The Thalassogens

The six most common elements in the universe as a whole are thought to be hydrogen, helium, oxygen, neon, nitrogen, and carbon, in' that order. Out of every 10.000 atoms in the universe about 9.200 are hydrogen, 790 are helium, 5 are oxygen, 2 are neon, 2 are nitrogen, and 1 is carbon. All the rest make up an insignificant scattering and for many purposes can be simply ignored.

With this information on hand, we can ask ourselves what the most common compound (i.e., a substance with a molecule made up of two or more different kinds of atoms) in the universe is. It stands to reason that the most common compound would be one with a small, very stable molecule made up of atoms of the two most common elements.

Since helium atoms don't form parts of any molecules at all, that leaves hydrogen and oxygen as the most common compound‑forming elements in the universe. One atom of each can combine to form "hydroxyl" (OH), which has been detected in the interstellar spaces of our galaxy and of at least one other. It can only exist in rarefied media such as that in space. Two hydrogen atoms and one oxygen form water (H2O), and that can exist at planetary densities ‑ and is undoubtedly the most common such compound in the universe.

Naturally, water wouldn't be common everywhere. It wouldn't exist at all in any normal star, of course. The molecule breaks up at stellar temperatures. On too‑small planetary bodies, water molecules would be too light and flitting to be held by the feeble gravitational force. Some might be held by chemical forces to the rocky crust, but this would represent a very small percentage of the total potential. It is not surprising then that the Moon, Mars, and, undoubtedly, Mercury are relatively dry.

On giant planets such as Jupiter and Saturn, where the gravitational field is intense and the temperature is low, there is a much more representative sampling of the material of the universe and surely water is by far the most common compound on such worlds.

Earth stands in an intermediate position. It is small enough and warm enough to have lost most of the water it might have possessed at the start. More likely, it failed to gather most of it in the first place out of the swirling cloud of dust and gas from which the planet formed. Even so, water on Earth is extremely plentiful.

In fact, in two respects, Earth's water is absolutely unique. In the first place, water is by far the most common liquid on Earth. Indeed, it is the only liquid on Earth present in quantity. (What is in second place? Petroleum, perhaps.)

Secondly, water is the only substance on Earth present, in quantity, in all three phases, solid, liquid, and gas. Not only is there an ocean full of water, but there are polar caps of miles‑deep ice, and there is water vapor making up a major (if variable) part of the atmosphere.

The question is then: Can any substance other than water serve? Can a planet exist with a large ocean of any substance other than water?

To answer that question, let's consider the requirements:

1) The ocean substance must be a plentiful component of the universe mixture. We can imagine oceans of liquid mercury, or liquid fluorine or liquid carbon tetrachloride, but we can't realistically imagine any planet with these particular substances present in such quantities as to spread out into oceans.

2) The ocean substance must have a prominent liquid phase. For instance, the Martian polar caps may well be frozen carbon dioxide, but there is no liquid carbon dioxide phase at Martian atmospheric pressure. The solid carbon dioxide vaporizes directly to gas, so there would be no carbon dioxide ocean even if there were enough carbon dioxide to form one.

3) Ideally, we would want a substance whose liquid phase could be transformed with reasonable ease to either solid or gas, if we are to make possible those properties of Earth's ocean which lead to ice caps, clouds, rain and snow. Thus 'an ocean of liquid gallium at the temperatures of water's boiling point, for instance, might produce gallium "ice caps" with ease, but at that temperature, gallium's vapor pressure would be so low that there would be no gallium vapor in the air to speak of, no gallium clouds, no gallium rain. On the other hand, if we had an ocean of liquid helium at a temperature of 2o above absolute zero (i.e., 2o K.) there would be plenty of helium vapor in the atmosphere (indeed, that would make up almost all the atmosphere) and helium rain would be common, but there is likely to be no helium ice or snow because solid helium doesn't form, even at absolute zero, except under considerable pressure, and we would be hard put to design a planet with sufficient atmospheric pressure at 2oK. to do the job.

In considering the requirements, let's begin with the first ‑ presence in oceanic quantities. For that, we had better work with the top six elements only: hydrogen, helium, oxygen, neon, nitrogen, and carbon. Any substance made up of anything but these six elements (singly or in combination) might have many virtues but would simply not be present in sufficiently overwhelming a quantity to make up an ocean composed entirely or nearly entirely of itself.*

* There is one conceivable exception on an Earthlike planet. Silicon dioxide is present in oceanic quantities but it is a solid and wouldn't be a liquid under anything but white heat. Scratch silicon dioxide.

Of these six elements, two, helium and neon, can exist in elemental form only. A third, hydrogen, can form compounds, but exists in such overwhelming quantities that on any planet capable of collecting more than a trace of it (i.e., on Jupiter, as opposed to Earth) it must exist mostly in elemental form for sheer lack of sufficient quantities of other elements with which to combine.

