Why does water change phase so easily?

Question

I did some undergraduate level Chemistry as part of my degree. It left me with the distinct impression that the polarity of water is responsible for its unusual phase properties: it's the only everyday substance that can exist in all three states under easily replicable conditions.

My daughter is learning about states of matter at school, and I thought this might be a fun "extra" for her homework, so I sat down to explain it to her. At which point I realised that I couldn't explain it myself. Polarity explains why water solidifies relatively easily, and why ice floats/expands - but I couldn't see/remember why it might have anything to do with transition into a gas.

So I taught her about phase diagrams instead, as a nice visual thing to take to class.

But it left me with the question: does water's polarity help explain its unusual phase properties? Or is it something else? Or are those phase properties perhaps not as unusual as I seem to remember?

Answer 1

does water's polarity help explain its unusual phase properties?

Not just its polarity, but its ability to form hydrogen bonds, together with the size and shape of its molecules.

Hydrogen bonds give water an unusually high boiling point for a molecule its size; many molecules of similar size that can't hydrogen-bond are gases at room temperature. They also give water a relatively high heat capacity and high enthalpies of fusion and vaporization. These latter two properties mean that ice has to absorb more heat to melt and liquid water has to absorb more heat to boil than would otherwise be necessary without hydrogen bonding.

Water's shape lets it form 4 tetrahedrally arranged hydrogen bonds. This lets it hydrogen-bond into networks of chair-shaped hexagonal rings in ice. The voids in the center of the rings make ice less dense than water. This is unusual; the solid form of a compound is most often denser than the liquid form.

Answer 2

Vapor pressure of water is quite small, thanks to H-bonding in the liquid. Therefore more energy is needed to evaporate (break H-bonds with neighboring molecules) a water molecule.

If you compare water to "nearest" non-H-bonded analog, dimethylether, $\ce{CH_3-O-CH_3}$ you'll find, that vapor pressure of ether is much higher (500 kPa) than water (2.3 kPa).