When a molecule has a built-in lean
Some molecules are lopsided not just for a flicker, but permanently. The reason is that atoms differ in how hard they pull on shared electrons — a property called electronegativity. When two different atoms bond, the greedier one hogs the shared electrons, taking on a slight negative charge while the other is left slightly positive. The bond now has a built-in lean: a bond polarity.
If a whole molecule's polar bonds don't cancel out, the molecule itself ends up with a positive end and a negative end. We measure that overall lean with the dipole moment: a bigger dipole moment means a more lopsided molecule. Water has a large one; carbon dioxide, though its bonds are polar, is straight and symmetric so the leans cancel — its dipole moment is zero.
Dipoles attract dipoles
Once molecules have permanent positive and negative ends, they line up the obvious way: the positive end of one nestles against the negative end of its neighbour, like little bar magnets snapping into rows. This is the dipole–dipole interaction, the second member of the van der Waals family. Because the dipoles are permanent — not fleeting like dispersion — the attraction is steadier and usually stronger for molecules of similar size.
So a polar molecule feels *both* dispersion (everyone gets that) *and* dipole–dipole on top. That is why, comparing two molecules of similar size, the more polar one usually boils hotter — it has more glue holding it together. This neatly extends the boiling-point story from the last guide: dispersion sets the baseline, and permanent dipoles add to it.
The standout strong one: hydrogen bonding
There is a special, unusually strong version of dipole–dipole attraction that earns its own name: the hydrogen bond. It happens in a precise situation — when a hydrogen atom is bonded to one of three very electron-greedy atoms (nitrogen, oxygen, or fluorine), and a nearby molecule has one of those greedy atoms with a lone pair of electrons to offer.
Why so strong? Hydrogen is the smallest atom, with just one electron. When that electron is pulled away by a greedy neighbour, the bare hydrogen nucleus — a naked proton — is left exposed, with almost nothing to shield it. That intense little patch of positive charge can grip a neighbouring lone pair tightly. A hydrogen bond is several times stronger than an ordinary dipole–dipole tug, though still far weaker than a covalent bond.
Why water is so strange
Water is the champion of hydrogen bonding. Each tiny molecule can form up to four hydrogen bonds — two through its hydrogens, two through the oxygen's lone pairs — knitting the whole liquid into a shifting, three-dimensional net. This single fact explains a long list of water's oddities, starting with its boiling point. Molecules of similar size, like methane, are gases far below freezing; water, knitted by hydrogen bonds, holds together all the way to 100 °C.
Hydrogen bonds also explain why ice floats — a genuinely rare and life-saving quirk. As water freezes, its molecules lock into an open, roomy cage that holds them slightly *farther* apart than in the jostling liquid. So ice is less dense than water, floats on top, and lakes freeze from the surface down instead of solid through — letting fish survive the winter beneath.
The same hydrogen bonds hold the two strands of DNA together and fold proteins into their working shapes. They are weak enough to be unzipped and re-formed, yet strong enough to hold the pattern — exactly the property life needs. Before we leave dipoles, it is worth naming one even stronger relative: when a full ion sits beside a dipole, you get an ion–dipole interaction, the powerful attraction that lets water pull salt apart, which we explore in the final guide.