FIG 1Fig. 1.
Fig. 1.
The result was encouraging, and led to the probability of the nitrogen being altered in some way, or of the presence of some new component of the atmosphere. An experiment was therefore begun on a larger scale, the atmospheric nitrogen being passed backwards and forwards from one large glass gasholder A to another B, through a tube filled with magnesium heated to redness G, to absorb nitrogen; over red-hot copper oxide (a) (b), so that any carbonaceous matter such as dust should be oxidised to carbon dioxide and water; and these, if produced, were absorbed by placing in the train of tubes, one filled with a mixture of soda and lime F and I, to absorb any carbon dioxide which might possibly be formed, and two filled with pentoxide of phosphorus D and H, to dry the gas, so that water-vapour, carried along with the gas from the gasholders (which contained water) might be removed before the gas passed over the red-hot magnesium; for water acts on hot magnesium, forming oxide of magnesium and hydrogen, and the gas would have become contaminated with the latter had this precaution not been taken.
The process was continued for ten days, by which time most of the nitrogen had become absorbed. The apparatus was then somewhat altered, so as to make it possible to work with a smaller quantity of gas; but the tubes destined to absorb nitrogen, hydrogen, etc., were filled with the same materials as before. In a few days more the volume was reducedto one-seventh of what it had been when the transference to the smaller apparatus was made, or about one-eightieth of the original volume of the atmospheric nitrogen taken.
The gas was then weighed, this time in a larger bulb, the weight being 0·2190 gram; and such is the possibility of precision in weighing on a good balance, that a difference of one two-thousandth of the whole weight was detectable. The density of the gas was now found to be 16·1. At this stage it was still believed that the new gas was an ozone-like modification of nitrogen, difficult to attack by magnesium. It was supposed that just as oxygen, when exposed to an electric discharge, undergoes a cleavage of its molecules, two-atom molecules becoming one-atom molecules for an instant, which then unite to form three-atom molecules, so the action of the magnesium on the nitrogen might be to withdraw one atom of nitrogen from the two-atom molecule, leaving a single uncombined atom, which might not improbably find two partners, each of its own kind, to form with them a three-atom molecule—a sort of nitrogen-ozone, in fact. Hence it was resolved to continue the absorption with fresh magnesium for a still longer time, in the hope of its being possible to isolate the three-atom nitrogen molecules. Butit became apparent that the bright metallic magnesium was now not much attacked; and on estimating the total amount of nitrogen absorbed, by treating the compound of nitrogen and magnesium with water, and liberating the nitrogen as ammonia, it appeared that only a small quantity of magnesium nitride had been formed. The density of this further purified gas was again determined, when it was found that a litre now weighed 1·7054 gram, corresponding to a density of 19·086.
A portion of this gas was mixed with oxygen and exposed to a rain of electric sparks in presence of caustic soda; in fact, Cavendish’s old plan of causing nitrogen to combine was now resorted to. Contraction occurred, and on removing the excess of oxygen, the diminution of volume was found to amount to 15·4 per cent of the original volume taken. Making the supposition that the gas of density 19 still contained nitrogen, and allowing for its influencing the density, it followed that the pure gas should be twenty times as heavy as hydrogen.
A tube such as is usually employed in examining the spectra of gases at low pressures was next filled with the gas of density 19. Such a tube, called a Pflücker’s tube, after its inventor, contains wires of platinum sealed through at each end, where it is about half an inch in width; the middle portion of the tube is about 3 inches long, and its bore is a fine capillary. When the platinum wires are connected with the secondary terminals of a Ruhmkorff’s coil, and the tube is partially exhausted, a brilliant glow appears in the capillary portion. If viewed through a glass prism, different gases show different sets of coloured lines crossing the usual gradation of colours of the spectrum. Thus hydrogen exhibits three striking lines, one bright red, one peacock blue, and one violet; nitrogen shows a large number of somewhat hazy bands, red, orange, yellow, and yellow-green in colour, besides a number of bands of a violet colour; but the new gas, while exhibiting the bands characteristic of nitrogen, showed in addition certain groups of red and green lines which did not appear to belong to the spectrum of any known gas.
FIG 2Fig. 2.
Fig. 2.
While these experiments were in progress, Lord Rayleigh was occupied in preparing nitrogen from other sources, and in determining its density; and in every case it was evident that nitrogen from all sources except the atmosphere weighed somewhat less than atmospheric nitrogen. He therefore proceeded to repeat Cavendish’s experiment, and like Cavendish, he obtained a small residue of gas which would not disappear on sparking with oxygen, in presence of caustic soda. The sparks, as they passed, could be observed through a spectroscope (which consists of an arrangement of prisms and lenses so designed as to examine the components of the light emitted by the sparks), and he, too, was struck with the unusual character of the spectrum. His experiments proved, besides, that the amount of residue was roughly proportional to the amount of air taken; thus, beginning with 50 cubic centimetres of air, the residue was 0·32 cubic centimetre; and from 5 cubic centimetres of air, only 0·06 cubic centimetre of gas was obtained.
These small amounts are not proportional to the quantities of air taken; but, as will afterwards be seen, the discrepancy is owing to the solubility of the new gas in water. Still they served to show that from a comparatively large amount of air, more of the new gas could be obtained than from a smaller amount.
