Intermolecular Force Myths: What Boiling Really Breaks

Ask a class why ethanol boils at a lower temperature than water and you’ll hear a range of confident answers. Some are right, many are half right, and a few rest on a picture of matter that simply isn’t true. Intermolecular forces are one of the topics where students often learn the vocabulary (London, dipole–dipole, hydrogen bonding) without the model underneath, so the wrong ideas survive every test that only asks for definitions.

This post takes eight of those ideas apart. For each one you’ll see the myth, why it’s so easy to believe, and what’s actually going on, with a real example you can check. If you want the general background first, read intermolecular forces; for bonding myths more broadly (energy release, full shells and so on), see chemical bonding misconceptions.

Myth 1: “Boiling breaks the bonds in the molecules”

Why people believe it. Everyday language blurs the line. We say water “breaks down” when heated, and diagrams of boiling often show loose atoms floating upwards. The word “bond” is also used for both the strong links inside a molecule and the weaker attractions between molecules, so it feels natural to assume boiling attacks the same thing.

What’s really happening. Boiling separates whole molecules from one another. The covalent bonds inside each molecule stay intact. Steam is made of H₂O molecules, exactly like liquid water; if the O–H bonds had broken, you’d get hydrogen and oxygen, and cooling the vapour would not give water back.

The energies make the point clearly. Breaking one mole of O–H bonds takes roughly 460 kJ, so pulling both hydrogens off every water molecule in a mole would cost more than 900 kJ. Vaporising a mole of water at 100 °C takes only about 41 kJ. That smaller figure is the cost of overcoming the hydrogen bonds and other attractions between molecules. The two processes are more than twenty times apart. The same distinction is laid out side by side in intramolecular vs intermolecular forces.

Myth 2: “Every substance has intermolecular forces”

Why people believe it. Textbooks often introduce states of matter with simple molecular substances, so students generalise: melting and boiling must always be about forces between molecules.

What’s really happening. Intermolecular forces only make sense when there are separate molecules. Sodium chloride has no NaCl molecules at all; it’s a giant lattice of Na⁺ and Cl⁻ ions, and melting it means overcoming ionic bonding throughout that lattice. Diamond and silicon dioxide are giant covalent structures, so melting them means breaking covalent bonds. Metals are held by metallic bonding.

That’s why the melting points are so different. Sodium chloride melts at about 801 °C, while iodine, a simple molecular solid, melts at about 114 °C. Before you talk about London forces or dipoles, check the structure type. If the substance is a giant structure, intermolecular forces are the wrong tool.

Myth 3: “London forces only exist in non-polar molecules”

Why people believe it. London (dispersion) forces are usually taught with examples like the noble gases, the halogens and hydrocarbons. Polar molecules are then introduced with dipole–dipole forces, as though they replace London forces.

What’s really happening. Every atom and molecule has London forces, because every one has an electron cloud that can be momentarily distorted. A fluctuation in one molecule induces a matching distortion in its neighbour, and the two attract. Polar molecules have London forces and permanent dipole attractions. Molecules that can hydrogen bond have all three.

This matters because students who think polar molecules lack London forces can’t explain trends like the one in the next myth. For a full treatment of how temporary dipoles arise, see London dispersion forces.

Myth 4: “Dipole–dipole forces are always stronger than London forces”

Why people believe it. Lists in textbooks often go “London < dipole–dipole < hydrogen bonding”, and it’s easy to read that as a rule for any pair of molecules rather than a rough guide for molecules of similar size.

What’s really happening. In small molecules, London forces can easily be the larger contribution, and in large molecules they usually dominate. The hydrogen halides are the classic evidence:

Compound	Boiling point	Polarity of the H–X bond
HCl	about −85 °C	most polar of the three
HBr	about −67 °C	less polar
HI	about −35 °C	least polar

If dipole–dipole forces controlled the order, HCl would boil highest. Instead, the boiling point rises as the halogen gets bigger. Iodine has far more electrons than chlorine and a larger, more easily distorted electron cloud, so London forces grow faster than the dipole shrinks. The permanent dipole is real, but it isn’t the main reason HI is the least volatile. You can compare the halogens themselves on the chlorine and iodine pages.

Myth 5: “A molecule that can hydrogen bond always boils higher than one that can’t”

Why people believe it. Hydrogen bonding is taught as the strongest intermolecular force, and water’s unusually high boiling point is the headline example. The shortcut “hydrogen bonding means high boiling point” follows.

What’s really happening. Hydrogen bonding is strong per interaction, but the total attraction between molecules depends on all the forces added together. Compare methanol and hexane:

• Methanol, CH₃OH, can hydrogen bond. It boils at about 65 °C.
• Hexane, C₆H₁₄, is non-polar and relies on London forces alone. It boils at about 69 °C.

