Intermolecular Forces — Print · Grade 11 Science

Intermolecular Forces

Learning goals

• I can identify the intermolecular forces present in a substance: London dispersion, dipole-dipole, and hydrogen bonding.
• I can use the relative strength of intermolecular forces to compare boiling points, melting points, and related properties.

Readings for this unit

• One warmed sample, several possible responses

There are two very different kinds of 'sticking together' in chemistry, and confusing them causes no end of trouble. Inside a molecule, atoms are held by strong chemical bonds — the covalent bonds that make an H₂O molecule, for instance. Between one molecule and its neighbors act much weaker attractions called intermolecular forces (IMFs). Weak though they are, IMFs are the quiet decision-makers of the physical world: they determine whether a substance is a gas, liquid, or solid at room temperature, how high its boiling and melting points are, how thick and syrupy it is, and how strong its surface tension is. When a pot of water boils, it is these IMFs breaking apart — not the sturdy O–H bonds, which survive untouched.

A molecular interaction present even without permanent polarity is London dispersion forces, and here is their surprising feature — all molecular species can exhibit dispersion, including nonpolar molecules without a permanent net dipole. At any instant the electrons in a molecule can slosh slightly to one side, creating a fleeting, temporary dipole that induces a matching dipole in a neighbor, and for a moment they attract. The more electrons a molecule has, and the larger and more polarizable its electron cloud, the stronger these forces grow. That single trend explains why, going down the halogen column, F₂ and Cl₂ are gases, Br₂ is a liquid, and I₂ is a solid at room temperature: same type of force, steadily more of it.

Dipole-dipole forces appear in polar molecules, which carry a permanent positive end and a permanent negative end because their electrons are shared unevenly. These molecules line up so that opposite ends attract, through electric charge-distribution interactions, and for molecules of similar size these contributions add to dispersion; their relative importance depends on the species and orientation. It is why a polar molecule generally boils at a higher temperature than a nonpolar molecule of comparable mass — there is an extra electrical grip to overcome before the molecules can fly apart into a gas.

The standout of the three is hydrogen bonding, a particularly important interaction with a specific donor/acceptor context in common introductory examples. For the usual school N/O/F examples, a hydrogen atom must be bonded directly to one of the three most electronegative atoms — nitrogen, oxygen, or fluorine — and that hydrogen must then be attracted to a lone pair on the N, O, or F of another molecule. When those conditions are met, the attraction is remarkably strong for an IMF. This is the secret behind water's strange and life-giving behavior: an unusually high boiling point for such a small molecule, a solid form (ice) that floats because hydrogen bonds hold the molecules in an open, lower-density lattice, and a surface tension strong enough for insects to walk on.

Property comparisons should combine dispersion, permanent polarity, hydrogen-bond context, size, shape and phase. A simple hydrogen-bonding > dipole > dispersion slogan does not rank every real pair: accumulated dispersion can be substantial, and crystal packing strongly affects melting. Stronger relevant attraction often raises the energy needed for molecular separation in a matched comparison, but a single category label cannot set every boiling point, viscosity or solubility.

Comparing real substances shows the idea in action. At ordinary atmospheric pressure, water (H₂O) boils near 100°C while methane (CH₄), of nearly identical mass, boils at −161°C — because water hydrogen-bonds and methane has only weak dispersion forces. Even two molecules with the *same* formula can differ dramatically: ethanol (C₂H₅OH) has an O–H group and boils at 78°C, while dimethyl ether (CH₃OCH₃), the same atoms rearranged with no O–H, lacks an O–H donor and boils near −24°C at ordinary atmospheric pressure; its oxygen can still accept suitable hydrogen-bond interactions. Reading boiling points like this makes IMFs a detective's tool. And it closes the loop with the rest of this unit: intermolecular forces are exactly the attractions an 'ideal' gas is assumed to lack, which is why gases with weak IMFs behave most ideally and why real-gas phase behavior depends on temperature, pressure and interactions; above the critical temperature, compression alone does not cross an ordinary liquid–gas coexistence boundary.