At this stage the two discoverers joined forces, and letters passed almost daily between them, describing the results of experiments which one or other had made. And just prior to the meeting of the British Association at Oxford in August 1895, it was decided that the proof of the existence of a new constituent gas in air was sufficiently clear to render it advisable to make to the Association a short announcement of the discovery. The statement was received with surprise and interest; chemists were naturally somewhat incredulous that air, a substance of which the composition had been so long and so carefully studied, should yield anything new. One of the audience inquired whether the name of this new substance had been discovered; as a matter of fact it was then under consideration.
But it was still conceivable, although improbable, that the new gas was being produced by the very processes designed for its separation, and attention was first turned to devising a complete proof of its actual presence in air. Now it is known that the rates of diffusion of gasesthrough a narrow opening, or through a number of minute holes, such as exist in a pipe of porous clay,e.g.a tobacco-pipe stem, are in inverse proportion to the square roots of the densities of the gases. Oxygen is, in round numbers, sixteen times as dense as hydrogen; the square roots of 16 and 1 being 4 and 1, it was found by Graham, who first carefully investigated this subject, that four times as much hydrogen would pass through a porous diaphragm, in a given time, as oxygen. Thecompoundof hydrogen and oxygen, however, in the state of gas,viz.steam, is not separated by such a process into its constituents; it diffuses as such, and since it is nine times as dense as hydrogen, the relative rates of diffusion of steam and hydrogen are as 1: √9, or as 1 to 3; that is, for every 3 parts of hydrogen passing through such a septum, 1 part of steam would pass in the same time.
An experiment was therefore devised, in which a large quantity of air was made to stream slowly through a long train of stems of churchwarden tobacco-pipes, placed inside a glass tube, the latter being closed at each end, except for the entrance and exit tubes of the tobacco-pipes;in the encasing glass tube a vacuum was maintained, and the gases, passing through the walls of the pipe-stems, were pumped off and discharged. According to what has just been said, these should be the lighter gases, nitrogen and oxygen, which ought to pass through the porous stems more quickly than the supposed heavier constituent of air; while the air issuing from the end of the train of pipes should contain relatively more of the heavier constituent, and should in consequence have a greater weight than an equal volume of air. But it was obviously convenient to remove the oxygen before weighing this sample of altered air, and this was done in the usual way by passing the mixed gases over red-hot copper. It was found that such nitrogen was even heavier than ordinary atmospheric nitrogen; not much, it is true, but still consistently heavier. The denser constituent could, in fact, be concentrated by this means. The proof was therefore indubitable that the new gas existed in air as such.
There is another method of proof, however, which was not left untried. Experiment showed that the solubility of the new gas in water is considerably greater than that of nitrogen, although less than that of oxygen. In 100 volumes of water at the ordinary temperature, about 1·5volumes of nitrogen will dissolve, about 4·5 volumes of oxygen, and about 4 volumes of the new gas, to which the name finally chosen for it, “argon,” may now be applied. Now the proportion in which the constituents of a mixture of gases will dissolve in a solvent is conditioned first by their relative solubilities, and second, by their relative proportion. Thus, if air be considered to be simply a mixture of 1 volume of oxygen and 4 volumes of nitrogen, the gas extracted from water which has been shaken with air will have the composition—
So that the proportion of oxygen to nitrogen in such a mixture of gases is considerably greater than in air: instead of being approximately 1 to 4, it is nearly 4·5 to 6. The discovery of this law concerning the composition of the gases dissolved in liquids was due to Dr. Henry, one of the biographers of Dalton.
The gases can be almost entirely extracted by boiling the water. But to boil large quantities of water at one operation in a vessel suitable for collecting the escaping gas is not easy. It is much simpler to cause the water to pass slowly through a can below which there is apowerful flame, so that the water in its passage becomes heated to the boiling-point, and gives off its gas before it escapes. Of course the gas collected contained oxygen, but this was easily removed by the usual method of passing it over red-hot copper. The density of the residual gas was determined, and it was found to be at least as much greater than that of “atmospheric” nitrogen as the density of “atmospheric” nitrogen exceeded that from chemical sources. Hence it was to be concluded that the new constituent of air, argon, was being concentrated by dissolving air in water, and extracting the dissolved mixture of gases. A third proof that argon exists in air will be given farther on.
FIG 3Fig. 3.
Fig. 3.
In order that the properties of the newly-discovered gas, argon, might be thoroughly investigated, it was necessary to prepare it on a much larger scale than had hitherto been attempted, and this was carried out by the two processes for removing the oxygen and nitrogen which have been already described. Supposing the new gas to have the density 20 compared with oxygen as 16, the density of the atmospheric mixture of nitrogen and argon compared with that of nitrogen alone shows that airshould, roughly speaking, contain less than one part of argon in one hundred. Hence, to obtain a litre of argon, it was necessary to work up a large quantity of atmospheric nitrogen. Now, as has just been said, there are two ways of doing this. (1) One is to produce an electric flame between two pieces of stout platinum in air, confined in a large glass balloon of about 6 litres capacity, over a weak solution of caustic soda. For this purpose a very powerful rapidly alternating current is necessary. The latest, and apparently the best, method of carrying this out, was described by Lord Rayleigh in his RoyalInstitution lecture in January 1896. The neck of the balloon is placed downwards, and connected by means of a glass tube, passing through a cork which closes the neck, with a rotating fan or paddle-wheel with curved blades, which forces through the tube a weak solution of caustic soda; another tube, also entering through the cork, conveys away the excess of soda to the fan, whence it is again forced into the balloon. The soda solution makes a fountain in the balloon, and flows in a uniform stream down its sides, covering its inner surface with a thin layer of liquid. Through the cork the two electrodes, with their thick platinum terminals, enter; and there is another tube besides, which conveys into the balloon a mixture of air and oxygen in such proportions that they combine completely on exposure to the flame. The layer of soda solution plays a double part. It prevents the undue heating of the glass balloon, which otherwise must be sunk in running water in order to keep it cool; and it exposes a very large and constantly renewed surface of soda to the nitrous fumes which are produced by the combination of the nitrogen and the oxygen, and so removes them as quickly as they are formed. It appears probable thatthe union results initially in the formation of nitric oxide, NO, which then unites partially with oxygen to form some nitrogen peroxide, NO2. This is absorbed by the soda, giving a mixture of nitrite and nitrate of sodium, NaNO2and NaNO3. Working in this way, from 7 to 8 litres of mixed gases can be made to combine per hour. The rapidly alternating current is best obtained by the use of a transformer; and as the heating effect on the platinum terminals is very great, they must be made of stout rods.
