The first to realize this clearly was the Danish physicist Niels Bohr. In 1913 Bohr pointed out that if an electron absorbed energy, it had to absorb it a whole quantum at a time and that to an electron a quantum was a large piece of en 'ergy that forced it to change its relationship to the rest of the atom drastically and all at once.

Bohr pictured the electron as circling the atomic nucleus in a fixed orbit. When it absorbed a quantum of energy, it suddenly found itself in an orbit farther from the nucleus - there was no in-between, it was a one-step proposition.

Since only certain orbits were possible, according to Bohr's treatment of the subject, only quanta of certain size could be absorbed by the atom-only quanta large enoug to raise an electron from one permissible orbit to another.

When the electrons dropped back down the line of per missible orbits, they emitted radiations in quanta. They emitted just those frequencies which went along with the size of quanta they could emit in going from one orbit to another.

In this way, the science of spectroscopy was rational ized. Men understood a little more deeply why each ele ment (consisting of one type of atom with one type of energy relationships among the electrons making up that type of atom) should radiate certain frequencies, and cer tain frequencies only, when incandescent. They also under stood why a substance that could,absorb certain frequen cies should also emit those same frequencies under other circumstances.

In other words, Yirchhoff had started the whole problem and now it had come around fuil-circle to place his em pirical discoveries on a rational basis.

Bohr's initial picture was oversimple; but he and other men gradually made it more complicated, and capable of explaining finer and finer points of observation. Finally, in 1926, the Austrian physicist Erwin Schri3dinger worked out a mathematical treatment that was adequate to an alyze the workings of the particles making up the interior of the atom according to the principles of the quantum theory. This was called "quantum mechanics," as opposed to the "classical mechanics" based on Newton's three laws of motion and it is quantum mechanics that is the founda- tion of Modern Physics.

There are fashions in science as in everything else. Con duct an experiment that brings about an unusual success and before you can say, "There are a dozen imitations!" there are a dozen imitations!

Consider the element xenon (pronounced zee'non), dis covered in 1898 by William Ramsay and Morris William Travers. Like other elements of the same type it was iso lated from liquid air. The existence of these elements in air had remained unsuspected through over a century of ardent chemical analysis of the air, so when they finally dawned upon the chemical consciousness they were greeted as strange and unexpected newcomers. Indeed, the name, xenon, is the neutral form of the Greek word for "strange," so that xenon is "the strange one" in all literalness.

Xenon belongs to a group of elements commonly known as the "inert gases" (because they are chemically inert) or the "rare gases" (because they are rare), or "noble gases" because the standoffishness that results from chemi cal inertness seems to indicate a haughty sense of seff importance.

Xenon is the rarest of the stable inert gas and, as a matter of fact, is the rarest of all the stable elements on Earth. Xenon occurs only in the atmosphere, and there it makes up about 5.3 parts per million by weight. Since the atmosphere weighs about 5,500,000,000,000,000 (five and a half quadrillion) tons, this means that the planetary supply of xenon comes to just about 30,000,000,000 (thirty billion) tons. This seems ample, taken in full, but picking xenon atoms out of the overpoweringly more corn,mon constituents of the atmosphere is an arduous task and so xenon isn't a common substance and never will be.

What with one thing and another, then, xenon was not a popular substance in the chemical laboratories. Its chem ical, physical, and nuclear properties were worked out, but beyond that there seemed little worth doing with it. It remained the little strange one and received cold shoulders and frosty smiles.

Then, in 1962, an unusual experiment involving xenon was announced whereupon from all over the world broad smiles broke out across chemical countenances, and little xenon was led into the test tube with friendly solicitude.

"Welcome, stranger!" was the cry everywhere, and now you can't open a chemical journal anywhere without find ing several papers on xenon.

What happened?

If you expect a quick answer, you little know me. Let me take my customary route around Robin Hood's barn and begin by stating, first of all, that xenon is a gas.

Being a gas is a matter of accident. No substance is a gas intrinsically, but only insofar as temperature dictates.

On Venus, water and ammonia are both gases. On Earth, ammonia is a gas, but water is not. On Titan, neither am monia nor water are gases.

So I'll have to set up an arbitrary criterion to suit my present purpose. Let's say that any substance that remains a gas at -1000 C. (-148' F.) is a Gas with a capital letter, and concentrate on those. This is a temperature that is never reached on Earth, even in an Antarctic winter of extraordinary severity, so that no Gas is ever anything but gaseous on Earth (except occasionally in chemical lab oratories).

Now why is a Gas a Gas?

I can start by saying that every substance is made up of atoms, or of closely knit groups of atoms, said groups being called molecules. There are attractive forces between atoms or molecules which make them "sticky" and tend to hold them together. Heat, however, lends these atoms or molecules a certain kinetic energy (energy of motion) which tends to drive them apart,.since each atom or mole cule has its own idea of where it wants to go. [I enjoy sin]

The attractive forces among a given set of atoms or molecules are relatively constant, but the kinetic energy varies with the temperature. Therefore, if the temperature is raised high enough, any group of atoms or molecules will fly apart and the material becomes a gas. At tempera tures over 60000 C. all known substances are gases.

