Water is truly the molecule of life as we know of no living systems that can exist (in the poorly defined state that we call "alive") without the presence of liquid water. You may be surprised to find out how much water is in various organisms. Lettuce is about 95% water, carrots—85% water, animal cells about 80% water. At the other end of the spectrum, there are dessicated seeds and spores that can have as little as 5% of their volume composed of water, although at these low water contents, the organisms exist in a state of suspended animation—with no biochemistry occuring.
What is Water?
Water is an oxygen atom with two hydrogen atoms attached to it by covalent bonds. It is a solid below 0° C, a liquid between 0° C and 100° C, and a gas above 100° C, at atmospheric pressure. Our Earth (which obviously should have been called Ocean) has 2/3 of its surface covered with water (more during ice ages), and our weather is largely determined by the water entering and leaving the atmosphere from the Earth’s surface. In living things, water serves as a structural element (due to its incompressibility), as a medium for diffusion (and hence communication between cells), as a solvent and as a component in biochemical reactions. Water is a singular liquid, unlike the hydrides of the elements close to oxygen in the periodic table (NH3, HF, H2S are all gases at room temperature, H2O should have a boiling point of –93° C) with unique properties that make it ideal for being the basis of living systems. These unique properties are consequences of the molecular structure of water, so we shall first look at structure and then go on to see how this leads to the strange behavior of water.
Oxygen has six valence electrons and each hydrogen atom has one, so the hydrogen atoms covalently bond to the oxygen leaving two lone pairs of electrons on the oxygen. The length of the bonds is 1Å and the angle between them is 105° , very near the angle between the vertices of a regular tetrahedron (109° and change). If the Schrödinger equation for a water molecule is solved to find the average density of the outer shell electrons, the following picture of the water molecule emerges:
[James Trefil. Meditations at 10,000 Feet. Macmillan Publishing co., New York, N.Y. 1986]
The two "horns" behind the oxygen molecule represent a higher electron density and so the average charge density there is slightly negative. This leaves the two protons with a lower electron density and a slightly positive charge density. If we were to plot these regions of altered charge density in three dimensions, we would find that they very nearly form the vertices of a regular tetrahedron.
The hydrogen bond is a chemical bond formed by the attraction between the electronegative region of an oxygen molecule and the electropositive region of a proton that is bonded to an oxygen. It is weaker than a covalent bond, having a binding energy of about 2-10 kcal/mol (whereas covalent bonds are on the order of 100 kcal/mol). In water, there are the same number of electronegative regions as electropositive, so each such region can participate in hydrogen bonding with a neighbor.
[James Trefil. Meditations at 10,000 Feet. Macmillan Publishing co., New York, N.Y. 1986]
With the tetrahedral geometry of the water molecule, an infinite array of such hydrogen bonding can be set up. In fact, this is the structure of ice Ih (the ice that forms under normal atmospheric pressures).
[ N.H. Fletcher. The Chemical Physics of Ice. Cambridge University Press, London. 1970]
This is the same geometrical structure as diamond. The rods in this picture represent hydrogen bonds between molecules, having an average strength of 4.5 kcal/mol. The structure of liquid water lies somewhere between this crystalline array and a random arrangement of the molecules such as exists in a normal liquid.
Models of the Structure of Liquid Water
Although we have the promises of leading physicists that superstring theory will bring us the end of physics, we still don’t have a physical model of liquid water that explains all of its behavior (phase changes are a real PITA). This behavior has been studied extensively and duly documented in learned treatises. For the purposes of understanding some of the properties of water that are important for living things, two simple models of the structure of liquid water will suffice.
The uniform model considers the liquid to be composed of a uniform subunit that represents the average interaction between molecules. This is similar to our standard model of fluids in which they are treated like a bag of marbles. This model predicts that as heat is added to water, the kinetic energy of each molecule will increase thereby increasing the average distance separating each molecule at the expense of the binding energy between molecules. The density of such a liquid is expected to decrease with temperature right from the solid-liquid phase change to the liquid-gas phase change. If we look at the density of water over this range we find the following:
If you look really closely at this picture (you won’t see anything wrong, but pretend that you do just to prevent any trouble) you will see that the density actually peaks above the melting point, at 4ºC, and then drops slightly as the temperature is lowered further. Even if you don’t see this, you know that it is correct since you can probably recall dropping ice into a cocktail last weekend and seeing it come bobbing up to the surface to float on top (this also works with pure water so don’t give me your quibbles about the altered density due to mixing alcohol into the solution). The molecules of liquid water are actually more densely packed than in the solid. This is fortuitous as well, since lakes, rivers, and oceans freeze on the surface first, thereby providing a barrier to evaporative, convective, and radiative heat loss that allows them to remain unfrozen underneath despite long exposures to sub-freezing temperatures at the surface. Fortuitous because life requires liquid water and the Earth would have frozen over billions of years ago were it not for this property of water.