(2) To prepare a large quantity of argon by the absorption of atmospheric nitrogen by magnesium is a somewhat tedious process. The air must be first freed from oxygen by means of red-hot copper, and the atmospheric nitrogen collected in a gasholder. Long tubes of combustion-glass tubing, which stands a bright red heat without becoming deformed, are packed with magnesium turnings and heated to redness in long gas furnaces, such as are used in organic analyses; and through these the “atmospheric nitrogen,” dried by passage over soda-lime and phosphorus pentoxide, is then passed. The magnesium begins to glow at that end of the tube nearest the entrance, owing toits combination with nitrogen, and a hot ring is seen to travel slowly down the tube to the other end, marking the place where such combustion is in progress. The gas issuing from the tube is collected in a small gasholder. When one tube of magnesium is exhausted, another is substituted for it. Each tube is capable of absorbing about seven litres of nitrogen, so that to obtain a litre of argon about one hundred litres of “atmospheric nitrogen” must be employed, and about fourteen tubes of magnesium are required. M. Maquenne, who has prepared the nitrides of several metals, has stated that a mixture of lime and magnesium, yielding metallic calcium, is more easily manipulated than pure magnesium, owing to the absorption of the nitrogen at a lower temperature. The process has not been tried on a large scale, but if the temperature of combination of magnesium and nitrogen could be thus reduced, it would much facilitate the operation, for the greatest care has to be taken not to overheat the combustion tube, else it softens, and blows into holes. Porcelain tubes are attacked by the magnesium, and crack on cooling; and iron tubes are difficult to clean out.
This preliminary operation does not yield pure argon; it merely removes a large portion of the nitrogen. To free the argon from the remainder, it is caused to circulate (by means of a specially contrived mercury-pump, where each drop of mercury in falling down a narrow glass tube carries before it a small bubble of gas) through tubes containing red-hot copper, red-hot copper oxide, red-hot magnesium, and cold soda-lime and phosphoric anhydride. The copper serves to remove traces of oxygen; the copper oxide yields up its oxygen to any hydrogen or carbon compound—dust and the like—which may happen to be present; the soda-lime absorbs any carbon dioxide produced by the combustion of the carbon compounds, and at the same time partially dries the gas; while the phosphoric anhydride effectually dries the gas, previous to its passage over the red-hot magnesium, which in its turn removes the nitrogen. It is necessary to continue this circulation for several days before the litre of gas is entirely freed from nitrogen.
It is difficult to choose between these two methods: both are troublesome, and require a considerable time, but in an ordinary laboratory the latter is probably the more easily set in operation, forthe former requires a suitable electric current, and power, so as to rotate the water-fan. Up to the present date, the only sources which have yielded argon are atmospheric air, gases extracted from mineral waters or from springs, one meteorite, and a few rare minerals. No animal or vegetable substance appears to contain it. Experiments were made in the summer of 1895 by Mr. George MacDonald and Mr. Alexander Kellas, in order to decide whether argon was a constituent of any living matter. Some peas were reduced to powder and dried; the carbon and hydrogen of the peas were burned to carbon dioxide and water by heating with oxide of copper, and under these circumstances the nitrogen is evolved in the state of gas. Had argon been contained in the vegetable, it too would have accompanied the nitrogen. The nitrogen was then, as usual, absorbed for the most part by means of magnesium, and the small unabsorbed residue was mixed with oxygen and exposed to electric sparks for many hours, in presence of caustic soda. There wasnoresidue left after absorbing the excess oxygen: the gas was completely removed. Similar experiments carried out on animal tissue led to a similar conclusion. Two mice were chloroformed, and when deadthey were dried in an oven until all the moisture of their bodies was completely driven off, and it was possible to reduce them to powder. It is interesting to note that one of these mice contained 73 per cent of water, and the other 70·5 per cent. The dried animals yielded about 11 per cent of their weight of nitrogen. Absolutely no residue of gas was obtained on causing this nitrogen to combine; hence it appears to be a legitimate conclusion that neither animal nor vegetable tissue contains any appreciable amount of argon.
But these experiments lead to a further result. They show that nitrogen, procured from its compounds, when treated in the same way as atmospheric nitrogen, yields no trace of argon. And it must therefore be taken as proved without doubt that argon is actually present in the atmosphere as such, and is not produced by any process to which the nitrogen has been submitted in order to extract it.
