Forces Between Molecules, Not Within Them
Why intermolecular forces are a completely different concept from bonds
Every molecule discussed so far in this Chemical Bonding unit has been held together internally by intramolecular forces — the covalent bonds within a single molecule, which are relatively strong and require significant energy to break. Intermolecular forces (IMFs) are a fundamentally different, generally much weaker category: the attractive forces that exist between separate, neighboring molecules, holding a collection of many individual molecules together as a bulk liquid or solid.
This distinction explains a common point of confusion: boiling a liquid, like water, does not break any of the covalent O-H bonds within individual water molecules — the water molecules remain fully intact H₂O molecules even as steam. What boiling actually overcomes is the intermolecular forces holding neighboring water molecules to each other, freeing individual molecules to escape into the gas phase as separate, independent particles. This is exactly why boiling points are governed by IMF strength, not by the strength of a substance's internal covalent bonds.
Three main categories of intermolecular force are covered in introductory chemistry, and they exist in a clear order of increasing strength, captured by the mnemonic LDH: London dispersion forces (the weakest, but present in every molecule without exception), dipole-dipole interactions (present specifically in polar molecules), and hydrogen bonding (the strongest of the three, a special, particularly strong case of dipole-dipole interaction that only occurs under specific structural conditions).
💡 Why IMF Strength Directly Predicts Boiling Point
Boiling a liquid requires supplying enough energy to overcome the intermolecular forces holding its molecules together, allowing individual molecules to break free into the gas phase. This means substances with stronger intermolecular forces require more energy — and therefore a higher temperature — to boil, giving a direct, predictable relationship between IMF strength and boiling point that holds true across an enormous range of substances.
Water's famously anomalous boiling point is the clearest illustration of this principle. Based purely on its very small molecular size and mass, water 'should' have a much lower boiling point, more comparable to other small molecules of similar mass. Instead, water boils at 100°C — dramatically higher than would be predicted from its size alone — specifically because water engages extensively in hydrogen bonding, the strongest of the three IMF categories, requiring considerably more energy to separate its molecules than weaker dipole-dipole or London dispersion forces alone would require. This single fact — that IMF strength, not molecular size, is the dominant factor determining boiling point — resolves what would otherwise be a confusing exception and turns it into a clear, predictable illustration of the underlying principle.
LD
London dispersion forces — present in every molecule
London dispersion forces are the weakest of the three IMF categories, but they are also universal — present to some degree in every molecule and atom, including completely nonpolar ones, since they don't depend on permanent polarity at all. They arise from temporary, momentary fluctuations in electron distribution: at any given instant, an atom or molecule's electrons might be distributed slightly unevenly by pure random chance, creating a fleeting, temporary dipole. This temporary dipole can induce a corresponding temporary dipole in a neighboring molecule, and the resulting brief, constantly fluctuating attraction between the two is the London dispersion force.
Although individually weak and momentary, London dispersion forces become stronger as molecular size (and specifically, the number of electrons present) increases, since larger electron clouds are more easily and more significantly distorted into a temporary dipole. This is why larger nonpolar molecules can actually have higher boiling points than smaller polar molecules in some cases — sufficiently strong, cumulative London dispersion forces can outweigh weaker dipole-dipole interactions in a smaller molecule.
Among the halogens, F₂ and Cl₂ are gases at room temperature, Br₂ is a liquid, and I₂ is a solid — a direct progression explained entirely by increasing London dispersion force strength as the number of electrons (and molecular size) increases down the group, since all four molecules are nonpolar and have no other IMF type available to them.
DD
Dipole-dipole interactions — between polar molecules
Dipole-dipole interactions occur specifically between polar molecules (covered in the Polarity lesson), arising from the permanent partial positive and partial negative charges each polar molecule possesses. Neighboring polar molecules orient themselves so the partial positive end of one molecule aligns near the partial negative end of an adjacent molecule, creating an attraction stronger than the momentary, temporary attraction of London dispersion forces alone. Every polar molecule experiences both dipole-dipole interactions AND London dispersion forces simultaneously (since dispersion forces are universal), meaning the two force types add together, or compound, rather than one type simply replacing the other.
HCl (polar, dipole-dipole present) has a higher boiling point than a similarly-sized nonpolar molecule would, precisely because it benefits from both dipole-dipole interactions and London dispersion forces working together, rather than dispersion forces alone.
HB
Hydrogen bonding — the strongest IMF, with specific structural requirements
Hydrogen bonding is not a true chemical bond in the same sense as a covalent or ionic bond — it's an especially strong, specific type of dipole-dipole interaction that occurs only under a precise structural condition: a hydrogen atom must be directly, covalently bonded to one of three highly electronegative, small atoms — nitrogen, oxygen, or fluorine (often remembered with the mnemonic 'N, O, F' or the phrase 'hydrogen bonds need N, O, F'). Because these three atoms are so strongly electronegative and so small, the H-N, H-O, and H-F bonds are unusually polar, and the resulting partial positive charge on the hydrogen atom is unusually concentrated and exposed — allowing it to interact especially strongly with a lone pair on a nitrogen, oxygen, or fluorine atom on a neighboring molecule.
This strict structural requirement means many polar molecules do NOT experience hydrogen bonding at all, even though they do experience ordinary dipole-dipole interactions — hydrogen bonding requires this specific N-H, O-H, or F-H combination, not just any polar bond involving hydrogen.
