Intermolecular Forces
The weak attractions BETWEEN molecules — not the bonds inside them — are what decide whether a substance is a gas, a liquid, or a solid.
Water boils at 100 °C, but hydrogen sulfide — a heavier, very similar-looking molecule — boils at about −60 °C. If mass were the whole story it should be the other way around. The answer is a set of weak but decisive attractions acting between molecules: intermolecular forces.
Inside vs between: two very different attractions
Every molecule is held together inside by strong covalent bonds. But molecules also feel weaker attractions to each other — these are intermolecular forces (IMFs). The distinction is the whole lesson: covalent bonds are intramolecular (within), IMFs are intermolecular (between).
IMFs are much weaker than covalent bonds — often 10 to 100 times weaker — but they punch above their weight, because they decide whether a substance is a gas, liquid, or solid at room temperature.
1. London dispersion forces — in everything
Even a nonpolar molecule has electrons in constant motion. At any instant they can bunch to one side, creating a fleeting temporary dipole that induces a matching dipole in a neighbour. That flickering attraction is the London dispersion force, and it acts between all molecules.
It's weak per encounter but grows with the number of electrons — so bigger, heavier molecules (and larger atoms) have stronger dispersion forces. It's why iodine (I₂) is a solid while fluorine (F₂) is a gas, and why boiling points climb down a column of noble gases.
2. Dipole-dipole forces — for polar molecules
A polar molecule has a permanent positive end and a negative end (recall the net dipole from the polarity lesson). Neighbouring polar molecules line up so the δ+ of one sits near the δ− of another — a dipole-dipole attraction. It's stronger than dispersion for molecules of similar size, which is why polar substances generally boil higher than nonpolar ones of comparable mass.
3. Hydrogen bonding — the strong one
Hydrogen bonding is an especially strong dipole-dipole attraction that appears in a specific case: when hydrogen is bonded to a small, highly electronegative atom — N, O, or F — the H becomes very positive and is drawn to a lone pair on the N, O, or F of a neighbouring molecule.
It's the strongest of the three IMFs (though still far weaker than a covalent bond), and it is why water has such a strangely high boiling point, why ice floats, and why DNA's two strands hold together yet unzip easily.
- Both molecules are polar, so both have dipole-dipole forces and London dispersion.
- Look for an –O–H group: ethanol has one (its H is bonded directly to oxygen); dimethyl ether does not (its H atoms are all on carbon).
- So ethanol molecules can hydrogen-bond to each other; dimethyl ether cannot.
- Hydrogen bonding is the stronger IMF, so more energy is needed to separate ethanol molecules — a much higher boiling point.
- Each of water's 2 hydrogens can hydrogen-bond to a neighbour's lone pair: that's 2.
- Each of oxygen's 2 lone pairs can accept a hydrogen from a neighbour: that's 2 more.
- 2 + 2 = 4 hydrogen bonds — the tetrahedral network that makes ice's open, lower-density structure (so ice floats).
Check your understanding
- Intermolecular forces (IMFs) act BETWEEN molecules; covalent bonds act WITHIN them and are much stronger.
- London dispersion acts in all molecules and grows with size/number of electrons.
- Dipole-dipole acts between polar molecules; hydrogen bonding (H on N, O, or F) is the strongest IMF.
- Stronger IMFs → more energy to separate molecules → higher melting and boiling points.
- Boiling and melting break IMFs, not the covalent bonds inside molecules.