Intermolecular Forces

Definition: Intermolecular forces (IMFs) are the attractive forces between separate molecules, as opposed to the intramolecular bonds holding atoms together within a molecule, that determine physical properties like boiling point, melting point, viscosity, and solubility.

How It Works

IMFs are all electrostatic at heart, but they arise from different sources of charge imbalance, roughly from weakest to strongest:

London dispersion forces exist between all molecules, polar or not.
Electron distribution around any atom or molecule fluctuates instant to instant, briefly creating a tiny, temporary dipole.
That momentary dipole induces a matching temporary dipole in a neighboring molecule, and the two weakly attract.
Dispersion forces grow stronger with more electrons and greater surface area.
Larger, more polarizable molecules have more “give” in their electron cloud.
Dipole-dipole forces occur between molecules with a permanent dipole, e.g. HCl.
A permanent dipole is an asymmetric distribution of electron density from differing electronegativities.
The partially positive end of one molecule aligns with the partially negative end of a neighbor.
Hydrogen bonding is a particularly strong subtype of dipole-dipole interaction.
It occurs specifically when H is bonded directly to N, O, or F, small, highly electronegative atoms.
These atoms pull electron density away from hydrogen so strongly that its partial positive charge attracts a lone pair on a neighboring N, O, or F unusually strongly.
Ion-dipole forces occur when an ion interacts with a polar molecule, such as Na+ surrounded by water molecules when salt dissolves.
These are typically stronger than any of the three purely molecular IMFs above.
Because these forces hold separate molecules together, overcoming them, melting or boiling a substance, doesn’t break any covalent bonds.
It just separates whole molecules from each other.
That’s why molecular substances generally have far lower melting and boiling points than ionic or covalent network solids, where breaking apart the substance means breaking primary bonds.

Under the Hood

Rough energy scales:

London dispersion: 1-40 kJ/mol (scales with molecular size)
Dipole-dipole:      5-25 kJ/mol
Hydrogen bonding:    5-30 kJ/mol (HF is unusually strong)
Covalent bond:     150-1,100 kJ/mol (for comparison)

All IMFs are an order of magnitude or more weaker than a typical covalent bond.
This is why phase changes require far less energy than breaking a molecule apart chemically.

Worked example: dispersion forces and boiling point.

Given: noble gases have no permanent dipole and no hydrogen bonding, only London dispersion

Element	Electrons	Boiling Point (°C)
He	2	-269
Ne	10	-246
Ar	18	-186
Kr	36	-153
Xe	54	-108

Answer: boiling point rises smoothly with electron count/atomic size.
This is direct evidence that more electrons give a larger, more polarizable electron cloud and stronger momentary dipoles.

Worked example: hydrogen bonding anomaly.

Given: hydrides of group 16, H2O (18 g/mol) and H2S (34 g/mol)
Step 1: H2O boils at 100°C, H2S boils at -60°C
Step 2: molar mass trends predict H2S should boil higher, since it’s heavier with more dispersion forces
Step 3: water’s O-H bonds hydrogen bond extensively, each water molecule can form up to 4 hydrogen bonds
Step 4: H2S’s much less electronegative sulfur barely hydrogen bonds at all
Answer: the anomaly is direct evidence that H-bonding, not just molecular weight, dominates water’s behavior

Worked example: predicting relative boiling points.

Given: butane (C4H10, nonpolar, MW 58), acetone (C3H6O, polar, MW 58), 1-propanol (C3H8O, hydrogen bonds, MW 60)
Step 1: all three have nearly identical molar mass, so dispersion forces are roughly comparable
Step 2: the deciding factor is the strongest IMF type each can form
Answer: butane (dispersion only, -0.5°C) < acetone (dipole-dipole, 56°C) < 1-propanol (hydrogen bonding, 97°C)
This matches measured values closely, showing IMF type, not molecular weight, is the dominant factor when masses are similar.
Molecular shape also matters independently of polarity.
More elongated molecules can align more surface area against a neighbor, increasing total dispersion contact.
This holds even between molecules of identical formula.
n-pentane boils at 36°C while its branched isomer neopentane boils at 9.5°C, despite the identical molecular formula C5H12.

