HOT WATER

In the previous chapter, I pointed out that water was the only possible thalassogen for a planet at the temperature of the Earth, the only compound that could possibly exist in sufficient quantity in the liquid phase to form an ocean.

To be liquid at the relatively low temperatures of Earth (as I explained), a substance would have to consist of covalent molecules, that is, molecules in which pairs of neighboring atoms more or less share electrons in neighborly fashion, rather than carrying through a transfer of one or more electrons from one atom to another, bodily.

In general, the larger the molecular weight of a covalent compound the higher the temperature range of its liquid phase. From that standpoint one might expect a substance which is liquid at water temperatures to have a molecular weight of perhaps 180. The molecular weight of water, however, is 18, just one tenth of what it "ought" to be. On the basis of molecular weight, liquid water is surprisingly warm; it is "hot water" indeed.

But why is that? Are we perhaps oversimplifying if we relate liquid‑phase temperatures to molecular weight alone?

Well, at the end of the previous chapter I listed molecular weights and boiling points without any attempt to pick and choose among them. That is probably unfair, for substances are made up of different elements, and these differ greatly among themselves in chemical and physical properties. Yet elements exist in families, and within these, the members are quite similar. It might be best to stick to members of a particular family and see what regularities we can find there.

For instance, consider the six elements of the noble gas family, their atomic weights, and their boiling points.

Here we have a smooth rise in boiling point with the atomic weight, which is what we would expect, looking at the matter in an unsophisticated way. After all, if the atoms grow heavier, it takes more energy in the form of heat to lift them away from each other and send them off separately in vapor form.

TABLE I
Element	Atomic Weight	Boiling Point (oK.)	
			
Helium (He)	4.0	4.2	
Neon (Ne)	20.2	27.2	
Argon (Ar)	39.9	87.4	
Krypton (Kr)	83.8	120.2	
Xenon (Xe)	131.3	166.0	
Radon (Rn)	222.0	211.3	

"oK." represents the absolute scale of temperature with the zero Point ("absolute zero") at ‑273.16 oC.

What if we shift to the four elements of another family, the halogens, a family as well defined as that of the noble gases (see Table 2). Here, too, the boiling point rises smoothly with the atomic weight. There is a fifth halogen, the last in the series, which is named astatine. It is a radioactive element and even its most long‑lived nuclear variety (with an atomic weight of 210) has a half‑life of but 8.3 hours. It has not yet been obtained in quantities large enough to allow a clear boiling point determination, but I am willing to bet, sight unseen, and any reasonable sum, that its boiling point is somewhere in the neighborhood of 570oK.

While the progression is smooth within an element family, observe what happens when we cross the line. Compare Tables 1 and 2. Neon and fluorine are not very different in atomic weight but fluorine's boiling point is three times as high as neon's. This goes all the way down, each halogen having a boiling point approaching three times that of the noble gas with similar atomic weight.

TABLE 2
Element	Atomic Weight	Boiling Point (oK)	
Fluorine (F)	19.0	85.0	
Chlorine (Cl)	35.5	238.5	
Bromine (Br)	79.9	331.9	
Iodine (I)	126.9	457.5	

If we could stop there, we'd be able to work up a hardand‑fast relationship between particle size (whether atomic or molecular) and boiling point. In science, though, it isn't fair to stop at any point where you find yourself with the answer you want. You have to be sporting enough to look further and try to spoil your own hypothesis.

That's not hard to do. Chlorine atoms combine by twos also, and chlorine is C12, with a molecular weight of 71. That is distinctly less than the atomic weight of krypton, and yet the boiling point of chlorine is just twice that of krypton.

So we had better not try to cross the family lines in working up our theories. For the rest of the article I will stick to families and it will only be anomalies within the families that will receive our attention.

But let's see, is it only the boiling points that vary smoothly with atomic (or molecular) weight? Is the variation always direct, so that the measure grows larger as the weight goes up? Let's consider a third well‑marked element family, that of the "alkali metals" and take their melting points this time.

Cesium's melting point is down to 301 oK., or 28.5 oC., which means that it will melt on a hot summer day. There is also a sixth alkali metal, francium, which is radioactive, with its most long‑lived nuclear variety (atomic weight, 223) having a half‑life of only 21 minutes. Its melting point has not been determined but you can bet, though, it is very likely about 290oK. and that it will melt on a balmy spring day.

