Why Water Has Such Strange Properties

Drop an ice cube into a glass of water and watch it float. This is so familiar that it takes effort to notice how strange it is. Almost every other substance on Earth contracts when it freezes, so the solid form sinks in its own liquid. Solid wax sinks in melted wax. Solid iron sinks in molten iron. Water does the opposite, and the reason traces back to the shape of a single molecule.

A water molecule is two hydrogen atoms bonded to one oxygen atom, but the three atoms do not lie in a straight line. Oxygen holds onto electrons more tightly than hydrogen does, a tendency called electronegativity. The shared electrons in each O–H bond spend more time near the oxygen, leaving the hydrogens slightly positive and the oxygen slightly negative. Oxygen also carries two pairs of electrons that are not involved in bonding, called lone pairs, and these push the hydrogens down into a bent shape with an angle of about 104.5 degrees. The result is a molecule with a definite positive end and a definite negative end. Chemists call such a molecule polar.

Polarity changes everything. Because the positive hydrogen of one water molecule is attracted to the negative oxygen of another, water molecules cling to each other through a particularly strong attraction called a hydrogen bond. A hydrogen bond is not as strong as the covalent bond inside a molecule, but it is far stronger than the weak attractions between most small molecules. Each water molecule can form up to four hydrogen bonds at once: two through its hydrogens, two through its lone pairs. Liquid water is therefore not a loose crowd of independent molecules but a shifting network, with bonds breaking and reforming trillions of times per second.

This network is what makes water peculiar. Consider the floating ice. When water cools toward freezing, the hydrogen bonds lock the molecules into a rigid lattice in which each molecule sits at the center of a small tetrahedron of neighbors. Holding that geometry requires the molecules to space themselves slightly farther apart than they were in the liquid. Ice is therefore less dense than the water it came from, and it floats. Lakes freeze from the top down rather than the bottom up, which lets fish survive winter under an insulating lid.

The same network explains water's stubborn resistance to temperature change. To warm water, you must supply enough energy not only to make the molecules move faster but also to break some of the hydrogen bonds holding them together. This is why coastal climates are milder than inland ones: a body of water absorbs and releases heat slowly, buffering the air around it. The same logic explains why sweating cools you. To evaporate, a water molecule has to break free of nearly all its hydrogen-bonded neighbors, which costs a great deal of energy, and that energy comes out of your skin.

Polarity also makes water an unusually good solvent. Salt dissolves because the positive ends of water molecules surround the chloride ions while the negative ends surround the sodium ions, prying the crystal apart. Sugars and many proteins dissolve because they too have polar regions that water can grip. Oils, which are nonpolar, do not dissolve, because water molecules would rather stay bonded to each other than make room for a molecule that offers nothing in return. This selectivity is the foundation of nearly every biological process: cells use water's preferences to fold proteins, organize membranes, and sort what enters and leaves.

None of these properties are exotic additions to water. They all fall out of one geometric fact and one electrical fact: the molecule is bent, and its ends carry opposite charges. The strangeness of water is not strangeness at all once you see the shape behind it.