HF, H₂O, and NH₃ all exhibit hydrogen bonding (satisfying the N/O/F requirement) and all have anomalously high boiling points for their small size, exactly parallel to water's case — while a molecule like HCl, though polar, does not hydrogen bond (chlorine doesn't qualify), and has a correspondingly much lower boiling point relative to its similar molecular size.
🔬 Applied Scenario — Predicting and Comparing Boiling Points Using IMFs
Ranking or predicting the relative boiling points of different substances is one of the most common practical applications of IMF theory.
A
Step one: identify which IMF types are present. For any given molecule, first determine whether it's polar or nonpolar (from its Lewis structure and geometry), and if polar, whether it specifically contains an N-H, O-H, or F-H bond qualifying it for hydrogen bonding — every molecule has London dispersion forces regardless, but dipole-dipole and hydrogen bonding are conditional on the molecule's specific structure.
B
Step two: identify the strongest IMF type present, since it generally dominates the comparison. A molecule capable of hydrogen bonding will generally have a higher boiling point than a similarly-sized molecule that only has dipole-dipole interactions, which in turn will generally have a higher boiling point than a similarly-sized nonpolar molecule with only London dispersion forces — as long as molecular size is roughly comparable across the substances being compared.
C
Step three: account for molecular size when comparing across very different molecule sizes. Because London dispersion forces increase substantially with molecular size, a large enough nonpolar molecule can have a higher boiling point than a small polar or even hydrogen-bonding molecule — size and IMF type both need to be considered together, not IMF type alone, when the molecules being compared are very different in size.
D
Applying this to real substances: comparing propane, acetone, and ethanol (roughly similar molecular weight). Propane (nonpolar, London dispersion only) has the lowest boiling point of the three. Acetone (polar, dipole-dipole plus dispersion, but no N-H/O-H/F-H bond) has an intermediate boiling point. Ethanol (has an O-H bond, so hydrogen bonding plus dipole-dipole plus dispersion) has the highest boiling point of the three — a direct, real-world confirmation of the LDH strength ordering among molecules of comparable size.
⚠️ Most Common IMFs Mistakes
IMFs are not the same as, and are generally much weaker than, the covalent bonds inside a molecule — this confusion is extremely common. Students sometimes think boiling water breaks the O-H covalent bonds within each water molecule. Boiling only overcomes the (much weaker) hydrogen bonds between separate water molecules; the covalent O-H bonds within each individual water molecule remain completely intact in steam.
Hydrogen bonding requires hydrogen bonded DIRECTLY to nitrogen, oxygen, or fluorine — not just "a polar molecule with hydrogen in it somewhere." Students sometimes assume any molecule containing both hydrogen and a polar bond qualifies for hydrogen bonding. A molecule like HCl is polar and contains hydrogen, but chlorine doesn't satisfy the N/O/F requirement, so HCl does not hydrogen bond — it only experiences ordinary dipole-dipole interactions.
Every molecule has London dispersion forces, including polar ones — the different IMF types add together rather than replacing each other. Students sometimes think a polar molecule "only" has dipole-dipole forces and no London dispersion at all. In reality, a polar molecule experiences both dipole-dipole interactions AND London dispersion forces simultaneously; a hydrogen-bonding molecule experiences all three types (hydrogen bonding, dipole-dipole, and London dispersion) simultaneously, compounding together.
✓ Quick Self-Test
1. What is the difference between an intermolecular force and an intramolecular (covalent) bond, and why does this distinction matter for understanding boiling?
2. What are the three main types of IMF, listed from weakest to strongest?
3. Why do London dispersion forces increase with molecular size, and what real-world example (like the halogens) demonstrates this?
4. What specific structural requirement must be met for a molecule to exhibit hydrogen bonding?
5. Why does water have an anomalously high boiling point compared to what its small molecular size would otherwise predict?
Answers:
1. Intermolecular forces (IMFs) are attractive forces between separate, neighboring molecules; intramolecular (covalent) bonds are the forces holding atoms together within a single molecule. This distinction matters because boiling a liquid only overcomes the intermolecular forces between molecules, not the covalent bonds within each individual molecule — the molecules themselves remain chemically intact after boiling.
2. From weakest to strongest: London dispersion forces (present in every molecule, arising from temporary electron distribution fluctuations), dipole-dipole interactions (present in polar molecules, from permanent partial charges), and hydrogen bonding (the strongest, a special case of dipole-dipole interaction requiring hydrogen directly bonded to nitrogen, oxygen, or fluorine).
3. London dispersion forces increase with molecular size because larger molecules have more electrons, and larger electron clouds can be more easily and more significantly distorted into a temporary dipole, creating stronger momentary attractions. The halogens demonstrate this directly: F₂ and Cl₂ are gases, Br₂ is a liquid, and I₂ is a solid at room temperature, purely due to increasing dispersion force strength as molecular size and electron count increase down the group, since all four are nonpolar.
4. Hydrogen bonding requires a hydrogen atom to be directly, covalently bonded to one of three highly electronegative, small atoms: nitrogen, oxygen, or fluorine (often remembered as 'N, O, F'). This specific combination creates an unusually polar bond with a concentrated, exposed partial positive charge on hydrogen, allowing an especially strong interaction with a lone pair on a neighboring molecule's N, O, or F atom.
5. Water has an anomalously high boiling point because its O-H bonds allow extensive hydrogen bonding between water molecules — the strongest of the three IMF types. Despite water's very small molecular size (which would otherwise predict a low boiling point based on weak dispersion forces alone), the considerable additional energy required to overcome its hydrogen bonds raises its boiling point dramatically higher than its size alone would suggest.