As for oxygen, nitrogen, and carbon, these, in the presence of a vast preponderance of hydrogen will exist only in combination with as much hydrogen as possible. Oxygen will exist as water (H2O) ‑ nitrogen, as ammonia (H3N); and carbon, as methane (H4C).

This gives us our list of the six possible thalassogens**: hydrogen, helium, water, neon, ammonia, and methane, in order of decreasing quantity.

** This is a word I have just made up. It is from Greek words ("sea‑producers") and I define it as "a substance capable of forming a planetary ocean."

The next step is to consider each in connection with its liquid phase. At ordinary pressures, equivalent to that produced by Earth's atmosphere, each has a clear‑cut boiling point temperature, above which it exists only as a gas. This boiling point can be increased when pressure is increased, but let's ignore that complication, and consider the boiling point, in degrees above absolute zero, at ordinary pressure.

It turns out that the boiling points of helium, hydrogen, and neon are, respectively, 4.2o K., 20.3o K., and 27.3o K.

But keep in mind that even distant Pluto has a surface temperature estimated to be roughly 60o K. In fact, I wonder if any sizable planet, such as the outer members of our solar system, can ever have extremely low temperatures. Internal heat arising from radioactivity must be sufficient to keep the surface temperature at Plutonian levels, at least, even in the complete absence of any sun. (Jupiter, for instance, according to a recent report I've seen, radiates three to four times as much heat as it receives from the Sun.)

In short, then, for any reasonable planet we can design, the temperature is going to be too high for the presence, in quantity, of helium, hydrogen, or neon in the liquid phase. Scratch them from the list and we have only three thalassogens left: methane, ammonia, and water.

And what are their boiling points? Why, respectively, 111.7, 239.8, and 373.2 oK.

If we consider these three, we come to these conclusions:

1) Water is the most common and is therefore the most likely to form an ocean.

2) Since methane is liquid across a range of 23 degrees, ammonia across 44 degrees, and water across 100 degrees, water, of the three, has by far the broadest temperature range for the liquid phase and, in its ocean‑forming propensities, is least sensitive to temperature deviation.

3) Most important of all, water forms its oceans at a higher temperature than the other two. You might expect methane oceans on a planet like Neptune or ammonia oceans on a planet like Jupiter. Only water, however, only water, could possibly form an ocean on an inner planet like Earth.

Well then, we depend for the existence of our ocean, and therefore for the existence of life, on the fact that water happens to have its liquid range at a far higher temperature than that of any other possible thalassogen. Is that just the way the ball bounces or is there something interesting to be wrung out of the water molecule?

Let's see --

When atoms combine to form molecules, the bond between them is formed through a kind of tug‑of‑war over the outermost electrons in those atoms. In many cases, one type of atom has the capacity to hold on to one or two electrons over and above those it normally possesses. Given half a chance it will grab on to such electrons. Since the atom itself is electrically neutral (positive charges in the interior, balancing negative charges on the outskirts) and since every electron has a negative charge, an atom which is capable of taking on one or more additional electrons then carries a net negative charge. Elements made up of atoms capable of doing this are therefore characterized as "electronegative."

The most electronegative of the elements, by far, is fluorine. Following it, in order, are oxygen, nitrogen, chlorine, and bromine. These are the only strongly electronegative elements.

Some atoms on the other hand have no strong ability to latch on to additional electrons. Indeed, they find it difficult to hold on to the electrons they normally possess and have a considerable tendency to give up one or two. Given half a chance they will do so. Once they lose such negatively charged electrons, what remains of the atom has a net positive charge. Such atoms are therefore "electropositive."

Most of the elements tend to be somewhat electropositive. The most electropositive elements are the alkali metals, of which sodium and potassium are the most common representatives. Calcium, magnesium, aluminum, and zinc are other examples of strongly electropositive elements.

When an electropositive element, like sodium, meets an electronegative one like chlorine, the sodium atom freely gives up an electron, which the chlorine atom as freely takes. What is left is a sodium atom with a positive charge (a sodium ion) and a chlorine atom with a negative charge (a chloride ion). The attraction between the two ions is the strong pull of an electromagnetic force and this is called "electrovalence." A number of chloride ions cluster around each sodium ion and a number of sodium ions cluster around each chloride ion. The result is an intricate and very orderly array of ions that hang on to each other tightly.

The commonest way of pulling ions apart is to use heat.

All ions, no matter how firmly held in place by some sort of attraction, are vibrating about that place. This vibration is related to temperature. The higher the temperature, the more energetic the vibration. If the temperature is high enough, the vibration becomes violent enough to pull the ions apart, however strong the electromagnetic force between them, and the substance then melts. (In the liquid phase, the ions are no longer held firmly in place and they move about freely.)