TABLE 3
Element	Atomic Weight	Melting Pt (oK)	
Lithium (Li)	6.9	452	
Sodium (Na)	23.0	371	
Potassium (K)	39.1	337	
Rubidium (Rb)	85.5	312	
Cesium (Cs)	132.9	301	

Other properties of other kinds vary in this regular fashion with atomic weights within element families, with values sometimes moving steadily upward and sometimes steadily downward.* The next question is, though, will this same sort of happy effect work within families of compounds‑that is, substances with molecules made up of more than one kind of atom?
* It is only fair to say that a rigidly steady variation is not always found. There are exceptions. However, modern chemists can usually account for them, and we will have an example later in this article.

Consider molecules made up of carbon and hydrogen. These come in many varieties, because carbon atoms can link together in chains and rings. Suppose, then, we consider a single carbon atom combined with hydrogen, a chain of two carbon atoms combined with hydrogen, a chain of three carbon atoms, four, and so on. The longer the chain the larger the molecular weight, and we can consider such a series of progressively larger molecules of very much the same kind to make up a family. What happens to the boiling point, in that case?

TABLE 4
Compound			Molecular Weight	Boiling Point (oK).	
Methane (C-H4)	16.0	111.7	
Ethane (C2-H6)	30.1	184.5	
Propane (C2-H8)	44.1	228.7	
Butane (C4-H10)	58.1	273.7	
Pentane (C5-H12)	72.2	309	
Hexane (C6-H14)	86.2	341	

As you see, boiling point rises smoothly with molecular weight in this case.

To be sure, the family of "hydrocarbons" considered in Table 4 is one in which all the members have molecules made up of the same elements. Would it be possible to set up families in which at least one of the elements changes from member to member?

Thus, carbon is the first member of an element family of which the next three, in order of increasing atomic weight, are silicon (Si), germanium (Ge), and tin (Sn). An atom of each of these higher members can combine with four hydrogen atoms to form well‑known compounds (silane, germane, and stannane, respectively) analogous to methane. Table 5 shows what happens to the boiling points there, and you see we get regularity in such a family, too.

TABLE 5
Compound	Molecular Weight	Boiling Point (oK)	
Methane (CH4)	16.0	111.7	
Silane (SiH4)	32.1	161.4	
Germane (GeH4)	76.6	184.7	
Stannane (SnH4)	122.7	221	

The problem, then, of finding out why water has the high liquid‑range temperatures it has, may become easier to handle if we work within some family of compounds that includes it.

Water molecules are made up of hydrogen and oxygen atoms (H2O). Of these two elements, hydrogen is a 'loner' and is not part of any clearly defined family (though it has certain relationship both to the halogens and to the alkali metals). Oxygen, on the other hand, is the first member of a family that includes sulfur (S), selenium (Se), and tellurium (Te) as later members. An atom of each of these three can combine with two hydrogen atoms to form molecules (H2S, H2Se, and H2Te, respectively) that are analogous in structure to water molecules.

TABLE 6
Compound	Molecular Weight	Boiling Point (oK)	
Water (H2O)	19.0	373.2	
Hydrogen sulfide (H2S)	34.1	213.5	
Hydrogen selenide (H2Se)	91.0	231.7	
Hydrogen telluride (H2Te)	129.6 	271.0	

If we look at the last three members alone, we see that the boiling point goes up with molecular weight. But water doesn't fit! Its boiling point should be, judging from the rest, something like 200oK. or ‑73oC. Only the coldest polar days should suffice to liquefy its vapor and yet here it is, boiling something like 170 degrees higher than it should. Hot water, indeed. 

There are two other compounds that, like water, don't fit their families in this respect.

A hydrogen atom will combine with one atom of any of the halogens. We might expect HF to have a boiling point about 170, but it doesn't. Its boiling point is 292.6 or about 120 degrees "too high."

TABLE 7
Compound		Boiling Point (oK)	
Hydrogen fluoride (HF)		292.6 (~120 too high)	
Hydrogen chloride (HCl)		199.2	
Hydrogen bromide (HBr)		206.5	
Hydrogen iodide (HI)		237.8	

Then, too, three hydrogen atoms will combine with one atom of the member of a family of elements that includes nitrogen (N), phosphor‑us (P), arsenic (As), and antimony (Sb). The first member of the series, ammonia (H3N), ought to have a boiling point of about 150, but it doesn't. Its boiling point is 239.8, which is about ninety degrees "too high."

TABLE 8
Compound		Boiling Point (oK)	
Ammonia (H3N)		239.8 (~90 too high)	
Phosphine (H3P)		185.5	
Arsine (H3As)		219	
Stibine (H3Sb)		256	

What, then, do these three too‑high‑boiling compounds, water (H20), ammonia (H3N), and hydrogen fluoride (HF), have in common?