This point having been settled, the actual percentage of argon in atmospheric air next invited inquiry. It is by no means very easy to absorb quantitatively the whole of the nitrogen from an accuratelymeasured sample of air, for small gains and losses are apt to occur. It is necessary to keep the air out of contact with water as much as possible, because argon, being more soluble than nitrogen, dissolves in larger proportional amount in the water, and is thereby partially removed. The air was therefore entirely manipulated over mercury. The processes were like those previously employed: most of the nitrogen was removed with magnesium, and the residue was freed from all nitrogen by sparking with oxygen. Experiments directed to this end were carried out by Mr. Kellas in Professor Ramsay’s laboratory, and independently by M. H. Schloesing in Paris. The results were identical. “Atmospheric nitrogen” consists of pure nitrogen mixed with 1·186 per cent of its volume of argon.
It is now possible, knowing the percentage of argon in atmospheric nitrogen and its density (19·94), to calculate whether Lord Rayleigh’s determinations of the density of atmospheric nitrogen were correct. The weight of one litre of pure nitrogen is 1·2511 gram, and of argon, 1·7818 gram; hence a litre of a mixture of 98·814 volumes of nitrogen with 1·186 volume of argon must possess the weight 1·2574 gram. Theactual number found by Lord Rayleigh was 1·2572 gram, which is almost exactly identical with the number calculated.
Mineral waters, as a rule, contain small quantities of argon mixed with oxygen, nitrogen, carbon dioxide, and in some cases sulphuretted hydrogen and helium, a gas of which more hereafter. The waters actually examined were the Bath waters, which contain much nitrogen, a little argon, and a trace of helium; the Buxton waters, containing nitrogen and a little argon; the water from “Allhusen’s Well,” Middlesborough, which evolved gas of an inflammable nature consisting mainly of nitrogen, but also containing marsh-gas, and argon to the extent of O·4 per cent; water from boiling springs in Iceland evolved gas containing somewhat more argon than air does, viz. 1·14 per cent; and lastly, water from the Harrogate sulphur springs yielded a gas largely consisting of a mixture of sulphuretted hydrogen, carbon dioxide and nitrogen, but giving also an appreciable amount of argon. Such determinations show that argon is not merely confined to the atmosphere above the earth, but that it penetrates the earth and is contained insubterraneous water. These results have been obtained by Lord Rayleigh, Professor Ramsay, Mr. Travers, and Mr. Kellas.[26]
Similar experiments have been made by Dr. Bouchard in Paris[27]on effervescing waters from Cauterets in the Pyrenees. One of those springs yielded a mixture of nitrogen with a small amount of argon and helium; another yielded only nitrogen and argon; while a third gave nitrogen and helium. Such are, up to the present, the sources of argon. It has been several times stated that the element helium, which is closely allied to argon in its physical properties, and in its inertness, is a normal constituent of our atmosphere, although in very small amount. This, however, is not the case. For the argon of the atmosphere has been very carefully examined for helium in two ways. First, Lord Rayleigh dissolved a large quantity of argon in water, leaving only a minute bubble undissolved. Now while 1000 parts of water dissolve 40 parts of argon, they dissolve only 7 parts of helium; and if helium were present, it should be found in the residue. But carefulspectroscopic examination failed to reveal the characteristic lines of helium. Next, Professor Ramsay and Dr. Collie diffused argon fractionally; and as the densities of argon and helium are very different (2 and 20), the helium should have been visible in the first portion which diffused. There was no trace to be detected. But, further, they submitted this portion of argon to a discharge of electricity for several hours in a vacuum-tube provided with platinum electrodes. This process, as they proved, carries out the helium with the platinum which is splashed on to the sides of the tube. The argon was then removed by pumping it off in the cold. On heating the splashed-off platinum, no helium spectrum could be observed. In a similar experiment, in which a very minute trace of helium had been added to the argon, there was no difficulty in detecting the helium on heating the tube. The conclusion therefore follows that no helium is present in our atmosphere.
It is, besides, exceedingly improbable that helium should be present. For, as shown by Dr. Johnstone Stoney, the rate of motion of a molecule of hydrogen is so rapid that, on finding its way to the confines of the atmosphere, it would escape, and travel through space until it found aplanet of sufficient attractive force to hold it. The same is true of helium. Were helium present in the atmosphere, it would ultimately leave us for the sun, or for some planet of much greater mass than the earth. This conception of Dr. Stoney’s tallies with the observation that the moon, a planet of small mass, is devoid of an atmosphere, and that the sun, a body with a mass 300,000 times as great as the earth, shows in its chromosphere the spectra both of hydrogen and of helium with great brilliancy.
It is now of interest to inquire what are the properties of argon and how it is related to other elements.
THE PROPERTIES OF ARGON
The density of a gas is one of its most characteristic and important properties. Avogadro’s law, which postulates that equal volumes of gases, at equal temperature and pressure, contain equal numbers of molecules, renders it possible to compare the weights of the molecules by determining the relative weights of the gases. Thus, as the ratio between the densities of nitrogen and oxygen is 7 to 8, a single molecule of nitrogen, the smallest portion which can exist in freedom, uncombined with other elements, is ⅞ths of the weight of a single molecule of oxygen. Hence a determination of the density of argon leads directly to a knowledge of the relative weight of a single molecule of this gas.
