Humble H20 not all that it seems

Water is all around us, among the most abundant material we encounter (and pouring copiously from the sky as I write this).

It is essential for life as we know it, as the medium for biochemical reactions and comprising from about 50 to 80 per cent by weight of human bodies (depending on age, gender and physique).

It is also essential for agriculture, widely used in many industrial processes, used in heat transfer (cooling or heating), as a solvent, for mechanical work (from water wheels to steam turbines), transport and leisure.

The chemical composition of water is deceptively simple. nyone with a rudimentary knowledge of chemistry knows its molecular formula, H2O, indicating two atoms of hydrogen bonded to one of oxygen.

But if that were the complete story, water would be a gas under “normal” conditions, like methane or carbon dioxide and not much good for conducting biochemical reactions, let alone bursting pipes, mixing with whisky or floating supertankers.

Properties
Understanding its physical properties requires a deeper look at bonding around the water molecule.

Firstly, consider the bonds between oxygen and hydrogen.

These are described as “covalent”, since each is made up by sharing one electron from the hydrogen and one from the oxygen atom.

These paired electrons then occupy a region (an orbital) that surrounds the hydrogen and oxygen atoms. But there’s a twist. Some covalent bonds are more equal than others. In the case of water, the oxygen atoms exert a greater pull on the electrons, developing a slight negative charge compared with the hydrogen atoms, which become slightly positively charged.

Next, there are four more electrons around the oxygen atoms, which are also available for bonding. But these are already paired up and can form “dative” bonds by being donated to another atom that wants them. Consequently, this is often called hydrogen bonding.

Hydrogen bonds are weaker than covalent bonds and rather promiscuous, in that the participating atoms can change partners relatively easily. Nevertheless, this joins the molecules together in a network, which makes water a liquid under normal conditions and a solid at low temperatures.

This hydrogen bonding network causes the characteristic volume increase (roughly 9 per cent) when water freezes. Conversely, increased pressure distorts the hydrogen bonding, lowering the freezing point, which can cause a lubricating layer of water under ice-skates and glaciers.

Although it becomes less extensive as water is warmed up, due to increasing thermal energy in the form of molecular motion, the hydrogen bonding network only breaks down completely when liquid water turns to vapour.

The bonds around the oxygen atom move as far apart as possible, to minimise electrostatic repulsion; consequently, they point towards the corners of a tetrahedron, with the oxygen atom at the centre.

Viewed one way, the oxygen atoms in ice become joined together in sheets of puckered hexagons, causing the six-fold symmetry familiar in snowflakes.

Viewed another way, the crystal structure in ice is similar to that of diamond.

Although the bonds in ice are not as strong as those between carbon atoms, high pressures can still occur.

At –10°C, a pressure of 1000kg per square cm can be generated, equivalent to parking a large car on a 5p piece, which is more than enough to burst pipes!


DR LAITY is Senior Research Fellow in the Department of Chemical and Biological Sciences, University of Huddersfield.