Of course, there are only a, few exceptional substances with interatomic or intermolecular forces so strong that it takes 6000' C. to overcome them. Some substances, on the other hand, have such weak intermolecular attractive forces that the warmth of a summer day supplies enough kinetic energy to convert them to gas (the common anes thetic, ether, is an example).

Still others have intermolecular attractive forces so much weaker still that there is enough heat at a tempera ture of -I 00' C. to keep them gases, and it is these that are the Gases I am talking about.

The intermolecular or interatomic forces arise out of the distribution of electrons within the atoms or molecules.

The electrons are distributed among various "electron shells," according to a system we can,accept without de tailed explanation. For instance, the aluminum atom con tains 13 electrons, which are distributed as follows: 2 in the innermost shell, 8 in the next shell, and 3 in the next shell. We can therefore signify the electron distribution in the aluminum atom as 2,8,3.

The most stable and symmetrical distribution of the electrons among the electron shells is that distribution in which the outermost shell holds either all the electrons it can hold, or 8 electrons-whichever is less. The innermost electron shell can hold only 2, the next can hold 8, and each of the rest can hold more than 8. Except for the situ ation where only the innermost shell contains electrons, * No, I am not implying that atoms know what they are doing and have consciousness. This is just my teleological way of talk ing. Teleology is forbidden in scientific'articies, 1-ut it s'o happens then, the stable situation consists of 8 electrons in the outermost shell.

There are exactly six elements known in which this situ ation of maximum stability exists:

Electron Electron Element Symbol Distribution Total helium He 2 2 neon Ne 2,8 10 argon Ar 2,8,8 is krypton Kr 2,8,18,8 36 xenon Xe 2,8,18,18,8 54 radon Rn 2,8,18,32,18,8 86

Other atoms without this fortunate electronic distribu tion are forced to attempt to achieve it by grabbing addi tional electrons, or getting rid of some they already pos sess, or sharing electrons. In so doing, they undergo chem ical reactions. The atoms of the six elements listed above, however, need do nothing of this sort and are sufficient unto themselves. They have no need to shift electrons in any way and that means they take part in no chemical reactions and are inert. (At least, this is what I would have said prior to 1962.)

The atoms of the inert gas family listed above are so self-sufficient, in fact, that the atoms even ignore one another. There is little interatomic attraction, so that all are gases at room temperature and all but radon are Gases.

To be sure, there is some interatomic attraction (for no atoms or molecules exist among which there is no attrac tion at all). If one lowers the temperature sufficiently, a point is reached where the attractive forces become dom inant over the disruptive effect of kinetic energy, and every single one of the inert gases will, eventually, become an inert liquid.

What about other elements? As I said, these have atoms with electron distributions of less than maximum stability and each has a tendency to alter that distribution in the direction of stability. For instance, the sodium atom (Na) has a distribution of 2,8, I. If it could get rid of the outer most electron, what would be left would have the stable 2 8 configuration of neon. Again, the chlorine atom (CI) b@s a distribution of 2,8,7. If it could gain an electron, it would have the 2,8,8 distribution of argon.

Consequently, if a sodium atom encounters a chlorine atom, the transfer of an electron from the sodium atom to the chlorine atom satisfies both. However, the loss of a negatively charged electron leaves the sodium atom with a deficiency of negative charge or, which is the same thing, an excess of positive charge. It becomes a positively charged sodium ion (Na+). The chlorine atom, on the other band, gaining an electron, gains an excess of nega tive charge and becomes a negatively charged chloride ion ["chlorine ion" as a convention of chemi amp;al nomenclature we might just as well accept with a weary sigh. Anyway, the "d" is not a typographical error] (CI-).

Opposite charges attract, so the sodium ion attracts all the chloride ions within reach and vice versa. These strong attractions cannot be overcome by the kinetic energy in duced at ordinary temperatures, and so the ions hold to gether firmly enough for "sodium chloride" (common salt) to be a solid. It does not become a gas, in fact, until a temperature of 1413' C. is reached. .Next, consider the carbon atom (C). Its electron dis tribution is 2,4. If it lost 4 electrons, it would gain the 2 helium configuration; if it gained 4 electrons, it would gain the 2,8 neon configuration. Losing or gaining that many electrons is not easy,_so the carbon atom shares electrons instead. It can, for instance, contribute one of its electrons to a "shared pool" of two electrons, a pool to which a neighboring carbon atom also contributes an elec tron. With its second electron it can form another shared pool with a second neighbor, and with its third and fourth, two more pools with two more neighbors. Each neighbor * The charged chlorine atom is called "chloride ion" and not ran set up additional pools with other neighbors. In this way, each carbon atom is surrounded by four other carbon atoms.