But with what should the density of argon be compared? What gas must serve as the standard of density? To answer this question, it is necessary to give a short sketch of the development of chemical theory regarding the atomic weights of elements and their relative volumes.
Dalton proposed to adopt as the unit of atomic weight the weight of the lightest atom, namely, that of hydrogen. Taking, for example, water as one substance containing hydrogen, its percentage composition by weight is—
If the smallest portion of water capable of free existence contains one atom of hydrogen and one of oxygen, then placing the weight of an atom of hydrogen as unity, the weight of an atom of oxygen is eight times as great. And although we do not know the absolute weight of any single atom, we are justified in supposing that an atom of oxygen is eight times as heavy as an atom of hydrogen. But have we any right to make the assumption that a molecule of water contains one atom of each element? Dalton came to the conclusion that this supposition was a justifiable one; but there are strong reasons against it. We havealready seen that Cavendish discovered approximately, and that Gay-Lussac and Humboldt determined accurately, that when hydrogen and oxygen unite to form water, two volumes of the former combine with one of the latter. Now it appears improbable on the face of it that any given volume of hydrogen should contain only half as many particles as an equal volume of oxygen; and it is still more improbable, when we take into consideration (1) Boyle’s discovery that if the pressure on a gas be increased, the volume of the gas, whatever it may be, diminishes in like proportion; and (2) Gay-Lussac’s and Dalton’s discovery, that all gases, when equally raised in temperature, expand equally. It would be very remarkable if one gas, containing twice as many particles in unit volume as another, should show exactly similar behaviour towards pressure and temperature. Hence it appeared not unreasonable to suppose that the composition of water was expressed by one particle of oxygen in union with two particles of hydrogen. (The word “particle” is here used in the meaning of “small portion”; such particles may be molecules or they may be atoms.)
When steam is formed by the union of hydrogen with oxygen, it has a volume equal not to the sum of the volumes of the hydrogen and the oxygen, but to two-thirds of the sum, or equal to that of the hydrogen alone, or twice that of the oxygen. And as steam, like hydrogen and oxygen, follows Boyle’s and Gay-Lussac’s laws, it must be supposed that in the steam there are as many particles as in the hydrogen from which it was formed. But the particles of steam must necessarily be more complex than those of the hydrogen, inasmuch as the steam contains oxygen as well as hydrogen.
These difficulties may, however, be easily overcome by the following supposition, which was first formulated by Avogadro in 1811. The ordinary particles of hydrogen and of oxygen are complex, each containing at least two atoms, or smaller particles, which usually exist in combination with each other, or with atoms of some other element. Two volumes of hydrogen, therefore, contain twice as many particles as one volume of oxygen; to such particles the name “molecules” is now universally applied. And as these molecules are themselves each made up of two smaller particles, now termed “atoms,” there exist in two volumes of hydrogen twice as many atoms as in onevolume of oxygen. On combination, the atoms in the molecules of hydrogen and oxygen rearrange themselves, so that two atoms of hydrogen and one atom of oxygen combine to form a molecule of water-vapour, containing three atoms. The steam now contains as many molecules as did the hydrogen before combination; but whereas the molecules of hydrogen originally consisted of two atoms each, the molecules of steam contain three atoms. It is this which causes the contraction from three volumes to two when hydrogen and oxygen molecules exchange partners in forming water molecules.
Of course the difficulty would meet with an equally good explanation if it were supposed that the hydrogen molecules and the oxygen molecules each contained four atoms, or eight atoms; but there is no need to increase the complexity of the molecule, and the assumption that these molecules are “diatomic” completely serves the purpose. The composition of water is therefore believed to be two atoms of hydrogen in combination with one atom of oxygen; and when hydrogen and oxygen unite to form water, a transaction similar to an exchange of partners is supposed to occur; the atoms of hydrogen and oxygen are imagined toleave their partners of like kind, and to rearrange themselves so that groups of atoms, or molecules, each containing two atoms of hydrogen and one of oxygen, are formed. To such an arrangement the formula H2O is applied, while ordinary hydrogen molecules may be represented as H2, and molecules of oxygen as O2.
It has been shown already (p. 150) how Lord Rayleigh obtained the number 15·882 for the density of oxygen compared with that of hydrogen. To determine the atomic weights of elements, the usual process has been to analyse their oxides, for only a few elements form compounds with hydrogen. Thus the analysis of copper oxide yields the numbers—
And as no compound of copper and hydrogen is known which lends itself to analysis, the atomic weight of copper is necessarily referred to that of oxygen. If the atomic weight of hydrogen be taken as unity, that of oxygen, from Lord Rayleigh’s determination, must be 15·882, because, in comparing the weights of equal volumes of the gases, a comparison is made of the weights of equal numbers of molecules; and asit is reasonable to suppose that each molecule of hydrogen and of oxygen contains two atoms, the number 15·882 represents the weight of an atom of oxygen compared with that of an atom of hydrogen taken as 1. But this number has not been regarded as sufficiently established by experiment. Other observers (for the importance of this ratio has been acknowledged since the beginning of the century) have obtained results differing from that given above, although not to any great extent. And as it is a matter of indifference what basis or standard be taken for atomic weights, which represent only relative numbers, it is common to accept the atomic weight of oxygen as 16, in which case that of hydrogen, if Lord Rayleigh’s determination of its density be regarded as accurate, would be 1·0074. Hence if we place the atomic weight of oxygen as 16, that of copper would be 63·34. And as with copper, so with most other elements. It is very seldom that the atomic weight of an element has been directly compared with that of hydrogen; it is, in fact, almost always ascertained by analysis of its chloride, bromide, or oxide; and the atomic weights of chlorine and bromine have been verycarefully compared with that of oxygen. There is, besides, another convenience in accepting 16 as the atomic weight of oxygen: it is that many atomic weights are then represented by whole numbers instead of by fractions; thus, sulphur has the atomic weight 32, if oxygen be made 16, whereas, if it were 15·882, the atomic weight of sulphur would be 31·764, a number much more difficult to remember.
