Intermolecular bonds are the forces that hold molecules together in a liquid or solid. They’re not the same as the bonds holding atoms together within a single molecule. Instead, they act between neighboring molecules, pulling them toward each other through electrical attraction. These forces are much weaker than the bonds inside molecules, but they determine some of the most familiar properties of matter: whether a substance is a solid, liquid, or gas at room temperature, how high its boiling point is, and whether it dissolves in water.
Intermolecular vs. Intramolecular Bonds
The distinction here is straightforward but important. Intramolecular bonds are the strong connections between atoms inside a single molecule. The bonds holding hydrogen to oxygen in a water molecule, for example, require about 464 kJ/mol of energy to break. Intermolecular bonds are the much weaker attractions between one water molecule and the next, requiring only about 19 kJ/mol to overcome. That’s roughly 25 times weaker.
When water boils, you’re breaking intermolecular bonds. The molecules themselves stay intact as H₂O. They just separate from each other and float away as gas. When you melt ice, the same thing happens at a smaller scale: intermolecular bonds loosen enough for molecules to slide past one another, turning a rigid solid into a flowing liquid. The atoms within each molecule never change during these transitions.
All intermolecular forces are electrostatic in nature, meaning they arise from the attraction between positive and negative charges. The differences between types come down to where those charges originate and how strong they are.
London Dispersion Forces
Every molecule experiences London dispersion forces, making them the most universal type of intermolecular attraction. They exist even between completely nonpolar molecules that have no permanent electrical imbalance. In helium, for instance, these forces are incredibly small, around 0.076 kJ/mol, but they’re still there.
The mechanism works like this: electrons are constantly moving around a molecule’s atoms. At any given instant, they might cluster slightly toward one side, creating a fleeting imbalance where one end of the molecule is slightly negative and the other slightly positive. This temporary dipole then influences the electrons in a neighboring molecule, nudging them to shift as well. For a brief moment, the two molecules attract each other. This happens trillions of times per second across countless molecules, and the cumulative effect is a real, measurable attraction.
Two factors control how strong these forces get. The first is size. Larger, heavier molecules have more electrons spread over a greater volume, and those electrons are farther from the nucleus and less tightly held. They shift more easily, creating stronger temporary dipoles. Chemists call this quality “polarizability.” The second factor is shape. Two molecules with the same weight and formula can have very different dispersion forces depending on their geometry. Pentane and neopentane both have the formula C₅H₁₂, but pentane is an elongated chain while neopentane is compact and roughly spherical. The elongated shape lets pentane molecules line up alongside each other with more surface contact, producing stronger dispersion forces. At room temperature, pentane is a liquid while neopentane is a gas.
Dipole-Dipole Forces
When a molecule has a permanent electrical imbalance, it carries a dipole: one end is slightly positive, the other slightly negative. This happens when atoms in the molecule don’t share electrons equally. Oxygen, nitrogen, and fluorine pull electrons toward themselves more strongly than most other atoms, so any bond between one of these and a less electron-hungry atom creates a lopsided charge distribution.
Having polar bonds isn’t enough on its own, though. The molecule’s geometry matters. Carbon dioxide has two polar bonds, but they point in exactly opposite directions and cancel each other out, leaving the molecule nonpolar overall. Hydrogen fluoride, on the other hand, has a single polar bond with no opposing force to cancel it, giving the molecule a strong permanent dipole.
When polar molecules come near each other, the positive end of one is attracted to the negative end of another. This dipole-dipole interaction is stronger than London dispersion forces for molecules of similar size. In hydrogen chloride, the dipole-dipole interaction contributes about 3.3 kJ/mol of attractive energy.
Hydrogen Bonds
Hydrogen bonding is a particularly strong version of dipole-dipole attraction, and it plays an outsized role in biology and everyday chemistry. It occurs when a hydrogen atom bonded to nitrogen, oxygen, or fluorine is attracted to a nitrogen, oxygen, or fluorine atom on a neighboring molecule. These three elements are among the most electronegative on the periodic table (fluorine at 4.1, oxygen at 3.5, nitrogen at 3.0, compared to hydrogen’s 2.1), so the bond between hydrogen and any of them is extremely polar. The hydrogen carries a large partial positive charge, and the other atom carries a large partial negative charge.
Despite the name, a hydrogen bond is not a true bond in the way a covalent bond is. It’s an intermolecular attraction. But it’s strong enough to have dramatic effects. Water’s unusually high boiling point for such a small molecule, its ability to dissolve so many substances, and the fact that ice floats all trace back to hydrogen bonding. Each water molecule can form up to four hydrogen bonds with its neighbors, creating an extensive network of intermolecular connections.
In biology, hydrogen bonds are critical for maintaining the structure of both DNA and proteins. The two strands of the DNA double helix are held together by hydrogen bonds between paired bases. In proteins, hydrogen bonds between the backbone atoms are the primary force that folds the chain into its characteristic shapes, including the coiled alpha helices and flat beta sheets that make up the core of most protein structures. About two-thirds of all hydrogen bonds in a typical protein are these backbone-to-backbone connections that define its three-dimensional form.
Ion-Dipole Forces
When you drop table salt into water, something specific happens at the molecular level. Each sodium ion (positively charged) attracts the negative end of nearby water molecules, while each chloride ion (negatively charged) attracts the positive end. These ion-dipole forces pull water molecules around each ion in an organized shell, gradually prying the salt crystal apart.
Ion-dipole forces are among the strongest intermolecular interactions because they involve a full electrical charge on the ion, not just the partial charges of a dipole. They weaken with the square of the distance between particles, which is slower than the dropoff for other intermolecular forces. Smaller ions produce stronger ion-dipole interactions because the surrounding water molecules can get closer to the charge center. This is why the energy released when small ions dissolve (the enthalpy of hydration) is greater than for larger ions of the same charge. Some salts hold onto their water molecules so tightly that they trap water in their crystal structure when they re-solidify, forming what chemists call hydrated salts.
The Van der Waals Umbrella
You’ll often see the term “van der Waals forces” used in textbooks and articles, and it can be confusing because it doesn’t refer to a single type of force. According to the IUPAC definition (the international chemistry naming authority), van der Waals forces include dipole-dipole interactions, dipole-induced dipole interactions, and London dispersion forces. It essentially covers all the attractive and repulsive forces between molecules except for actual chemical bonds and the interactions involving fully charged ions. Some sources use the term loosely to mean any nonspecific intermolecular attraction, so if you see it, check the context to know which forces are being discussed.
How These Forces Affect Physical Properties
The hierarchy of intermolecular force strength, from strongest to weakest among the common types, runs: hydrogen bonding, then dipole-dipole, then London dispersion. This ranking directly predicts physical properties. Substances with stronger intermolecular forces need more heat energy to pull their molecules apart, so they have higher boiling points and higher melting points. They also tend to be more viscous because their molecules resist flowing past one another, and they have lower vapor pressure because fewer molecules have enough energy to escape into the gas phase at any given temperature.
This is why small nonpolar molecules like methane are gases at room temperature (only weak dispersion forces hold them together), while water, a similarly small molecule with strong hydrogen bonds, is a liquid. It’s also why rubbing alcohol evaporates faster than water from your skin: alcohol has weaker intermolecular forces, so its molecules escape into the air more readily. Whenever you’re comparing two substances and wondering why one boils at a higher temperature or stays liquid while the other evaporates, the answer almost always comes down to which one has stronger intermolecular forces.