Nevertheless the temperature, by ordinary standards, must be quite high before the strong attractions between the sodium ions and the chloride ions can be overcome. Sodium chloride (ordinary table salt) has a comparatively high melting point, therefore: 1074oK. (For orientation, a pleasant spring day with the temperature at 70oF. is at 294oK.)

Still higher temperatures are required to pull the ions apart altogether and send them in pairs (one sodium ion and one chloride ion) into the nearly total independence of the gas phase, so that the boiling point of sodium chloride is 1686oK.

This is more or less true of all electrovalent compounds, which form by the transfer of one or more electrons from one atom to another. Molybdenum oxide has a melting point of 2893 oK and a boiling point of 5070 oK.

What happens, though, when one electropositive element meets another? Sodium atoms, for instance, can form bonds among themselves by allowing the outermost electron each possesses (and which they hold on to only very loosely) to be shared among them all. This is a stabler situation than would exist if each were responsible for its own outermost electron only, as in sodium gas. Consequently sodium atoms cling together and sodium is a solid at ordinary temperatures. To be sure, it doesn't take much to pull the atoms apart and sodium melts at a temperature of 370oK., just under that of boiling water. It doesn't boil, though, and obtain complete atomic independence till 1153oK.

(Those outermost electrons wander easily from atom to atom. Their existence accounts for the fact that sodium, and metals generally, conduct heat and electricity so much better than non‑metals.)

Metals made up of less electropositive atoms get together more snugly and some of them end up by forming bonds as tight as those of any electrovalent compound. Tungsten metal has a melting point of 3640oK. and a boiling point of 6150oK.

Yet though metallic atoms fit together well, there is a greater tendency for them to transfer electrons to the electronegative atoms, particularly to oxygen, which is by far the most common of all the strongly electronegative elements. For this reason, there is virtually no free metal in the Earth's crust.*

* Earth has a metallic core because it contains so much iron that there just aren't enough electronegative atoms to take care of it all. The metallic excess, denser than the oxygen‑containing electrovalent compounds, settled to the Earth's center in the soft, youthful days of the planet.

In general, then, we can say that metals and electrovalent compounds are so high‑melting as to offer no chance of a liquid phase at any reasonable planetary temperature, up to and including that of Mercury. Those few which might (like sodium metal or tin tetrachloride) cannot possibly be present in large enough quantity to form an ocean.

So we must look for something else. What happens if one electronegative atom meets another? What happens if one fluorine atom meets another, for instance? Each of the fluorine atoms can handle one electron over and above its usual assignment, but neither is in a position to give up one of its own in order to satisfy the other. What does happen is that each atom allows the other a share in one of its own electrons. There is a two‑electron pool to which each contributes and in which each shares. Both fluorine atoms are then satisfied.

In order for this pool to exist, though, the two fluorine atoms must remain at close quarters! To pull them apart takes a lot of effort, for it means breaking up that two-electron pool. Consequently, under ordinary circumstances, fluorine in elementary form exists in molecules made up of atom pairs (F2). The temperature must rise well over 1300oK, even to begin to break up the fluorine molecule and shake the individual atoms apart. The attraction between atoms represented by shared electrons is called a "covalent bond."

Two fluorine atoms, once they have formed their two-electron pool, have no reason to share any electrons with any other atoms, much less transfer electrons to them or even receive electrons from them. The two‑electron pool completely satisfies their electron needs. Consequently, when one fluorine molecule meets another fluorine molecule, they bounce off each other, with very little tendency to stick together.

If there were no tendency to stick together at all, the fluorine molecule would remain independent of its neighbors however far down the temperature might drop. The molecules would move more and more sluggishly, bounce off one another more and more feebly, but they would never stick.

However, there are what are called "Van der Waals forces," named for the Dutch chemist who first studied them. Without going into the matter in detail we can simply say that there are weak attractive forces between atoms or molecules even when there is no outright electron transfer or electron sharing.

Thanks to Van der Waals forces, fluorine molecules are slightly sticky, and if the temperature drops low enough, the energy that keeps them moving will not be great enough to make them break away after colliding. Fluorine will condense to a liquid.

The boiling point of liquid fluorine is 85 oK. If the temperature drops further still, the fluorine molecules lock firmly into an orderly array and fluorine becomes a solid. The melting point of solid fluorine is 50 oK.

The same thing happens with the other electronegative elements. Chlorine, oxygen, and nitrogen also form electron pools between two atoms. We therefore have chlorine molecules, oxygen molecules, and nitrogen molecules, each made up of atom pairs (Cl2 O2, and N2). Even hydrogen atoms, which are not particularly electronegative, form molecules by pairs (H2).