We see then that it is convenient to refer the density of argon to oxygen taken as 16. The density obtained by Professor Ramsay in February 1895, using a globe of small capacity (only 160 cubic centimetres), was 19·94; exactly the same result was given by Lord Rayleigh’s experiments in June 1895 on argon prepared by means of the electric discharge, with a balloon of much greater capacity, which held over two litres of gas. Now as a molecule of oxygen consists of two atoms, the weight of a molecule is twice the atomic weight, or 32; and as a given volume of argon must contain as many molecules as the same volume of oxygen, the weight of a molecule of argon must be twice 19·94, or 39·88.
But this gives no information regarding the relative weight of anatomof argon. To ascertain this important quantity two methods may be chosen. One is to make compounds of the element, and this will be first considered. Since an atom of an element is defined as the smallest amount which can exist in combination, then, if numerous compounds of an element be examined, that one which contains proportionally the least amount of the element may be regarded as containing an atom, unless there are reasons to the contrary. For example, reverting to the former instance of water, the relative proportions by weight of oxygen and hydrogen are, in round numbers, 16 to 2. Reasons have already been given showing why its formula should be H2O and not HO; its molecule must contain two atoms of hydrogen. But another compound of oxygen and hydrogen is known in which the proportions are 16 parts by weight of oxygen to 1 part by weight of hydrogen. Here also there are reasons for believing that this compound, hydrogen peroxide, contains two atoms of hydrogen; whence it follows that it must contain two atoms of oxygen, or 32 parts by weight to 2 parts by weight of hydrogen, and must therefore have the formulaH2O2. No other compound of oxygen and hydrogen is known; and it may be stated briefly that no compound of oxygen with any element whatever is known in which less than 16 parts by weight enters, compared, of course, with the atomic weight of the other element or elements in the compound. Hence 16 is accepted on this ground as the atomic weight of oxygen.
If now it were possible to prepare compounds of argon, similar reasoning might be applied to them, and that compound containing least argon would be regarded as indicating its atomic weight. Many attempts were therefore made to induce argon to enter into combination. And the consistent failure of these attempts led to the choice of the name “argon” or “idle” for the newly discovered element. The methods employed to prepare argon free from nitrogen, namely, by exposing the mixed gases to the action of oxygen in a discharge of electric sparks, and by passing them over red-hot magnesium, show that it cannot be induced to combine with one of the most electro-negative of elements, oxygen, and one of the most electro-positive, magnesium. It also refuses to combine with hydrogen or with chlorine when sparked with these gases; nor is it absorbed or altered in volume by passage througha red-hot tube along with the vapours of phosphorus, sulphur, tellurium, or sodium. Red-hot caustic soda, or a red-hot mixture of soda and lime, which attacks the exceedingly refractory metal platinum, was without action on argon. The combined influence of oxygen and an alkali in the shape of fused potassium nitrate or red-hot peroxide of sodium was also without effect. Gold would, however, have resisted such action, but it would have been attacked by the next agent tried, viz. persulphide of sodium and calcium. This mixture was exposed at a red-heat to a current of argon, again without result. Nascent chlorine, or chlorine at the moment of liberation, obtained from a mixture of nitric and hydrochloric acids, and from permanganate of potassium and hydrochloric acid, was without action. A mixture of argon with fluorine, the most active of all the elements, was exposed to a rain of electric sparks by M. Moissan, the distinguished chemist who first succeeded in preparing large quantities of fluorine in a pure state, without his observing any sign of chemical combination.
An attempt was also made to cause argon to combine with carbon by making an electric arc between two rods of carbon in an atmosphere ofargon. It was at first believed that combination had taken place, for expansion occurred, the final volume of gas being larger than the volume taken; but subsequent experiments have shown that the expansion was due to the formation of some oxide of carbon from the oxygen adhering to the carbon rods. On absorption of this oxide by the usual absorbent, a mixture of cuprous chloride and ammonia, the argon was recovered unchanged.
M. Berthelot, the celebrated French chemist, has stated that, on exposing argon mixed with benzene vapour to a rain of electric sparks, he has succeeded in causing argon to combine. Its volume certainly decreases, but whether this decrease is to be attributed to true combination or not is very doubtful. The benzene is converted into a resinous mass, which coats the walls of the tube; and it is not improbable that the argon may be dissolved, or even mechanically retained, in the resinous deposit. Helium, a gas closely resembling argon in properties, may be made to enter into a similar combination with metallic platinum, if combination it can be called; but the amount absorbed in both cases is extremely minute, and the gas is evolved unchanged on heating the resin or the metal.
