Intermolecular Forces — Hydrogen Bonding, Dipole and Dispersion
Water boils at 100 °C. Hydrogen sulfide, a heavier molecule with the same shape, boils at −60 °C. Nothing about the covalent bonds inside those molecules explains a 160-degree gap — the answer is entirely about the forces between molecules. Intermolecular forces decide melting point, boiling point, viscosity, surface tension and solubility, and they are worth far more marks than the single page most textbooks give them.
The one idea behind all of them
Every intermolecular force is electrostatic: a region of partial negative charge attracting a region of partial positive charge on a neighbouring molecule. The three types differ only in where those charges come from and how long they last.
They are all far weaker than the covalent bonds inside a molecule, and that is the key to a question students constantly get wrong. When water boils, no O–H bonds break. Boiling separates whole H₂O molecules from each other. Breaking O–H bonds would need about 460 kJ/mol; pulling the molecules apart needs about 41 kJ/mol.
1. London dispersion forces
Electrons move. At any instant they can be unevenly spread across a molecule, creating a momentary dipole, which induces an opposite dipole in the neighbour next to it. The two attract. The instant passes, but the effect repeats constantly, so on average there is a real attraction.
Dispersion forces exist in every substance — polar, non-polar, atoms, ions, everything. They are the only force available to non-polar molecules, and they strengthen with:
- More electrons (roughly, higher molar mass), because a larger electron cloud is more easily distorted — this is called higher polarisability.
- A longer, flatter shape, because more surface area can lie against a neighbour.
Two clean pieces of evidence. The noble gases have no dipoles at all, yet their boiling points climb steadily with size: He −269 °C, Ne −246 °C, Ar −186 °C, Kr −153 °C, Xe −108 °C. And the halogens go from gases to a solid for the same reason: F₂ −188 °C, Cl₂ −34 °C, Br₂ +59 °C, I₂ +184 °C. Iodine is a solid at room temperature purely because of dispersion.
2. Dipole–dipole forces
A molecule with a permanent dipole — a polar bond plus an unsymmetrical shape — has a fixed δ+ end and a fixed δ− end. Neighbouring molecules line up head to tail and attract. This is added on top of dispersion, never instead of it.
Being polar requires both conditions. CO₂ has two strongly polar C=O bonds, but it is linear, the two bond dipoles cancel exactly, and the molecule is non-polar. H₂O has the same two polar bonds arranged at 104.5°, they do not cancel, and it is strongly polar.
3. Hydrogen bonding
Hydrogen bonding is a specially strong dipole–dipole attraction, and it has strict requirements:
Those three elements qualify because they are small and highly electronegative, so the H is left almost bare — a proton with very little electron density shielding it, sitting extremely close to the neighbour's lone pair. Chlorine is nearly as electronegative as nitrogen but much bigger, and HCl does not hydrogen bond appreciably.
Note the name is misleading: a hydrogen bond is not a chemical bond. At roughly 10–40 kJ/mol it is about a tenth the strength of a typical covalent bond.
Relative strengths — the ranking table
| Force | Present in | Typical energy (kJ/mol) | Example |
|---|---|---|---|
| London dispersion | All substances | 0.05 – 40 | CH₄, I₂, Xe |
| Dipole–dipole | Polar molecules | 5 – 25 | HCl, CH₃Cl, acetone |
| Hydrogen bonding | H on N/O/F + lone pair | 10 – 40 | H₂O, NH₃, HF, ethanol |
| Ion–dipole | Ion dissolved in a polar solvent | 40 – 600 | Na⁺ surrounded by H₂O |
| Covalent bond (for comparison) | Within a molecule | 150 – 1000 | O–H ≈ 460 |
The ranges overlap deliberately, because the usual exam ranking "dispersion < dipole–dipole < hydrogen bond" is only true when the molecules being compared are of similar size. For large molecules, dispersion wins outright. That is why HI (bp −35 °C) boils higher than the far more polar HCl (bp −85 °C): iodine's huge electron cloud produces more dispersion than HCl's dipole contributes.
