Intermolecular forces are the attractions between molecules — they are weaker than covalent or ionic bonds but they determine melting points, boiling points, and whether a substance is a solid, liquid, or gas at room temperature. Larger molecules have stronger intermolecular forces, which is why longer-chain alkanes have higher boiling points than shorter ones.
What are intermolecular forces?
Intermolecular forces are the weak electrostatic attractions between separate molecules. They must not be confused with the covalent bonds within a molecule, which are strong and are not broken during melting or boiling. When a molecular substance melts or boils, it is the intermolecular forces — not the covalent bonds — that are overcome.
This distinction is crucial for GCSE chemistry. Water molecules, for instance, contain very strong O–H covalent bonds (464 kJ/mol), but water boils at only 100 °C because only the intermolecular attractions between water molecules need to be overcome to vaporise it.
What type of intermolecular forces appear in GCSE chemistry?
At GCSE level, the main intermolecular force you need to know is the van der Waals force (also called London dispersion forces or temporary dipole–induced dipole interactions). These arise as follows:
- Electrons in a molecule are in constant random motion. At any instant, they may be distributed unevenly — creating a momentary, temporary dipole (one end slightly negative, the other slightly positive).
- This temporary dipole induces a dipole in a neighbouring molecule — attracting or repelling the electrons there.
- The induced dipole then attracts the temporary dipole in the first molecule, creating a weak but real intermolecular attraction.
These van der Waals forces exist between all molecules, whether polar or non-polar.
How does molecular size affect intermolecular forces?
The strength of van der Waals forces depends on the number of electrons in a molecule. More electrons mean:
- More "surface area" for temporary dipoles to form across
- Larger temporary dipoles when they occur
- Stronger attractions between neighbouring molecules
This explains a crucial GCSE trend: as the number of carbons in a chain of alkanes increases, the boiling point increases.
| Alkane | Formula | Number of electrons | Boiling point (°C) |
|---|---|---|---|
| Methane | CH₄ | 10 | −162 |
| Ethane | C₂H₆ | 18 | −89 |
| Propane | C₃H₈ | 26 | −42 |
| Butane | C₄H₁₀ | 34 | −1 |
| Pentane | C₅H₁₂ | 42 | +36 |
As the chain length increases, the number of electrons increases, van der Waals forces become stronger, more energy is needed to separate the molecules, and the boiling point rises.
This also explains why the first four alkanes (methane to butane) are gases at room temperature while longer-chain alkanes are liquids or solids — their intermolecular forces are strong enough to hold them together.
Why do simple molecular substances have low melting and boiling points?
Simple molecular substances (e.g. H₂O, CO₂, O₂, HCl, simple organic molecules) have low melting and boiling points compared to giant structures such as ionic compounds or giant covalent structures. The reason is that the atoms within each molecule are held together by strong covalent bonds, but the separate molecules are held to each other only by weak intermolecular forces.
To melt or boil the substance, you only need to supply enough energy to overcome the intermolecular forces — you do not need to break the covalent bonds. Because intermolecular forces are weak, a relatively small amount of energy is needed, so melting and boiling points are low.
By contrast, sodium chloride (ionic) melts at 801 °C because you must break the strong electrostatic attractions between ions throughout the entire giant lattice.
How does molecular shape affect intermolecular forces?
Straight-chain molecules can pack closely together, maximising the contact area between adjacent molecules. This allows more van der Waals interactions across a larger surface area. Branched molecules of the same formula (same number of electrons) are more compact and spherical; they pack less efficiently and have less contact area between neighbours.
As a result, branched isomers have lower boiling points than their straight-chain equivalents.
Example: pentane (straight chain, bp +36 °C) vs 2,2-dimethylpropane (neopentane, highly branched, bp +9 °C) — same molecular formula C₅H₁₂, same number of electrons, but the branched isomer has weaker overall intermolecular forces.
Why does water have an unusually high boiling point for its size?
Water (H₂O) boils at 100 °C — far higher than you would predict purely from van der Waals forces for such a small molecule. Hydrogen sulfide (H₂S), a much larger molecule, boils at −60 °C.
The reason is hydrogen bonding — a particularly strong type of intermolecular attraction that forms between the δ+ hydrogen on one water molecule and the lone pair on the δ− oxygen of a neighbouring water molecule. Hydrogen bonds are roughly 20 times stronger than typical van der Waals forces, so significantly more energy is needed to separate water molecules, giving water its anomalously high boiling point.
GCSE chemistry mentions hydrogen bonding in the context of water without requiring you to know the full theory; the key fact is that water's unexpectedly high boiling point is caused by unusually strong intermolecular forces.
Frequently asked questions
What is the difference between intramolecular bonds and intermolecular forces?
Intramolecular bonds (such as covalent bonds) are the strong attractions within a molecule between individual atoms. Intermolecular forces are the much weaker attractions between separate molecules. When a molecular substance melts or boils, only intermolecular forces are overcome — the covalent bonds inside each molecule stay intact. This is why melting or boiling does not change the chemical identity of a simple molecular substance.
Why does CO₂ not conduct electricity even when melted?
CO₂ is a simple molecular substance. Its molecules carry no overall charge, so even in the liquid state there are no free-moving charged particles to carry an electrical current. This is a general property of molecular substances — they do not conduct electricity in any state. This contrasts with ionic compounds, which conduct when molten or dissolved because ions become free to move.
How do you explain the trend in boiling points of the noble gases?
The noble gases — helium, neon, argon, krypton, xenon — are monatomic and have no permanent dipoles. Their only intermolecular forces are van der Waals forces. As you go down Group 0, atomic (and electron) number increases, van der Waals forces strengthen, and boiling points increase: He (−269 °C), Ne (−246 °C), Ar (−186 °C), Kr (−153 °C), Xe (−108 °C). This is a clean demonstration that van der Waals force strength increases with electron count.
Does a higher boiling point always mean stronger intermolecular forces?
Yes — within a class of substances with the same type of intermolecular force. A higher boiling point means more energy is required to separate the molecules, which directly reflects stronger attractions between them. The comparison becomes more complex when different types of intermolecular force are at play (e.g. comparing a large non-polar molecule with a small polar one), but the principle holds within any series of structurally similar compounds.
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