Hexane has fifty electrons to methanol’s eighteen, and its long chain gives a large contact area between neighbours. Those London forces add up to slightly more than methanol’s mix of hydrogen bonds and weaker forces. Long-chain waxes are an even more extreme case: purely non-polar, yet solid at room temperature. Always weigh the size of the molecule, not just the “best” force it can form. The article on boiling points and intermolecular forces walks through more comparisons.

Myth 6: “Heavier molecules always have higher boiling points”

Why people believe it. Within a homologous series, such as the straight-chain alkanes, boiling point does rise smoothly with molar mass. Students then treat mass itself as the cause.

What’s really happening. Mass is a stand-in for the real factors: the number of electrons (which sets how polarisable the cloud is) and the shape of the molecule (which sets how much surface two molecules can share). When mass is held constant and shape changes, the boiling points spread out. The three isomers of C₅H₁₂ all have the same molar mass:

• Pentane (a straight chain) boils at about 36 °C.
• 2-Methylbutane (one branch) boils at about 28 °C.
• 2,2-Dimethylpropane (nearly spherical) boils at about 10 °C.

A straight chain can lie alongside its neighbours over its whole length, like pencils in a box. A compact, ball-shaped molecule touches its neighbours at only a few points, so the total London attraction is smaller. Same mass, same electrons, a range of over 25 °C.

Myth 7: “If a molecule has polar bonds, it’s a polar molecule”

Why people believe it. Bond polarity is taught first, using electronegativity differences, and the idea that “polar bonds make polar molecules” seems like a natural next step.

What’s really happening. A molecule’s overall polarity depends on how its bond dipoles add up in three dimensions. Carbon dioxide has two strongly polar C=O bonds, but the molecule is linear, so the two dipoles point in opposite directions and cancel. CO₂ has no overall dipole and only London forces between its molecules, which is part of why it’s a gas at room temperature. Tetrachloromethane, CCl₄, is similar: four polar C–Cl bonds arranged tetrahedrally, with zero net dipole.

Water, by contrast, is bent, so its two O–H dipoles don’t cancel. To judge polarity, you need the shape first; VSEPR and molecular geometry gives you the method, and dipole moments shows how the vectors combine.

Myth 8: “Intermolecular forces only matter for boiling points”

Why people believe it. Exam questions on this topic very often ask students to explain a boiling point, so the link between intermolecular forces and boiling becomes the whole story.

What’s really happening. The same forces control many other properties:

• Solubility. Substances tend to dissolve where the new attractions to the solvent can compensate for the attractions being overcome. That’s why ethanol mixes with water but hexane doesn’t.
• Surface tension. Water beads up on a waxed surface because molecules at the surface are pulled inwards by strong attractions to the molecules below.
• Viscosity. Glycerol, with three –OH groups per molecule, flows far more slowly than water.
• Physical state of solids. Ice, iodine crystals and candle wax are all held together by intermolecular forces.
• Biology. The shapes of proteins and the pairing of DNA bases depend heavily on hydrogen bonds and other weak interactions.

Thinking of intermolecular forces as “the boiling point topic” makes each of these feel like a separate fact to memorise, when they’re one idea applied in different places.

A quick way to reason about any comparison

When a question asks you to compare two substances, work through these steps in order:

1. Structure type. Simple molecular, giant covalent, ionic or metallic? Only simple molecular substances are explained by intermolecular forces.
2. Size. Count electrons (or compare molar masses within similar molecules). More electrons means stronger London forces.
3. Shape. Long and flat molecules pack closely; compact ones don’t.
4. Polarity. Use the shape to decide whether bond dipoles cancel.
5. Hydrogen bonding. Is there H on N, O or F, and a lone pair on N, O or F to accept it?

Then decide which factor dominates, and say so explicitly. Examiners reward the phrase “more energy is needed to overcome the stronger forces between molecules”, and they penalise any hint that bonds inside the molecules break. The intermolecular forces practice questions are a good place to test yourself.

Key takeaways

• Boiling and melting a simple molecular substance separate molecules; covalent bonds inside them stay intact.
• Giant ionic, covalent and metallic structures have no intermolecular forces to speak of; melting them breaks strong bonds throughout.
• London forces act between all molecules and often outweigh dipole–dipole forces, as the HCl–HBr–HI trend shows.
• Hydrogen bonding doesn’t automatically beat a larger non-polar molecule: hexane boils above methanol.
• Mass is a proxy. Electron count and molecular shape are what actually set London force strength.
• Polar bonds don’t guarantee a polar molecule; symmetry can cancel them.
• The same forces explain solubility, surface tension, viscosity and much of biochemistry, not just boiling points.