The mixture model of water structure (due to Roentgen) attempts to account for the parabolic dependence of density with temperature. The mixture model posits that some of the ice-like structure remains after the phase transition and that the clusters of water molecules are able to pack themselves more densely without the hydrogen bonding network between them. The latent heat of fusion (the heat required to go from ice at 0ºC to water at 0ºC) is 6 kJ/mol (note the change in units (so sue me)) whereas the energy of breaking all the hydrogen bonds would be 40 kJ/mol, therefore the heat of fusion implies that less than 15% of the hydrogen bonds are actually broken upon melting of ice. So we have two competing processes affecting the density as the temperature increases. The added heat will go into breaking hydrogen bonds (thereby making the average size of the clusters smaller) and increasing the kinetic energy of the clusters. As the cluster size decreases, the water becomes more dense but as the kinetic energy of each cluster increases, the water becomes less dense. Since large clusters will need a lot of kinetic energy to significantly alter their velocity, the predominating effect near 0ºC is the collapse of clusters. At higher temperatures where the cluster size is small enough that the energy of a hydrogen bond provides a significant increase in the kinetic energy of a cluster, expansion predominates. The maximum density occurs at 4ºC.
The heat required to separate molecules from a liquid and move them into the adjacent gas phase without a change in temperature is called the heat of vaporization. Not only must the hydrogen bonds be broken, but the molecule must have enough kinetic energy, upon reaching the surface, to break free of the Van der Waal’s forces that exist between the molecule and other molecules in the liquid. In water, the heat of vaporization at 100ºC is 40.7 kJ/mol, the highest of any known liquid. This is mainly due to the large number of hydrogen bonds that remain when the liquid reaches the boiling point.
Properties of Water
The specific heat capacity of a substance is the energy required to change the temperature of 1 gram of material by 1ºC. It is a measure of how much heat a substance can store, the thermal inertia, if you will. The calorie is defined as the amount of heat required to raise the temperature of 1g of water by 1ºC at 15ºC, so the specific heat of water is 1 (in appropriate units). Liquid ammonia is the only liquid with a higher specific heat. The high specific heat capacity of water is due to the gradual breaking of hydrogen bonds with increasing temperature. Biologically, the high specific heat of water is advantageous as it allows the temperature of an organism to be buffered against rapid changes with alterations in the environment. Water is the heat reservoir of choice for solar energy enthusiasts due to this property, and if you regard plants as wonderfully efficient solar energy factories, their success can be partly attributed to being made from the right stuff.
Viscosity is a measure of resistance to shear within a liquid; the lower the viscosity, the easier it flows. The viscosity of water at room temperature is 0.009 poise, very low considering the large clusters that exist in liquid water. At room temperature, 70% of water molecules are in clusters containing an average of 57 molecules each. These clusters are so short-lived, however, that the viscosity is not affected. Many solutes increase the viscosity of aqueous solutions (consider this over Thanksgiving when you’re desperately trying to thicken the gravy with the turkey already cooling on the table) and it also increases as the temperature decreases. Viscosity greatly affects the growth rate and morphology of crystals during freezing, so this is an important property for cryobiology. As the viscosity increases, a point can arise when crystal growth is no longer allowed kinetically, even though ice nuclei may be forming. Continued cooling past the glass-transition temperature will result in a glassy, amorphous solid. As this solid is warmed, the ice nuclei will grow once the viscosity is low enough so crystallization actually occurs on warming rather than cooling (the process of devitrification); a phenomenon we’ll look at more closely later.
Immiscibility with non-polar liquids
Liquids that have no polarization (no partial separation of charges like there are on the water molecule) have a very low solubility in water. An oil-vinegar salad dressing is a good example of what happens when you try to mix the two. The non-polar molecules would have to break the hydrogen bonds between water molecules in order to fit in between them, thus mixing is energetically unfavourable. This effect is what allows the "hydrophobic" bonds to form in aqueous solutions—the primary stabilizing factor for all biological membranes.