Professor Ramsay has also made experiments on the action of a silent electric discharge upon a mixture of argon with the vapour of carbon tetrachloride; the latter decomposes, giving, not a resin, but crystals of hexachlorobenzene and free chlorine; but the volume of the argon was unchanged. It was all recovered without loss. Lastly, the rare elements titanium and uranium have been heated to redness in a current of argon with no alteration or absorption of the gas. In short, all likely agents have been tried as absorbents for argon, but in no case has any true chemical combination been obtained.
These failures to produce compounds make it impossible to gain any knowledge regarding the atomic weight of argon by a study of its compounds, for it forms none. It is, indeed, in the highest degree improbable that, had compounds existed, none should have been found in Nature. There are, it is true, a few elements, such as platinum and those resembling it, which always occur native,i.e.in the elementary state; but even they yield to the attack of the agents tried with argon. It cannot, of course, be stated with absolute certainty that no element can combine with argon; but it appears at least improbable that any compounds will be formed.
It was therefore necessary to adopt some other method in attempting to determine the atomic weight of argon,—some method dependent on its physical rather than its chemical properties, for argon, unlike almost all other elements, appears to be devoid of chemical properties.
In order better to follow the train of reasoning based on experiment, it will be well to begin with an account of why the atomic weight of mercury is accepted as 200. The amount of mercury which combines with 16 parts by weight of oxygen is easily found by heating a weighed quantity of oxide of mercury, as Priestley and Scheele did, and weighing the residue of metal. The results of the most accurate experiments show that 200·36 grams of mercury combine with 16 grams of oxygen, and if the compound consists of one atom of each element, 200·36 must be the atomic weight of mercury. The first idea which naturally occurs is to find out the relative weight of mercury gas. This has been done, and it is found to have the ratio to that of oxygen of 100 to 16. Doubling these numbers will give the molecular weights, since a molecule of oxygen consists of two atoms, and must thereforepossess twice the weight of one atom. We thus obtain the numbers 200 and 32 as the molecular weights of mercury and oxygen respectively. It might therefore be concluded that 200 is not the true atomic weight of mercury, but 100, and that the compound of mercury with oxygen contains not one but two atoms of mercury, and should therefore be represented by the formula Hg2O, not HgO. But on surveying all known compounds of mercury, there is not one which contains less than 200 parts by weight of mercury in a molecule of the compound, or in which the mercury cannot be conceived to replace 2 parts by weight of hydrogen. And on weighing as gases the compounds of mercury with other elements, where such compounds do not decompose on heating like the oxide, the amount of mercury present must be always taken as 200, in order to add up to the molecular weight found. For example, a compound of mercury with carbon and hydrogen, named mercury methide, has a density of 120 compared with oxygen taken as 16, hence the comparative weight of its molecule must be 240. Now it is known to contain two atoms of carbon and six atoms of hydrogen, the atomic weights of which are 24 + 6 = 30.And deducting 30 from 240, 210 remains as an approximation to the atomic weight of mercury. It might, it is true, be the weight of two atoms of mercury, but if so it is singular that no compound contains a smaller proportion; and there is another reason, which follows immediately, that leads us to believe that 200 is correctly taken as the true weight of an atom.
It was discovered by Dulong and Petit, early in the century, that the higher the atomic weight of an element the less heat is required to raise its temperature through a given number of degrees. This heat can be measured by dropping a fragment of the element, carefully weighed and heated to a known temperature, into a known weight of cold water, and ascertaining what rise of temperature the water undergoes, owing to the heat communicated to it by the element. These comparative amounts of heat, if water is chosen as the standard, are termed specific heats. And as the specific heats of elements have been found by experiment to be inversely as their atomic weights, the product of the specific heat of any element and its atomic weight will give a constant number. If the quantity of element weighed is one gram, and its rise of temperature one degree, the numerical value of this product is about 6·4.
Now the specific heat of mercury has been found to equal 0·032; that is to say, it requires only a fraction of the value of 0·032 of heat to raise the temperature of say 1 gram of mercury through one degree, whereas the amount of heat necessary to raise 1 gram of water through one degree is represented by the number 1. Hence this number, 0·032, multiplied by the atomic weight of mercury, should yield the product 6·4; and it is seen at once that that number must be 200, for 200 × 0·032 = 6·4. This is an additional reason for believing that the atomic weight of mercury must be represented by the number 200.
We come next to a confirmatory piece of evidence which greatly strengthens the view that the atomic weight of mercury must be 200; but before entering into detail let us see what an atomic weight of 200 involves. The density of mercury gas is 100, and its molecular weight must be 200. But if its atomic weight is also 200, it follows of necessity that its molecule and its atom must be identical; that unlike oxygen and hydrogen, its molecule consists, not of two atoms, but of one single atom. There is nothing strange in this conclusion; there isno evident reason why single atoms should not act as molecules, or independent particles, able to exist in a free state, uncombined with each other or with any other molecules.
The specific heat of a gas is measured in much the same manner as that of a solid. A known volume of the gas is caused to pass through a spiral tube, heated to a certain definite high temperature; it then enters a vessel containing a known weight of water, traverses a spiral tube immersed in the water, and parts with its heat to the water. Knowing, therefore, the weight of the gas and its initial temperature, and also the rise of temperature of the water, the specific heat of the gas can be compared with that required to raise an equal weight of water through one degree. But gases are found to possess two specific heats. If the volume of the gas is kept constant, so that the gas does not contract during its loss of heat, one number for its specific heat is obtained; while if it is allowed to alter its volume a higher figure represents its specific heat. It will be necessary to consider the cause of this difference, in order to understand what conclusions can be drawn respecting the molecular nature of argon from a determination of the ratio between its two specific heats—that at constant pressure and that at constant volume.
