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Methane (CH₄) has a molar mass of 16 g/mol and boils at −162 °C. Water (H₂O) has a molar mass of 18 g/mol — almost the same — and boils at 100 °C. The difference of more than 260 degrees has nothing to do with the bonds inside the molecules. It’s all about the forces between them.
Intramolecular vs. intermolecular
- Intramolecular forces are the bonds holding atoms together inside a molecule — covalent bonds. They’re strong: breaking a C–H bond takes around 400 kJ/mol.
- Intermolecular forces (IMFs) are the attractions between separate molecules. They’re much weaker, typically from under 1 to about 40 kJ/mol.
When water boils, the O–H bonds don’t break. The molecules just pull away from each other. So boiling points, melting points, viscosity and solubility are all governed by intermolecular forces.
1. London dispersion forces
Electrons are always moving. At any instant, a molecule’s electron cloud can be slightly lopsided, making a brief temporary dipole. That momentary charge imbalance distorts the electron cloud of the next molecule, creating an induced dipole, and the two attract. Then the pattern shifts and happens again.
These fleeting attractions are called London dispersion forces, after Fritz London, who explained them in 1930.
- Every molecule and atom has them — even noble gases.
- They’re the only force between nonpolar molecules.
- They get stronger with more electrons, because larger electron clouds are easier to distort (more polarizable).
That’s why, down the halogens, fluorine and chlorine are gases, bromine is a liquid and iodine is a solid at room temperature: bigger molecules, more electrons, stronger dispersion forces. The same trend explains why the noble gases’ boiling points rise steadily from helium to radon.
Shape matters too. Long, straight molecules have more surface contact than compact ones. Pentane (straight chain) boils at 36 °C; its compact isomer neopentane, with the same formula C₅H₁₂, boils at about 10 °C.
2. Dipole-dipole forces
In a polar molecule, one end is permanently slightly negative (δ−) and the other slightly positive (δ+), because of a difference in electronegativity (see polar vs. nonpolar molecules). The positive end of one molecule attracts the negative end of its neighbour.
These dipole-dipole forces act in addition to dispersion forces, so polar molecules usually have higher boiling points than nonpolar molecules of similar size.
Example: propane (C₃H₈, nonpolar, 44 g/mol) boils at −42 °C. Acetaldehyde (CH₃CHO, polar, 44 g/mol) boils at 20 °C.
3. Hydrogen bonding
A hydrogen bond is an especially strong kind of dipole-dipole attraction. It happens when hydrogen is bonded directly to nitrogen, oxygen or fluorine — the three small, highly electronegative atoms. The hydrogen becomes strongly δ+ and is attracted to a lone pair on an N, O or F atom of a neighbouring molecule.
Hydrogen bonding explains the methane-water mystery: water molecules are held together by a network of hydrogen bonds, so boiling it takes far more energy. It also explains why DNA’s two strands zip together and why ice floats. We give it its own article: hydrogen bonding explained.
Ranking the forces
For molecules of similar size:
hydrogen bonding > dipole-dipole > London dispersion
But size can overturn that order. Iodine, I₂, is nonpolar with only dispersion forces, yet it’s a solid at room temperature, while water — with hydrogen bonds — is a liquid. I₂ has 106 electrons per molecule; water has 10. For large molecules, dispersion forces can add up to more than a few hydrogen bonds.
Ion-dipole forces — between an ion and a polar molecule — are stronger still. They’re what lets water dissolve salt: water molecules cluster around Na⁺ and Cl⁻ ions and pull them out of the crystal.
What IMFs control
| Property | Stronger IMFs mean… |
|---|---|
| Boiling and melting point | Higher |
| Viscosity | Thicker, slower-flowing liquid (honey, glycerol) |
| Surface tension | Higher (water striders can walk on water) |
| Vapour pressure | Lower (evaporates more slowly) |
| Heat of vaporization | Higher |
Solubility follows the rule “like dissolves like”. Polar substances dissolve in polar solvents because the new solute–solvent attractions can replace the old ones. Oil (nonpolar) and water (polar, hydrogen-bonded) don’t mix because oil molecules can’t offer water anything as good as the hydrogen bonds it would have to give up.
Quick answers
Are intermolecular forces the same as chemical bonds? No. They’re attractions between molecules, typically much weaker than the covalent bonds inside them.
Which intermolecular force is present in all molecules? London dispersion forces.
Why does a hydrogen bond need N, O or F? Those atoms are small and very electronegative, which leaves the hydrogen with a large partial positive charge and allows a close approach to the lone pair.
Related
See how boiling points vary across the whole periodic table with the boiling point trend map, or explore states of matter.
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