Also due to the hydrogen bonding network, water supports an extremely high surface tension. Insects can freely walk on water (or pick up a ball of water between their legs and drink it). A steel needle can be placed carefully on the surface and it will float.
When liquid water comes in contact with a surface with which it can form hydrogen bonds, it adheres to this surface. If the surface encloses a volume, then the column of water inside that volume will slightly climb the outer edge of the container to form a concave meniscus. If the radius of the container becomes small with respect to the radius of curvature of the meniscus, then the column of water can actually rise against gravity due to adhesion to the vessel walls. This phenomenon is called capillary rise and is used by plants to bring water from the ground up to the leaves.
The high degree of hydrogen bonding between water molecules gives the liquid a high tensile strength (cohesion). A column of water can support an negative pressure (tension) as long as cavitation (production of a surface, a bubble) doesn’t occur. The tensile strength of a water column is about 10% than of aluminium or copper. In large trees, since the water columns can be much higher than 9.81 m, the water columns are held together by the cohesive properties of water.
Since water molecules are polarized, they have a dipole moment (placing them in an electric field would produce a torque about their center of mass). With an applied electric field, water molecules will tend to align themselves in the field with their positive aspects toward the negative side of the field and their negative aspects toward the positive side of the field. In a random orientation, the electric fields produced by the charge separation on the water molecules would cancel, but in this coordinated alignment, the electric fields of the water molecules tend to combine and oppose an external electric field. The degree to which a substance does this is called its dielectric constant. Water has a high dielectric constant which allows it to act as a solvent for ionic compounds such as sodium chloride and potassium chloride. In the case of NaCl, the electric field that provides the ionic bond is reduced by 78 times when placed in water, so that the bond strength is lower than kT allowing thermal motion to separate the ions into solution. All biological cells require ionic gradients to be maintained, thus water is an admirable solvent for this purpose.
Another aspect of the high dielectric constant of water that is familiar to most people is the difficulty in defrosting foods with a microwave oven. Microwaves produce an oscillating electric field which tends to jiggle those molecules that can align themselves with the field. Thus water is easily heated by microwaves, but ice, trapped in its lattice is unable to have each molecule oriented with the external field. This leads to runaway heating of the liquid components (and subsequent cooking of small regions) while the frozen components don’t even warm up.
The Polywater Episode
In the Correspondence section of the October 11, 1969 issue of Nature, the following letter appeared under the heading, "Anomalous" Water:
SIR,—A report on the properties of "anomalous" water appeared recently in Nature (222, 159; 1969). The probable structure of this phase was reported by Lippincott et al.1 who refer to the phase as polywater, a terse descriptive of the structure.
The properties of polywater are negligible vapour pressure, density ~1.4 g/cm3, partial miscibility with normal water (depending on the length of the polymer chains) and stability to temperatures ~500ºC. The polymer chains are some 250-420 kJ/mole (60-100 kcal/mole) of monomer more tightly bound than normal water.
I need not spell out in detail the consequences if the polymer phase can grow at the expense of normal water under any conditions found in the environment. Polywater may or may not be the secret of Venus’s missing water. The polymerization of Earth’s water would turn her into a reasonable facsimile of Venus.
There are examples of phases in other systems which are difficult to nucleate. Once the nuclei are present, the phases grow readily, often by mechanisms other than those required to form the nuclei. It is almost a truism that, under conditions where both a stable phase and a metastable phase may form, the metastable phase forms first. In this case the metastable phase would be normal water.
After being convinced of the existence of polywater, I am not easily persuaded that it is not dangerous. The consequences of being wrong about this matter are so serious that only positive evidence that there is no danger would be acceptable. Only the existence of natural (ambient) mechanisms which depolymerize the material would prove its safety. Until such mechanisms are known to exist, I regard the polymer as the most dangerous material on earth.
Every effort must be made to establish the absolute safety of the material before it is commercially produced. Once the polymer nuclei become dispersed in the soil it will be too late to do anything. Even as I write there are undoubtedly scores of groups preparing polywater.
Scientists everywhere must be alerted to the need for extreme caution in the disposal of polywater. Treat it as the most deadly virus until its safety is established.
Pennsylvania 18703, USA.
1Lippincott, E.R., Stromberg, R.R., Grant, W.H., and Cessac, G.L. Science, 164, 1482 (1969).
And then the shit hit the fan...