Boiling Point vs Melting Point: What's the Difference?
Benzene melts at 5.5 °C and boils at 80.1 °C. Toluene — the same ring with a methyl group stuck on — boils higher, at 110.6 °C, but melts at a spectacular −95 °C. One molecule is bigger and holds on more tightly, yet it falls apart as a solid 100 degrees sooner. If melting and boiling measured the same thing, that could not happen.
The short answer: the melting point is the temperature at which a solid becomes a liquid, and the boiling point is the temperature at which a liquid becomes a gas throughout its bulk — when its vapour pressure equals the surrounding pressure. Melting only loosens a structure enough to let particles slide; boiling has to pull them apart completely, which is why boiling point is a much better guide to intermolecular force strength.
Quick comparison at a glance
| Feature | Melting point | Boiling point |
|---|---|---|
| Phase change | Solid → liquid | Liquid → gas |
| What must be overcome | Enough of the lattice to let particles move past each other | Essentially all the attractions between particles |
| Physical condition | Solid and liquid in equilibrium | Vapour pressure equals external pressure |
| Also called | The freezing point — same temperature, opposite direction | The condensation point, in reverse |
| Depends on | Intermolecular forces and how well the molecules pack | Intermolecular forces, almost alone |
| Pressure sensitivity | Very slight | Strong — changes with altitude |
| Energy absorbed | Enthalpy of fusion (water: 6.0 kJ/mol) | Enthalpy of vaporisation (water: 40.7 kJ/mol) |
| Water at 1 atm | 0 °C | 100 °C |
Look at those two energies. Vaporising water costs nearly seven times what melting it does — a direct measure of how much more work is involved in separating particles completely than in merely letting them slide.
What happens at the melting point
In a solid, particles are locked into fixed positions, vibrating in place. Heat them and the vibrations grow until the particles have enough energy to break out of their lattice positions and move past one another. The solid melts.
Crucially, they do not move far apart. A liquid is typically only about 5–10 % less dense than its solid, and the intermolecular forces are still very much doing their job. Melting loosens the arrangement; it does not dismantle it.
That's why melting point depends on two things: how strong the forces are, and how neatly the molecules stack into a crystal. A symmetrical molecule slots into a tidy, efficient lattice that takes a lot of energy to disrupt. An awkwardly shaped one packs badly and its lattice gives way early.
Benzene and toluene are exactly this. Benzene is a flat, perfectly symmetrical hexagon that packs beautifully — so it melts at a respectable 5.5 °C. Toluene's protruding methyl group ruins that symmetry, the crystal is poor, and it melts at −95 °C despite having stronger overall attractions.
What happens at the boiling point
Boiling is a much more demanding transition. To become a gas, a particle has to escape its neighbours entirely — the attractions must be overcome, not just loosened. There is no packing shortcut, because a gas has no structure to pack into.
A liquid boils when its vapour pressure equals the external pressure. Below that temperature, any bubble that forms inside the liquid is immediately crushed; at that temperature, bubbles survive and rise, and you see boiling.
Because nothing but the strength of the attractions is being tested, boiling point tracks intermolecular forces cleanly:
| Substance | Dominant forces | Boiling point |
|---|---|---|
| CH₄ | Dispersion only, tiny molecule | −161 °C |
| HCl | Dispersion + dipole-dipole | −85 °C |
| H₂O | Extensive hydrogen bonding | 100 °C |
| NaCl | Full ionic bonding | 1465 °C |
That last row is a different league because ionic compounds have no discrete molecules — the whole crystal is one electrostatic network (NaCl melts at 801 °C). Whenever you see melting or boiling points running to many hundreds or thousands of degrees, you are looking at ionic, metallic, or covalent-network bonding rather than intermolecular forces — nothing held together by intermolecular forces alone survives that kind of heat.
How to tell them apart in practice
On a heating curve, both show up as flat plateaus — temperature stops rising while the added heat goes into breaking attractions rather than speeding particles up. The first, shorter plateau is melting; the second, much longer one is boiling.
In the lab, they serve different purposes:
- Melting point is the purity test. A pure crystalline solid melts sharply over about 1 °C. Impurities disrupt the lattice, so the sample melts lower and over a wider range — often several degrees. Chemists identify unknown solids this way.
- Boiling point is the separation tool. Distillation exploits boiling-point differences to separate liquids, and a nonvolatile solute raises a liquid's boiling point rather than lowering it.
Note the asymmetry, because it is a favourite exam question: dissolved salt lowers the melting point of ice (which is why roads are gritted) and raises the boiling point of water. Both follow from the same principle — the solute lowers the vapour pressure of the liquid phase.
Worked examples
1. Which is higher for a given substance, melting point or boiling point? Boiling point, always — separating particles completely takes more energy than loosening them.
2. A solid melts at 121–122 °C; your sample melts at 114–119 °C. What does that tell you? It's impure. Impurities depress the melting point and broaden the range.
3. Why does ethanol (bp 78 °C) boil below water (bp 100 °C) despite both having –OH groups? Water forms up to four hydrogen bonds per molecule; ethanol's bulky ethyl group allows far fewer, so less energy is needed to separate the molecules.
4. Why does dry ice have no normal melting point? Solid CO₂ sublimes — it goes straight from solid to gas at −78.5 °C at 1 atm. It only melts under pressures above about 5 atm.
Common mistakes to avoid
- Assuming melting point predicts boiling point. Toluene beats benzene on boiling and loses badly on melting. Melting point is hostage to crystal packing and symmetry; boiling point is not.
- Thinking the temperature climbs during a phase change. It doesn't. On the plateaus, every joule goes into breaking attractions, not into raising the average kinetic energy.
- Confusing melting with dissolving. Melting is heat turning a solid into its own liquid. Dissolving is a solid dispersing into a different substance, the solvent — sugar in cold tea is dissolving, not melting.
FAQ
What is the difference between melting point and boiling point?
Melting point is the temperature at which a solid turns into a liquid; boiling point is the temperature at which a liquid turns into a gas throughout its bulk, when its vapour pressure equals the external pressure.
Is the boiling point always higher than the melting point?
For a given substance at the same pressure, yes. Some substances, such as carbon dioxide at 1 atm, sublime instead and have neither a normal melting point nor a normal boiling point at that pressure.
Does melting point change with pressure?
Only very slightly — a solid and a liquid have similar volumes, so pressure barely shifts the balance. Boiling point changes a great deal, because a gas occupies vastly more volume than the liquid it came from.
Why do impurities lower the melting point?
Foreign particles disrupt the regular crystal lattice, so less energy is needed to break it apart. They also broaden the melting range, which is why a sharp melting point indicates a pure sample.
The takeaway
Melting loosens a structure; boiling destroys it. That is why boiling point is the cleaner measure of how strongly molecules attract one another, while melting point carries extra information about how neatly they pack into a crystal — and why two similar molecules can rank one way on boiling and the opposite way on melting.
Round out the topic with What Is a Phase Change? Melting, Boiling, Sublimation, What Are the States of Matter? Solid, Liquid, Gas, and Ionic vs Covalent Bonds: What's the Difference?. This one closes the intermolecular-forces series that began with [Dipole-Dipole vs London Dispersion Forces].
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