TL;DR:
Boiling a liquid means pulling its molecules apart, so stronger attractions between molecules give a higher boiling point.
From weakest to strongest: dispersion forces, dipole-dipole attractions, hydrogen bonds.
Water (molar mass 18) boils about 260 °C above methane (molar mass 16) because water has hydrogen bonds and methane has only dispersion forces.
Methane and water are made of molecules of nearly the same weight. Yet methane is a gas that boils at about −162 °C, and water is a liquid that does not boil until 100 °C. Something other than size must be holding water molecules together much more tightly. That something is a family of forces between molecules.
Inside the molecule and between molecules
A covalent bond holds atoms together within a molecule, and it is strong. Intermolecular forces are the much weaker attractions between separate molecules. When a liquid boils or a solid melts, the molecules stay intact. Only the attractions between them are overcome.
So the boiling point is a measure of how hard it is to pull molecules apart. Stronger intermolecular forces mean more energy is needed, so a higher boiling point.
The three kinds
Dispersion forces
Every molecule has them. Electrons move around, and at any instant they may bunch up a little on one side, making a brief, tiny dipole. That induces a matching dipole in a neighbour, and the two attract. Larger molecules with more electrons are easier to distort, so their dispersion forces are stronger.
Dipole-dipole attractions
A polar molecule has a permanent positive end and negative end because its shape and bonds are uneven (see predicting molecule shapes with VSEPR). The positive end of one is attracted to the negative end of another.
Hydrogen bonding
A hydrogen bond is a strong kind of dipole attraction. It occurs when hydrogen is bonded to nitrogen, oxygen or fluorine, which are very electronegative. The hydrogen is left with a strong partial positive charge and is attracted to a lone pair on a nearby N, O or F atom. Despite the name, it is an attraction between molecules and is far weaker than a covalent bond.
Worked example: water versus methane
Molar masses: water
2 × 1.008 + 16.00 = 18.02 g/mol, methane12.01 + 4 × 1.008 = 16.04 g/mol. They are nearly the same, so size cannot explain the gap.Methane: carbon and hydrogen have a small electronegativity difference (0.35), and the four bonds are arranged symmetrically in a tetrahedron, so the molecule is nonpolar. Only dispersion forces act.
Water: the O-H bonds are polar and the molecule is bent, so it is polar. Each hydrogen is bonded to oxygen, so water forms hydrogen bonds, and each molecule can make several of them.
Boiling points: water 100 °C, methane about −162 °C. The difference is
100 − (−162) = 262 °C.
Similar mass, very different attraction. The extra energy needed to break water's hydrogen bonds is the reason a pan of water must reach 100 °C before it boils, while methane vaporises at a temperature colder than any freezer.
A second check shows dispersion forces growing with size, among nonpolar molecules. Ethane (30 g/mol) boils at about −89 °C, propane (44 g/mol) at about −42 °C and butane (58 g/mol) at about −1 °C. As molecules get heavier with more electrons, the dispersion attractions strengthen and the boiling point climbs.
Boiling and flammable liquids: Do not heat flammable liquids such as alcohols or fuels over an open flame, and never heat a closed container. Boiling water also scalds, so keep pots and kettles out of reach of children.
Hydrogen bonding explains many familiar facts about water. It is liquid at room temperature when much heavier molecules of similar type are gases. It takes a lot of energy to heat, because some of that energy goes into disturbing hydrogen bonds. The hydrogen bonds also hold water molecules together at the surface, which is why a small insect can walk on a pond.
Remember that intermolecular forces also depend on contact area. Two molecules with the same mass can have different boiling points if one is long and thin, so it touches its neighbours along more of its length, and the other is compact and round. The long one has the higher boiling point.
How to check your answer
Ladder. Rank forces: hydrogen bonding is stronger than dipole-dipole, which is generally stronger than dispersion of molecules of similar size.
Compare size first. If molecules have very different masses, dispersion forces may outweigh a small dipole.
Check for O-H, N-H or F-H. Their presence suggests hydrogen bonding.
Shape matters. A molecule with polar bonds can still be nonpolar if they cancel, as in CO₂.
Mistakes that cost marks
Saying boiling breaks the covalent bonds in the molecule.
Calling a hydrogen bond a normal covalent bond, or claiming any hydrogen counts. It must be on N, O or F.
Forgetting that all molecules have dispersion forces, including polar ones.
Ignoring molecule size when comparing.
Saying a polar bond always means a polar molecule.
Where you will use it next
The same forces explain why oil and water do not mix, why ice floats and how proteins and DNA hold their shape. They also decide which substances dissolve, as in solubility rules and precipitates, and connect to the bonding ideas in ionic and covalent bonds.
Frequently asked questions
Why does ice float?
In ice, hydrogen bonds hold water molecules in an open, regular pattern that takes up more room than in the liquid, so ice is less dense than liquid water.
Are intermolecular forces the same as bonds?
No. Bonds hold atoms inside a molecule. Intermolecular forces are the weaker attractions between separate molecules.
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