Why It Matters

“Like dissolves like” is a direct consequence of IMF matching.
Polar solvents dissolve polar/ionic solutes because the new solute-solvent interactions are comparably strong to what’s being broken.
Mixing polar and nonpolar substances, oil and water, fails because water’s strong hydrogen-bond network would have to be disrupted for no compensating energetic benefit.
IMFs also govern protein folding, hydrophobic side chains cluster away from water.
Protein folding is driven partly by London forces and partly by water’s own hydrogen-bond network.
IMFs govern DNA’s double helix, base pairs held by hydrogen bonds precise enough to encode genetic specificity.
IMFs govern surface tension, why water beads up and why some insects can walk on it.

Common Pitfalls

Confusing intermolecular forces, between molecules, govern phase changes, with intramolecular bonds, within a molecule, govern chemical identity.
Boiling water breaks hydrogen bonds between water molecules, not the O-H covalent bonds within them.
Assuming nonpolar molecules have “no” intermolecular forces; they still have London dispersion forces.
This is why nonpolar substances like nitrogen or octane can still be liquefied or solidified given a low enough temperature.
Forgetting that all molecules experience London dispersion simultaneously with any other IMFs present.
A polar molecule like HCl has both dipole-dipole and dispersion forces acting together, not just one or the other.
Requiring only “any H near any O” for hydrogen bonding.
The H must be covalently bonded directly to N, O, or F on one molecule, and interact with a lone pair on N, O, or F on another.
Assuming molecular size alone predicts boiling point without checking for hydrogen bonding.
As the H2O/H2S comparison shows, a lighter molecule can boil far higher if it hydrogen bonds and its heavier counterpart doesn’t.
Overlooking ion-dipole forces when discussing solutions; they’re often the strongest IMF present when an ionic solute dissolves in a polar solvent.

Comparison

Force	Requires	Relative Strength	Present In
London dispersion	Any molecule (temporary dipole)	Weakest (scales with size)	All molecules
Dipole-dipole	Permanent molecular dipole	Moderate	Polar molecules (e.g. HCl, acetone)
Hydrogen bonding	H bonded to N, O, or F	Strong	Water, ammonia, alcohols, DNA, proteins
Ion-dipole	Ion + polar molecule	Strongest of the four	Ionic solutes dissolved in water

Compound	Formula	MW (g/mol)	Dominant IMF	Boiling Point
Butane	C4H10	58	Dispersion only	-0.5°C
Acetone	C3H6O	58	Dipole-dipole	56°C
1-Propanol	C3H8O	60	Hydrogen bonding	97°C
n-Pentane	C5H12	72	Dispersion only, straight chain	36°C
Neopentane	C5H12	72	Dispersion only, branched	9.5°C

Real-World Application

Soap and detergent molecules exploit IMF mismatch to clean grease with water.

Each soap molecule has a long nonpolar hydrocarbon “tail” that dissolves into grease and oil via dispersion forces.
Each soap molecule also has a charged or polar “head” that dissolves into water via ion-dipole or hydrogen bonding.
In water, soap molecules spontaneously arrange into micelles, spherical clusters.
Tails point inward, away from water; heads face outward into it, trapping grease inside the nonpolar core.
The whole micelle, now effectively polar on its exterior, disperses into the surrounding water and rinses away.
This lets a fundamentally polar solvent remove a fundamentally nonpolar substance it couldn’t dissolve directly.

Example

Gecko feet are covered in millions of microscopic hair-like structures (setae) that maximize contact surface area with a wall or ceiling; the cumulative London dispersion force from that huge total surface area, not any chemical bond or suction, is strong enough to support the animal’s body weight, letting it walk up glass.

FAQ

Do ionic bonds count as an intermolecular force?

No, ionic bonding is an intramolecular, within-compound, interaction holding a lattice together.
Ion-dipole forces are the intermolecular category, describing how an already-formed ion interacts with separate polar molecules like water.

Why does water have unusually high surface tension?

Its extensive hydrogen-bond network pulls surface molecules strongly toward the bulk liquid, minimizing surface area.
This is why water beads up on a waxy, nonpolar, surface instead of spreading out.

Can a molecule’s shape affect its intermolecular forces even if its polarity stays the same?

Yes, more elongated or flatter molecules can pack closer together.
They align more of their surface area against a neighbor, increasing dispersion force contact even between molecules of identical formula.
This is part of why straight-chain alkanes generally boil higher than their branched isomers.

Referenced by

Dig deeper

Spot an error in this note?