Boiling points, explained by the table above
| Hydride series | Period 2 | Period 3 | Period 4 | Period 5 |
|---|---|---|---|---|
| Group 14 (bp °C) | CH₄ −162 | SiH₄ −112 | GeH₄ −89 | SnH₄ −52 |
| Group 15 (bp °C) | NH₃ −33 | PH₃ −88 | AsH₃ −63 | SbH₃ −17 |
| Group 16 (bp °C) | H₂O +100 | H₂S −60 | H₂Se −41 | H₂Te −2 |
| Group 17 (bp °C) | HF +20 | HCl −85 | HBr −67 | HI −35 |
Read the group 14 row first: CH₄ to SnH₄ rises smoothly, because those molecules are non-polar, only dispersion is at work, and dispersion grows with size. That row is the baseline.
Now read the first column. NH₃, H₂O and HF all sit far above where the trend predicts — each is the lightest member of its group yet boils highest, and those three are exactly the hydrides that hydrogen bond. Water is the most extreme because each molecule has two O–H hydrogens and two lone pairs, so it can form up to four hydrogen bonds at once. HF forms stronger individual hydrogen bonds but has only one H per molecule, so it makes fewer of them and boils 80 degrees lower.
Same molar mass, three different forces.
Propane C₃H₈ (44.1 g/mol), non-polar, dispersion only → boils at −42 °C
Dimethyl ether CH₃OCH₃ (46.1 g/mol), polar, dipole–dipole → boils at −25 °C
Ethanol C₂H₅OH (46.1 g/mol), has an O–H, hydrogen bonds → boils at +78 °C
Dimethyl ether and ethanol are isomers — identical atoms, identical molar mass. The only difference is that ethanol's hydrogen sits on the oxygen, and that single structural detail is worth 100 degrees of boiling point.
Same formula, different shape. The three isomers of C₅H₁₂ all have molar mass 72.15 g/mol and all are non-polar:
n-pentane, a straight chain → 36.1 °C
2-methylbutane, one branch → 27.8 °C
neopentane, a compact sphere → 9.5 °C
More branching means a more spherical molecule, less surface contact with its neighbours, weaker dispersion and a lower boiling point. Note that melting point runs the other way — neopentane melts at −17 °C, higher than n-pentane's −130 °C, because a symmetrical sphere packs into a crystal lattice much more neatly.
Why water is strange, in three lines
- Ice floats. Each molecule is locked into four hydrogen bonds in an open tetrahedral network, holding the molecules further apart than in the liquid.
- High specific heat capacity (4.18 J g⁻¹ K⁻¹). Added energy first breaks hydrogen bonds rather than speeding molecules up, so water heats and cools slowly.
- High surface tension. Surface molecules are pulled inward by hydrogen bonds with no outward partners, so the surface behaves like a skin.
- Saying covalent bonds break when a liquid boils. Only intermolecular forces are overcome. This wording alone loses marks in board answers.
- Claiming polar molecules "do not have" dispersion forces. Every molecule has dispersion. Polarity adds a second force on top.
- Calling any H–X…Y interaction a hydrogen bond. The H must be on N, O or F, and the acceptor must be N, O or F. H bonded to carbon does not qualify at this level.
- Forgetting shape when judging polarity. CO₂, CCl₄, BF₃ and CH₄ all contain polar bonds and are all non-polar molecules because of symmetry.
- Applying "hydrogen bond > dipole > dispersion" to molecules of very different size. Compare like with like, or the rule fails — as it does for HI versus HCl.
- Mixing up intramolecular and intermolecular hydrogen bonding. In o-nitrophenol the H is tied up inside the molecule, so it cannot bond to neighbours — which is why o-nitrophenol is far more volatile than p-nitrophenol.
Where intermolecular forces appear in exams
| Exam | Typical use |
|---|---|
| CBSE/ICSE Class 11–12 | Chemical bonding, states of matter, "arrange in order of boiling point" with reasons |
| JEE / NEET | Isomer boiling-point ordering, solubility "like dissolves like", real-gas deviations |
| IIT-JAM / CUET-PG | Anomalous hydride behaviour, van der Waals constants, solvent choice |
| GATE / CSIR-NET | Protein folding and base pairing, chromatography retention, solvent effects on reaction rate |
Every one of these arguments starts with electronegativity. The Interactive Periodic Table lets you check electronegativity, atomic size and group position for any element before you decide whether a bond is polar and whether hydrogen bonding is even possible.
Open the Interactive Periodic Table →Bonding and states of matter are where Class 11 chemistry either clicks or falls apart. ABC Chemistry runs Class 11–12 coaching at its Gurugram centre and online across India, plus IIT-JAM, GATE and CSIR-NET batches — details at abcchemistry.in.