1) All three are made up of molecules consisting of hydrogen atoms and one other kind of atom.

2) The other atoms involved - oxygen, nitrogen, and fluorine ‑ just happen to be the three most electronegative atoms there are; that is, the atoms most capable of snatching electrons from other atoms.

A fluorine atom, the most electronegative of all, can, for instance, take an electron away from a sodium atom altogether, assuming sole ownership and leaving the sodium atom utterly minus one electron.

The hydrogen atom is not quite the same. It holds on to its single electron more tightly than the sodium atom does to its one outermost electron. The fluorine atom does not take hydrogen's electron away altogether, but it does take over the lion's share of it. The electron, so to speak, is closer to the center of the fluorine atom than to the center of the hydrogen atom.

This means that if you imagine a line drawn down the center of the hydrogen fluoride molecule, with the hydrogen atom on one side and the fluorine atom on the other; the fluorine side, having more than its equal share of electrons, has what amounts to a small negative electric charge, while the hydrogen side has an equally small positive electric charge.

Much the same can be said of the water molecule and the ammonia molecule. In each case, the side of the hydrogen atoms carries a small positive charge, while the side of the oxygen (or nitrogen) atom carries a small negative charge.

All three molecules are "polar molecules." That is, they have poles at which electric charge is concentrated.

This is not true of H2S, for instance, which is otherwise so similar to H2O in structure. Sulfur just isn't as electronegative as oxygen 'and it cannot hog more than its fair share of the electrons of the hydrogen atoms. Hydrogen sulfide is therefore not particularly polar. Neither is hydrogen chloride or phosphine.

If we now consider polar molecules, those with a positively charged end and a negatively charged end, we must inevitably start thinking of the possibility of attraction between molecules. What if the positively charged end of one molecule should be near the negatively charged end of another molecule of the same kind? Would they not stick together a bit?

Yes, they would, particularly since the positively charged end involves the hydrogen atom. Why? Because the hydrogen atom is the smallest of all the atoms and its center can therefore be most closely approached. The strength of attraction between two oppositely charged objects varies inversely as the distance between them. The closer they come together, the stronger the attraction.

It follows, then, that the water molecule, the hydrogen fluoride molecule, and the ammonia molecule are "sticky molecules." They tend to line up positive end to negative end, and it takes significantly higher temperatures to pry them apart than if they were non‑polar; that is, lacking the concentration of charge on two opposite sides, and held together only by the Van der Waals forces mentioned in the previous chapter.*

*Van der Waals forces are also the result of electrical asymmetry in atoms and molecules, with momentary concentrations of electric charge in one place or another. In non‑polar molecules, however, the concentration‑shifts from place to place lead to overall polarizations that are tiny indeed, much less than those in polar molecules, where the concentration of charge is persistent and definitely localized.

Usually, the water molecules are pictured with a hydrogen atom attached to the oxygen atom of its own molecule by a solid bond representing an ordinary chemical linkage, while it is attached to the oxygen atom of a neighboring molecule by a longer, dashed bond to indicate the electromagnetic attraction of opposite charges.

Because the hydrogen atom is thus between two oxygen atoms, one of its own and one of a neighboring molecule (or, in similar fashion, between two fluorine atoms, between two nitrogen atoms, between a nitrogen atom and an oxygen atom, and so on), the situation is commonly referred to as a "hydrogen bond".

The hydrogen bond is only about one twentieth as strong as an ordinary chemical bond, but that is enough to add up to 170 degrees to the temperature required to tear the molecules apart and set the liquid to boiling. Water molecules are sticky enough, thanks to the hydrogen bonds, to boil at 373o K. instead of 200o K., and that, combined with the fact that hydrogen and oxygen are the two most common compound‑forming atoms in the universe, makes it possible for oceans of liquid to exist on a planet the temperature of the Earth.

What's more, it is because of the stickiness of the water molecules that it is possible for water to absorb so much heat for each degree rise in temperature or give off so much heat for each degree fall. We say, therefore, that water has an unusually high "heat capacity."