If a gas is allowed to expand into a vertical cylinder so as to drive up a piston loaded with a weight, it is said to “do work.” The work is measured by the weight on the piston, and also by the height to which it is raised. Thus, if the weight is one pound, and the height one foot, one foot-pound of work is done; if the mass is one gram and the height one centimetre, one gram-centimetre of work is done. During this process the gas must expand; and if it were enclosed in some form of casing through which heat could not pass—we know of no such casing, but we can contrive casings through which heat passes very slowly—the temperature of the gas would fall during its expansion, and it would lose heat. For each loss of one heat-unit or calory—i.e.the amount of heat given off by 1 gram of water in cooling through 1° Centigrade—the gas would perform 42,380 gram-centimetres of work; it would raise a weight of nearly 4¼ kilograms, or about 9⅓ lbs., through 1 centimetre, or nearly half an inch.
When a gas expands into the atmosphere it may be regarded as “raising the atmosphere” through a certain height, for the atmosphere possesses weight, equal on the average to 1033 grams on each square centimetre of the earth’s surface, or between 15 and 16 lbs. on each square inch. Suppose a quantity of air, weighing 1 gram, to be enclosed in a long cylindrical tube of one square centimetre in section. At the usual pressure of the atmosphere on the earth’s surface, and at 0° Centigrade, the volume of the air would be 773·3 cubic centimetres; and, as the sectional area of the tube is 1 square centimetre, the air would occupy 773·3 centimetres’ length of the tube. If heat be given to this air, so that its temperature is raised from 0° to 1°, it will expand, as Gay-Lussac showed, by1⁄273rd of its volume. Now the product of 773·3 and1⁄273is 2·83 centimetres; the level of the surface of the air will rise in the tube through that amount. In doing so it will perform the work of raising 1033 grams through 2·83 centimetres, or 2927 gram-centimetres. Careful measurements have shown that, in order to do this work, heat to the amount of 0·0692 calory must be given to the gas. But it has been found that to heat the air through one degree, without allowing it to expand, requires 0·1683 calory; that is, thesame amount of heat which would raise a gram of air through one degree, its volume being kept constant, will raise a gram of water through 0·1683°; or, in other words, the specific heat of air is 0·1683. But if allowed to expand, more heat is required—an additional 0·0692 calory must be given it; consequently its specific heat at constant pressure is greater; it is actually the sum of these two numbers, 0·1683 + 0·0692 = 0·2375.
We have thus—
This ratio is termed the ratio between the specific heats of air, and such a ratio is represented usually by the letterγ.
But it is not necessary to determine both kinds of specific heat in order to arrive at a knowledge of the value of this ratio. One plan, adopted by Gay-Lussac and Désormes at the suggestion of Laplace,[28]is to actually measure the fall of temperature by allowing a known volumeof gas, of which the weight can of course be deduced, to expand from a pressure somewhat higher than that of the atmosphere to atmospheric pressure. It is true that heat will rapidly flow in through the walls of the vessel; but by choosing a sufficiently large vessel, and surrounding its walls with badly-conducting material, the entry of heat will be so slow that it may, for practical purposes, be neglected. The number for this ratio, actually found by Gay-Lussac and Welters for air, was 1·376; but subsequent and more accurate experiments have given as a result 1·405, which is almost identical with that calculated above.
This method, however, can be employed only when an unlimited supply of gas is at disposal, for it entails the use of large vessels, and the compressed gas must be allowed to escape into the atmosphere, and is lost. There is, fortunately, another method by which the same results can be obtained, and which requires only a small amount of gas.
Sir Isaac Newton calculated that the velocity of sound in a gas was dependent on its pressure and on its density, in such a manner that
wherecstands for velocity (celerity),pfor pressure, anddfor density. When waves of sound are transmitted through air, the air is compressed in parts and rarefied in parts, in such a manner that compression follows rarefaction very rapidly, that part which is compressed at one instant being rarefied at the next, compressed again at a third, and rarefied at a fourth, and so on. Laplace was the first to point out that during such rapid changes of pressure as occur while a sound-wave is passing, the pressure will not rise proportionally to the density, as would be the case if Boyle’s law were followed; for on sudden rise of pressure the temperature of the compressed portion of the gas will be increased; and, correspondingly, on sudden fall of pressure, the wave of compression having passed, the temperature will fall. He showed that instead of two pressures being inversely proportional to their two volumes, under such circumstances, as they are according to Boyle’s law, or
they must be inversely proportional to the volumes raised to a power, the numerical expression of which is the ratio of the specific heats of the two gases,γ, thus:
or as
v1:v::d:d1,
The ratio of the two specific heats can therefore be determined by finding the velocity of sound in the gas, and by noting at the same time its density and its pressure.
To determine the velocity of sound in a gas, it is not necessary to adopt the plan which has been successfully carried out with air; that is, to make a sudden sound at one spot and to measure the interval of time which the sound takes to travel to another spot some miles distant. There is a simpler method, depending on the fact that the lengths of the waves of compression and rarefaction are proportional to the velocity of the sound. So that, knowing the velocity of sound in air, the velocity in any other gas may be found by determining the relative length of the sound-waves in air and in that gas.