Chlorine gas (Cl₂) boils at -34°C, while hydrogen chloride (HCl) boils at -85°C, a difference of about 51 degrees. This surprises many chemistry students because HCl is a polar molecule with dipole-dipole attractions, while Cl₂ is completely nonpolar. The explanation comes down to one thing: Cl₂ has far more electrons than HCl, giving it stronger London dispersion forces that outweigh HCl’s advantage in polarity.
The Intermolecular Forces in Each Molecule
HCl is a polar molecule with a dipole moment of 1.109 Debye. Chlorine pulls electron density away from hydrogen, creating a partially negative end (Cl) and a partially positive end (H). This means HCl molecules attract each other through dipole-dipole forces, where the positive end of one molecule lines up with the negative end of another. HCl also experiences London dispersion forces, as all molecules do.
Cl₂, on the other hand, is completely nonpolar. Two identical chlorine atoms share electrons equally, so there’s no permanent charge separation and no dipole-dipole attraction. The only intermolecular force holding Cl₂ molecules together is London dispersion forces.
At first glance, this makes it seem like HCl should have the higher boiling point. It has dipole-dipole forces on top of dispersion forces, while Cl₂ has dispersion forces alone. But the strength of those dispersion forces depends heavily on something else: how many electrons the molecule has.
Why Electron Count Matters So Much
London dispersion forces arise from temporary, fleeting imbalances in how electrons are distributed around a molecule. At any given instant, the electrons might cluster slightly toward one side, creating a momentary dipole. That temporary dipole induces a dipole in a neighboring molecule, and the two attract each other briefly. This happens constantly and rapidly across all the molecules in a substance.
The key factor is how easily a molecule’s electron cloud can be distorted, a property called polarizability. Bigger electron clouds with more electrons are easier to distort, which creates stronger and more frequent temporary dipoles. Cl₂ has 34 electrons (17 from each chlorine atom), while HCl has only 18 (17 from chlorine and 1 from hydrogen). That’s nearly twice as many electrons in Cl₂, spread across a larger molecular surface. The result is significantly stronger London dispersion forces between Cl₂ molecules.
A helpful comparison from Khan Academy illustrates the pattern: bromine (Br₂), which has even more electrons than Cl₂, boils at 59°C, far higher than chlorine’s -34°C. The trend holds consistently. More electrons mean stronger dispersion forces and higher boiling points.
When Dispersion Forces Beat Dipole-Dipole Forces
There’s a common misconception that dipole-dipole forces are always “stronger” than London dispersion forces. Textbooks sometimes list London dispersion forces as the “weakest” intermolecular force, and that’s true on a per-interaction basis for small molecules. But the total strength of dispersion forces in a substance scales with molecular size and electron count. For a large, electron-rich molecule, the cumulative dispersion forces can easily surpass the dipole-dipole attractions in a smaller polar molecule.
That’s exactly what happens here. HCl’s dipole-dipole interactions add some attraction between molecules, but HCl is a small, lightweight molecule with relatively few electrons. Its total intermolecular attraction, combining both dipole-dipole and dispersion forces, is still weaker than the dispersion-only attraction between the larger, more electron-rich Cl₂ molecules. The 51°C gap in their boiling points reflects that difference directly.
Molecular Size and Surface Area
Electron count isn’t the only contributor. The physical size and shape of a molecule also affect how strongly London dispersion forces operate. Longer molecules with more surface area have more points of contact with neighboring molecules, which means more opportunities for temporary dipoles to form and interact.
Cl₂ is a larger molecule than HCl. The bond length in HCl is about 1.27 angstroms, and hydrogen contributes almost no size to the molecule. Cl₂, with two full-sized chlorine atoms bonded together, presents a bigger target. That extra surface area gives neighboring Cl₂ molecules more contact, reinforcing the already-strong dispersion forces from its 34 electrons.
The General Principle
This comparison between Cl₂ and HCl is a textbook example of a broader rule in chemistry: the type of intermolecular force matters, but so does the magnitude. A nonpolar molecule with many electrons can easily have a higher boiling point than a polar molecule with fewer electrons. Polarity gives a molecule an extra source of attraction, but it doesn’t guarantee that source will be the dominant one.
You can see this pattern across many pairs of molecules. Nonpolar hydrocarbons like octane (boiling point 125°C) boil far above small polar molecules like formaldehyde (boiling point -19°C), for the same reason. The sheer number of electrons and the molecular surface area available for dispersion interactions overwhelm the relatively modest contribution of a permanent dipole. In the Cl₂ versus HCl comparison, the numbers just happen to make the lesson especially clear.