Water is the most familiar substance on Earth, and also one of the strangest. Based only on its small molar mass, water 'should' be a gas at room temperature, like other small hydrides such as H₂S. Instead it's a liquid, it expands when it freezes, and it dissolves an enormous range of substances. Every one of these unusual behaviours comes from the same root cause: water is bent, polar, and forms extensive hydrogen bonds. Look at what each molecule can do. The oxygen carries two lone pairs and a strong charge; each hydrogen is a nearly bare nucleus. So one water molecule can accept two hydrogen bonds through its lone pairs and donate two through its hydrogens, linking to up to four neighbours at once. Few molecules can build such a dense three-dimensional web, and this lesson traces each of water's strange properties back to that web.
A boiling point far too high for its size
Compare water with the other Group 16 hydrides, which have the same bent shape. From up to the boiling point falls as the molecules get smaller, exactly as dispersion forces predict. Following that trend, water would boil somewhere near −80 °C and be a gas on Earth. Instead it boils at 100 °C. Sulfur, selenium and tellurium are not electronegative enough for hydrogen bonding, but oxygen is, and those extra attractions add roughly 180 degrees to water's boiling point. Without them there would be no liquid oceans and no life as we know it.
| Hydride | Molar mass (g/mol) | Boiling point | Strongest IMF |
|---|---|---|---|
| 18.02 | 100 °C | Hydrogen bonding | |
| 34.08 | −60 °C | Dipole-dipole | |
| 80.99 | −41 °C | Dipole-dipole | |
| 129.62 | −2 °C | Dipole-dipole |
Ice is less dense than liquid water
In liquid water, hydrogen bonds constantly break and reform, letting molecules pack together relatively closely. As water freezes, its molecules lock into a rigid, open hexagonal lattice, with each molecule hydrogen-bonded to exactly 4 neighbours in fixed positions. This ordered structure actually takes up more space than the jumbled liquid — so ice is less dense than liquid water and floats. This is highly unusual: almost every other substance is denser as a solid than as a liquid. Ice has a density of about 0.917 g/mL, while liquid water is about 1.00 g/mL (it is densest at 4 °C). So water expands by roughly 9% when it freezes, which is enough to burst pipes, crack rocks as water freezes in their cracks, and heave roads into potholes during Canadian winters. Because freezing keeps the mass the same, you can find the new volume from , rearranged to .
High specific heat capacity
Water absorbs a large amount of thermal energy for a relatively small temperature change, because much of that energy goes into breaking and re-forming hydrogen bonds rather than directly speeding up molecular motion. This is why large bodies of water moderate the climate of nearby land, and why water is used as a coolant. It takes 4.18 J to warm 1 g of water by 1 °C, compared with about 2.4 J for ethanol and only 0.90 J for aluminum. That is why a beach's sand gets scorching on a sunny afternoon while the lake stays cool, and why towns on the Great Lakes have milder winters and cooler summers than places farther inland. Water's high heat of vaporization works the same way: evaporating sweat has to break many hydrogen bonds, so it carries a lot of heat away from your skin.
| Substance | Specific heat capacity (J/(g·°C)) | Main IMF or bonding |
|---|---|---|
| Liquid water | 4.18 | Hydrogen bonding |
| Ethanol | about 2.4 | Hydrogen bonding (one O–H per molecule) |
| Aluminum | 0.90 | Metallic bonding |
| Copper | 0.39 | Metallic bonding |
High surface tension, cohesion and adhesion
Molecules at water's surface are pulled inward and sideways by hydrogen bonds to their neighbours, but have no water molecules above them pulling outward. This net inward pull creates a taut 'skin' at the surface — strong enough to support small insects and cause water to bead up. The attraction of water molecules to each other is called cohesion; their attraction to a different surface, such as glass or the walls of a plant's xylem, is called adhesion. When adhesion to a narrow tube is strong, water climbs up the walls and drags the cohesive column behind it. This capillary action helps move water from roots to leaves and explains why the water surface in a glass graduated cylinder curves into a U-shaped meniscus that you read at the bottom.
The universal solvent
Water's polarity lets it dissolve a huge range of substances. Its partially negative oxygen end surrounds cations, and its partially positive hydrogen ends surround anions (ion-dipole attraction), pulling ionic compounds apart into solution. Its ability to hydrogen bond also lets it dissolve other polar molecules, like sugars and alcohols. Nonpolar substances (like oils), which cannot hydrogen bond or form strong dipole interactions with water, remain insoluble — 'like dissolves like'.
- Ionic compounds such as NaCl (ion-dipole attraction)
- Small polar molecules with O–H or N–H such as sugar, ethanol, NH₃ (hydrogen bonding)
- Acids such as HCl, which form ions in solution
- Nonpolar molecules such as oils, fats and waxes
- Hydrocarbons such as gasoline and hexane
- Nonpolar gases such as N₂ and O₂ (only slightly)