Intermolecular forces: why water boils at 100 °C and methane at −161 °C

Attraction without a bond

If particles did not attract one another at all, every substance would be a gas, with no liquids or solids. Why water and methane boil so far apart is therefore a question about those attractions and how well they stand up to heat.

When two atoms attract strongly enough to form an enduring unit, the unit is a molecule and the attraction is a chemical bond, acting inside it: an intramolecular force. When particles attract too weakly to be bound into a new unit, the attraction is a non-bonding one, called an intermolecular force. Its particles are usually molecules, but they can also be single atoms.

Heat works against the intermolecular force. In a gas, thermal energy keeps disrupting it and the particles stay apart. In a liquid, with the particles crowded close together, the same attraction has an effect that is easy to see. The attraction between the ions of a lattice is a different matter: it belongs with the bonds, and everything below concerns substances made of molecules or of single atoms.

Dispersion forces: a momentary dipole that averages to zero

The noble gases and completely nonpolar molecules such as H₂ and N₂ still condense to liquids. Some attraction between particles must therefore exist that needs no permanent dipole moment.

The answer lies in what “nonpolar” means. A nonpolar particle is one whose dipole moment is zero on average over time. At any instant, nothing guarantees that the negative charge around an atom’s centre is spread perfectly evenly. For that instant, the atom is a small dipole.

These lopsided moments are very short-lived, and over time they cancel to zero. While one lasts, though, it can induce a new dipole in a neighbouring atom or molecule. The induced dipole in turn helps to sustain the first. The result is an attraction between two particles with no permanent charge separation at all: the dispersion force, also called the London force, a spontaneous dipole–induced dipole interaction.

So the dipole is real at every instant and zero on average. The force it produces is weak for any one pair of particles, yet acts in every substance. Its strength depends on how easily an electron cloud is pushed out of shape, a property called polarisability: the size of the temporary dipole a nearby charge can induce. The larger an atom, the more loosely it holds its outer electrons, and the more readily its cloud distorts.

The noble gases: more electrons down the group

The noble gases show the dispersion force with nothing else beside it. Their ability to condense to liquids and to freeze to solids depends entirely on dispersion forces. Argon and helium are known as single atoms: quantum mechanics allows argon no stable Ar₂ molecule, and He₂ has never been observed. For a single atom there is no shape and no polarity to set aside, only the atom itself.

At ordinary temperatures, the attraction between argon atoms cannot withstand the disruption of thermal energy, so argon is a gas. Argon’s figure is one atom, its 18 electrons in three shells, with no second atom bonded to it.

Down the group the electron count climbs: 2 in helium, 10 in neon, 18 in argon, 36 in krypton, 54 in xenon and 86 in radon. The boiling points climb with it, to the nearest degree −269, −246, −186, −153, −108 and −62 °C. The strength of the dispersion force depends largely on the number of electrons. Each row down the group adds an electron shell, so each atom is larger than the one above and holds its outer electrons more loosely.

The halogens: two identical atoms, and the same rule

Fluorine, chlorine, bromine and iodine exist as molecules of two atoms: F₂, Cl₂, Br₂ and I₂. Between two identical atoms the sharing of electrons must be exactly even. None of these molecules therefore has a permanent dipole, and none can take part in a dipole-dipole force, the attraction between two permanent dipoles. What remains to hold their liquids and solids together is the dispersion force, the same attraction that condenses the noble gases, now acting between molecules instead of single atoms.

Their atoms carry 9, 17, 35 and 53 electrons. To the nearest degree, the four boil at −188, −34, 59 and 184 °C. At room temperature fluorine and chlorine are gases, bromine is a liquid and iodine is a solid: three states of matter within one group. The rule that ordered the noble gases orders the halogens too. Each larger atom holds its outer electrons more loosely, the dispersion force grows, and by iodine it holds a solid together at room temperature.

Permanent dipoles, and the hydrogen bond

A polar molecule carries a permanent dipole, and two such molecules attract in a second way. As two permanent dipoles approach, they tend to turn so that the positive end of one lies next to the negative end of the other. This dipole-dipole force can be significant when the dipole moments are large.

The hydrogen bond needs a particular arrangement of atoms. A hydrogen atom bonded to one of the more electronegative atoms has its electron cloud distorted by that atom, leaving it a partial positive charge. The hydrogen atom is so small that this charge is dense enough to interact with a lone pair on a nearby electronegative atom of another molecule. That attraction between molecules is the hydrogen bond.

It involves a measure of electron sharing, so it has some covalent character and is not simply a dipole-dipole force. In water it is about 23 kJ/mol, against 492 kJ/mol for the covalent O–H bond inside each molecule.

Ammonia is polar by measurement, with a gas-phase dipole moment of 1.47 D. Its hydrogens sit on nitrogen, which carries one lone pair, and liquid and solid ammonia are hydrogen-bonded. Even so, it boils at −33.35 °C.

Water against methane

Liquid methane, CH₄, with a molar mass of 16 g/mol, boils at −161 °C. All four of its hydrogens are attached to carbon, which is not usually considered particularly electronegative. Methane’s figure shows its five atoms, with every hydrogen bonded to the one carbon at the centre.

Water, H₂O, has a molar mass of 18.015 g/mol, close to methane’s. Extending the boiling points of the other group 16 hydrogen compounds to H₂O suggests it should be a gas under normal conditions, yet it boils at 99.974 °C. Its hydrogens sit on oxygen, one of the more electronegative atoms, and the oxygen carries two lone pairs. Each water molecule can therefore hydrogen-bond to its neighbours. Water’s figure shows three atoms joined by two O–H bonds, with both hydrogens on the oxygen.

In steam, ordinary thermal motion swamps the attraction between water molecules completely. Once the molecules are packed close together in the liquid, the same attraction has a marked effect, and that effect is what the name hydrogen bonding describes.

Water and ammonia are the lightest molecules in their series for which hydrogen bonding is possible, and both have anomalously high boiling points. Water and methane both have dispersion forces. Water’s hydrogens sit on oxygen, where the conditions for the hydrogen bond are met, and methane’s sit on carbon. Their molar masses differ by 2 g/mol, and their boiling points by about 261 degrees.