hydrogen bonding
/ HY-dro-jen /
Watch a water strider walk on a pond, or notice how water beads up on a freshly waxed car: water behaves as if it had a skin, far stickier than its tiny molecules should allow. The secret is hydrogen bonding — a special, unusually strong attraction that links each water molecule to its neighbors, like countless tiny magnets snapping together.
A hydrogen bond forms when a hydrogen atom already bonded to a small, hungry-for-electrons atom (nitrogen, oxygen, or fluorine) is drawn toward a lone pair of electrons on another such atom nearby. Because that hydrogen is left almost bare of its electron, it carries a strong positive patch that locks onto the negative lone pair. It is the strongest of the ordinary intermolecular forces — several times stronger than plain dipole–dipole attraction, though still far weaker than a real chemical bond.
Why it matters: hydrogen bonding gives water its astonishingly high boiling point, lets ice float, holds the two strands of DNA together, and folds proteins into shape. The honest caveat is that it sits in a gray zone between an intermolecular force and a true bond — strong and directional like a bond, yet weak and reversible like an attraction — which is exactly why it is so useful in living things.
Hydrogen fluoride (HF) boils at 20 °C, while the heavier hydrogen chloride (HCl) boils at minus 85 °C — HF's far higher boiling point comes from the strong hydrogen bonds its molecules form.
Hydrogen bonding pushes boiling points abnormally high.
A hydrogen bond is not a chemical bond despite the name. The hydrogen stays covalently attached to its own molecule; the 'bond' is the much weaker attraction to a neighboring electron lone pair. Only H attached to N, O, or F gives strong hydrogen bonds.