Model check: Hydrogen bonding is not simply every bond containing H; ordinary boiling does not break covalent molecules apart. All particles have dispersion attractions; larger polarizable electron clouds can produce substantial attraction.

Worked examples

Avoid these mistakes

• Hydrogen bonding is not simply every bond containing H; ordinary boiling does not break covalent molecules apart.

Words to know

intermolecular forces (IMFs)

The attractions between separate molecules, as opposed to the much stronger bonds within a molecule.

London dispersion forces

Weak attractions present in ALL molecules, caused by fleeting, temporary dipoles from shifting electrons; stronger in larger molecules with more electrons.

dipole-dipole forces

Attractions between the permanently positive end of one polar molecule and the negative end of another.

hydrogen bonding

An especially strong dipole-dipole attraction that forms when hydrogen is bonded to N, O, or F and is drawn to a lone pair on the N, O, or F of a neighboring molecule.

polarity

An uneven sharing of electrons that leaves a molecule with a positive end and a negative end (a dipole).

polarizability

How easily a molecule's electron cloud can be distorted; larger, electron-rich clouds are more polarizable and give stronger dispersion forces.

The number-one chemistry mix-up: bonds vs. attractions

When water boils away in a pot, what breaks? Not the water molecules — the sturdy O–H covalent bonds inside each H₂O survive completely intact, which is why steam is still water and not loose hydrogen and oxygen. What actually pulls apart are the far weaker intermolecular forces between separate molecules. The gap in strength is huge: a hydrogen bond is only about a twentieth as strong as the covalent bond it sits next to. Confusing the two — thinking boiling or melting breaks chemical bonds — is the single most common error in this whole unit.

Hydrogen bonds hold your DNA together

The famous double helix of DNA is zipped shut by hydrogen bonds between its base pairs: adenine grips thymine with two hydrogen bonds, and guanine grips cytosine with three. That strength is perfectly tuned — individually the bonds are weak enough that an enzyme can 'unzip' the two strands to copy them, yet collectively, over millions of base pairs, they are strong enough to keep your genetic code stable for a lifetime. The same hydrogen bonds fold proteins into the precise shapes that let them do their jobs. Life quite literally runs on intermolecular forces.

Everything strange about water traces to one force

Water breaks the rules for a molecule its size, and hydrogen bonding is behind nearly every oddity. It is why ice floats — the hydrogen bonds lock frozen water into an open, lower-density lattice, so lakes freeze from the top down and fish survive the winter beneath. It gives water an enormous capacity to soak up heat, which moderates Earth's climate, and a surface tension strong enough for a water strider to walk across a pond. That same clinginess even lets water climb the inside of the tallest trees, hauling itself up to the highest leaves.

Who discovered these invisible forces?

The weak attractions between molecules are named for Johannes van der Waals, the Dutch physicist who, in correcting the ideal gas law, realized real molecules must gently pull on one another — work that won him the 1910 Nobel Prize in Physics. The strongest of them, the hydrogen bond, was identified and named later, by the American chemists Wendell Latimer and Worth Rodebush in 1920, to explain water's peculiar behavior. It is a satisfying loop: the very forces that make the ideal gas law imperfect are the ones that give liquids and solids their character.

Why water beads on a waxed car but soap makes it spread

Splash water on a freshly waxed hood and it pulls itself into tight beads — the hydrogen bonds between water molecules tug inward, minimizing the surface in a display of surface tension. Add a drop of dish soap and the beads instantly collapse and spread flat. Soap molecules are surfactants, with one end that loves water and one end that flees it, so they wedge into the surface and disrupt the water's grip on itself. That same trick is how soap lifts greasy dirt: it surrounds oil droplets that plain water, held together by its own strong forces, simply cannot touch.

In the real world

Geckos climb glass using nothing but London dispersion forces: millions of microscopic hairs on their toes make so much combined contact that these individually feeble forces add up to support the animal's whole weight.

Exit check

Why can nonpolar molecules still condense?