In every case the melting and boiling points are low, with the exact value depending on the strength of the Van der Waals forces. Hydrogen, with its very small atoms, possesses a liquid range at a considerably lower temperature than that of fluorine. The boiling point of liquid hydrogen is 21oK, and the freezing point of solid hydrogen is 14 oK.

A few varieties of atom happen to possess a satisfactory number of electrons to begin with. They have little tendency to give up any electrons they have and still less to accept additional electrons from outside. They do not therefore tend to form compounds. These are the so‑called "noble gases."

There are six of these altogether and, of them, the three with the largest atoms can form compounds (not very stable ones) with the most electronegative elements, such as fluorine and oxygen. The three with the smallest atoms ‑ argon, neon, and helium (in order of decreasing size) won't do even that much under any conditions yet discovered. Nor will they form electron pools among themselves. They remain in sullen isolation as individual atoms.

Yet they, too, experience the mutual attraction of Van der Waals forces and, if cooled sufficiently, become liquids. The smaller the atom, the smaller the forces and the more strongly cooled they must be to liquefy.

Helium, with the smallest atoms of the noble gases, experiences such small attractions that of all known substances it is the most difficult to liquefy. The boiling point of liquid helium is phenomenally low, only 4.2 oK. Solid helium doesn't exist at all, even at 0 oK. (absolute zero), except under considerable pressure.

So far, though, these gaseous substances I have discussed, that are covalent in nature, and that have liquid ranges far down the temperature scale, are all elements - elements that either exist in the form of isolated atoms, as in the case of helium, or as isolated two‑atom molecules, as in the case of hydrogen.

Is it possible for molecules of two different atoms to be covalent in nature and to be low‑melting and low‑boiling. 

‑Yes, it is!

Consider carbon. The carbon atom is neither strongly electropositive nor strongly electronegative. It has a tendency to form two‑electron pools with each of four other atoms. It could form those pools with four other carbon atoms, each of which can form pools with three others, each of which with still three others, and so on indefinitely. In the end, uncounted trillions of carbon atoms may be sticking firmly together by way of strong covalent bonds. The result is that carbon has a higher melting point than that of any other known substance‑nearly 4000o K.

But the carbon atom may form a two‑electron pool with each of four different hydrogen atoms. The hydrogen atoms can only form one two‑electron pool apiece and so that ends it. The entire molecule consists of a carbon atom surrounded by four hydrogen atoms (H4C), and this is methane.

Methane molecules have little attraction for each other except by way of weak Van der Waals forces. The boiling point of liquid methane is 112 oK, and the melting point of solid methane is 89 oK.

Similarly, a carbon atom can form a molecule with one oxygen atom. This would be carbon monoxide (CO). Its boiling point and melting point are, respectively, 83 oK and 67 oK.

Now we can come to a general conclusion. Unlike metallic substances and electrovalent compounds, covalent compounds have low melting points and boiling points and only they can conceivably be thalassogens at reasonable planetary temperatures.

This gives us our first answer as to why water is a thalassogen at all: it is a covalent compound essentially. All right, that's something to begin with. Yet so many covalent compounds are, if anything, liquid at too low a range for planetary purposes and certainly for earthly purposes specifically. Why is liquid water so warm then?

One possibility rests in the fact that, in general, the larger the covalent atom or molecule, the stronger the Van der Waals forces and the higher the boiling point. Consider the following table, in which the size of the molecule is measured by its molecular weight (or, in the case of helium and neon, atomic weight).

The table isn't perfect, for helium, which has a larger atomic weight than hydrogen's molecular weight, nevertheless has a lower boiling point than hydrogen. Then, too, fluorine, which has a larger molecule than oxygen has, is nevertheless lower‑boiling. Still, the table seems to show that there is a kind of rough and ready relationship between molecular weight and boiling point in the case of covalent compounds.

Substance	Atomic or Molecular Weight	Boiling Point (oK)	
Hydrogen (H2)	2	17	
Helium (He)	4	4	
Neon (Ne)	20	27	
Nitrogen (N2)	28	77	
Carbon monoxide (CO)	28	83	
Oxygen (02)	32	90	
Fluorine (F2)	38	85	
Oxygen fluoride (OF2)	54	138	
Nitrogen fluoride (NF3)	71	153	
Chlorine (Cl2)	71	239	
Pentane (C5H12)	72	309	
Chlorine heptoxide (Cl2O7)	183	355	

We might conclude therefore that water, which has a boiling point at 373 oK, ought to have a molecular weight somewhat higher, or at least not particularly lower, than chlorine heptoxide. Its molecular weight ought to be, say, 180, as a minimum.

Except that it isn't. The molecular weight of water is 18, just one tenth what it "ought" to be.

Something, obviously, is terribly wrong - or right, perhaps, for it is to whatever causes this anomaly that we owe our life‑giving ocean. What that wrongness/ rightness might be we'll discuss in the next chapter.