scale up with molecular size because larger molecules have more electrons that can shift around, making them more “polarizable,” meaning easier to distort into temporary dipoles. This is why small hydrocarbons like methane are gases at room temperature while larger ones like octane are liquids.
Shape matters too. Pentane and neopentane have the same molecular formula (C₅H₁₂) and the same weight, but pentane is a liquid at room temperature while neopentane is a gas. Pentane’s long, stretched-out shape lets molecules line up closely, maximizing contact area. Neopentane is roughly spherical, so molecules can’t pack together as tightly and the dispersion forces between them are weaker.
Dipole-Dipole Interactions
When atoms in a molecule share electrons unevenly (because one atom pulls harder on the electrons than the other), the molecule develops a permanent positive end and a permanent negative end. It becomes a polar molecule. Two polar molecules near each other will orient so that the positive end of one faces the negative end of the other, creating a steady attractive force. The dipole-dipole interaction in hydrogen chloride (HCl), for example, is about 3.3 kJ/mol.
In a liquid or gas, molecules are constantly tumbling and moving, so these alignments aren’t perfect. But statistically, polar molecules spend more time in favorable orientations (positive near negative) than unfavorable ones, because the low-energy arrangements are more probable. This gives polar compounds noticeably higher boiling points than nonpolar compounds of similar size.
Hydrogen Bonds
Hydrogen bonding is a special, stronger version of dipole-dipole interaction. It occurs when a hydrogen atom bonded to a highly electronegative atom (fluorine, oxygen, or nitrogen) is attracted to a lone pair of electrons on another fluorine, oxygen, or nitrogen atom nearby. The hydrogen in this arrangement acts like a bridge between two electron-hungry atoms.
What makes hydrogen bonds stronger than ordinary dipole-dipole forces is that they involve two reinforcing effects at once. First, there’s the straightforward attraction between the partially positive hydrogen and the partially negative atom it’s reaching toward. Second, there’s an orbital interaction where electrons from the neighboring atom partially overlap with the hydrogen’s bond, giving the connection a slight covalent character. Most hydrogen bonds fall in the range of 10 to 40 kJ/mol, with extreme cases like the fluorine-hydrogen-fluorine interaction reaching about 40 kJ/mol, roughly half the strength of a full carbon-carbon covalent bond.
The strength of a hydrogen bond also depends on how electronegative the atom bonded to hydrogen is. An oxygen-hydrogen bond to nitrogen is stronger than a nitrogen-hydrogen bond to nitrogen, because oxygen pulls harder on the shared electrons, leaving hydrogen with a larger partial positive charge and creating a bigger antibonding orbital for the donor atom to interact with.
Ion-Dipole Forces
When an ion (an atom or molecule carrying a full electric charge) meets a polar molecule, the resulting attraction is called an ion-dipole force. These are the strongest of the common intermolecular forces because a full charge interacts much more powerfully than a partial one.
This is exactly what happens when you dissolve table salt in water. The sodium and chloride ions separate, and water molecules immediately surround each ion, forming what’s called a hydration shell. Water molecules orient with their oxygen (the negative end) pointing toward sodium ions and their hydrogen (the positive end) pointing toward chloride ions. A sodium ion carries about 7 to 14 water molecules in its outer shell at any given moment. These shells are dynamic, with water molecules swapping in and out in picoseconds, but the overall structure persists. For doubly charged ions like magnesium, the shell is even more tightly held, with six water molecules locked in an inner shell and 12 to 14 in a second shell.
How Intermolecular Forces Shape Physical Properties
Boiling point is essentially a direct readout of intermolecular force strength. To boil a liquid, you need to give molecules enough kinetic energy to escape the pull of their neighbors. Compounds with stronger intermolecular forces need more energy, so they boil at higher temperatures. This is why water (with its extensive hydrogen bonding) boils at 100°C, while methane (relying only on weak dispersion forces) boils at -161°C despite having a similar molecular weight.
The pattern is consistent: in any group of similar compounds, those capable of hydrogen bonding have the highest boiling points, followed by polar molecules with dipole-dipole interactions, followed by nonpolar molecules relying on dispersion forces alone. Within each category, larger molecules boil at higher temperatures because their dispersion forces are stronger.
Surface tension follows the same logic. Water has a surface tension of 72.8 millinewtons per meter at room temperature, the highest of any common non-metallic, non-ionic liquid. This is because water molecules at the surface are pulled inward by hydrogen bonds with their neighbors below and beside them, creating a kind of elastic “skin.” That high surface tension, combined with water’s ability to adhere to other surfaces, drives capillary action. Trees depend on this: water climbs through narrow xylem vessels in trunks and branches because cohesion (water molecules sticking to each other through hydrogen bonds) holds the water column together while adhesion keeps it attached to vessel walls.
Intermolecular Forces in Biology
Hydrogen bonds hold the two strands of DNA together. Each rung of the double helix consists of paired bases connected by either two or three hydrogen bonds, depending on the pair. Individually these bonds are weak enough to be unzipped when a cell needs to read or copy its genetic code, but collectively, millions of them along a strand provide enormous stability.
Proteins fold into their functional shapes through a combination of intermolecular forces. Hydrogen bonds help form the spirals and sheets that make up a protein’s backbone structure. Dispersion forces between nonpolar amino acid side chains help drive the protein’s core to fold inward, away from water. Ion-dipole interactions at the protein’s surface keep it dissolved in the watery environment of your cells. Disrupting these forces (with heat, for instance) causes proteins to unfold and lose function, which is why cooking an egg turns it from liquid to solid.