There is, similarly, an unusually high heat absorption at the melting point or boiling point, due to the necessity for breaking all those hydrogen bonds. That is, it takes much more heat than one would expect to convert ice at 273o K. to water at the same temperature, or to convert water at 373o K. to steam at the same temperature. Working in reverse, an unusual amount of heat is given off when steam condenses to water or water freezes to ice. (Water has an unusually high "latent heat of fusion and vaporization," in other words.)
This is more than a mere matter of statistics. Water acts as a huge heat‑sponge. It takes up and gives off more heat than any other common substance for a given change of temperature, so that the ocean rises in temperature, much more slowly under the beating rays of the Sun than the land does, and drops in temperature much more slowly in the absence of the Sun.

With a vast ocean of water on its surface, the Earth therefore has a much more equable temperature than it would without it. In the summer, the sluggishly warming ocean acts as a cooling device; in the winter, the sluggishly cooling ocean is a warming device. And if you want to see what that means in a practical sense, consider the temperature ranges over the day‑night interval and the summer‑winter interval of a land area far from any ameliorating stretch of ocean (North Dakota) with those of one that is surrounded by ocean on all sides (Ireland).

Since at any temperature, the evaporation of water absorbs more heat per gram of vapor formed than is true of any other common liquid, water is a particularly cheap and effective air‑conditioning device.

Perspiration is almost pure water and as it evaporates a great deal of heat must be absorbed from the object closest to that water ‑ which happens to be the skin on which the perspiration rests. In this way, the body is cooled.

Then, too, there is the matter of solvent properties. In a substance like sodium chloride (common salt), the sodium atoms lose an electron each to the chlorine atoms, which therefore gain an electron each. The sodium atoms carry a unit positive charge and the chlorine atoms a unit negative charge, and are hence called ions. The two sets of ions cling together through the attraction of opposite charges.*
* The sodium chloride combination is much more polar than the water molecule is and this is reflected in its extremely high boiling point.

When particles of salt are dumped into water, the presence of positive and negative poles on the water molecules sets up an electromagnetic field which tends to neutralize that which is set up by the charged sodium and chloride ions. The ions cling to each other with far less verve in the presence of water than in the open air and have a pronounced tendency to fall apart and go swimming in the water on their own. To put it briefly, sodium chloride dissolves in water.

So do a surprising variety of other electrovalent compounds, that is, compounds made up of oppositely charged ions after the fashion of sodium chloride.

Polar compounds, which are not built up of outright ions but have molecules with separated concentrations of charge (like water itself), also lose a considerable part of their tendency to cling together in the presence of water and therefore tend to dissolve. This includes many common substances of importance to life which have the oxygen‑hydrogen or nitrogen‑hydrogen linkage that makes polarization possible.

This includes various alcohols, sugars, amines, and other organic compounds.

No other liquid is so versatile a solvent as water; no other liquid can dissolve appreciable quantities of so wide a variety of substances. For sure, water cannot dissolve appreciable amounts of all electrovalent compounds, since electrovalency is not the only property that is important. And, of course, it cannot dissolve non‑polar compounds such as hydrocarbons, fats, sterols, and so on.

The importance of water's versatile solvent action is this --

The body's most important substances, the proteins and the nucleic acids, together with its most important fuels, the starches and sugars, are loaded with oxygen‑hydrogen and nitrogen‑hydrogen linkages and, if not polar altogether, have important polar regions within their molecules. Such compounds can therefore dissolve in water, or, at least, can attach water molecules intimately to various portions of their structure and undergo changes in connection with these attached water molecules.

In short, the body's chemistry can go on against the intimacy of a water background. This background is so essential to life as we know it, that life could only have reasonably begun in the ocean, and now, even where it has adapted itself to dry land, the tissues remain approximately 70 per cent water.

So consider water. Consider its high liquid‑range temperatures, its capacity to act as a temperature‑ameliorating heat‑sponge and as an efficient air conditioner, its ability to dissolve a wide variety of substances and, therefore, to act as a medium within which the reactions necessary to life can proceed, and you may well say, "Surely, this is no accident. Surely water is a substance that has been carefully designed to meet the needs of life."

But that is placing the cart before the horse, I'm afraid. Water existed, to begin with, as a substance of certain properties, and life evolved to fit those properties. Had water had other properties, life would have evolved to fit those other properties. If water had a lower liquid‑range temperature, for instance, life might have evolved on Jupiter. And if water had not existed at all, life might have evolved to fit some other substance altogether.

In every case, though, life would have evolved so neatly to fit whatever was at hand, that any form of that ]if e high enough to consider the situation with sufficient subtlety would well feel justified in believing that intelligent and purposeful supernatural design was involved. in something which, probably, the blind and random forces of evolution had produced.

------------------------
From Isaac Asimov : The left hand of the electron