Nov. 21, 1878] 



NATURE 



65 



The Italian chemist, Amadeo Avogadro, in discussing the 

 discoveries of Gay-Lussac respecting the simple relations which 

 exist between the volumes in which gases combine, was the first 

 to recognise that there likewise exists a simple relation between 

 the volumes of gases and the number of molecules which they 

 contain. The simplest hypothesis, said he, that can be made 

 regarding this matter consists in supposing that all gases contain 

 in equal volumes equal numbers of " integrant molecules. " By 

 this term he denoted what we now call simply molecules, and he 

 distinguished these integrant molecules from the "elementary 

 molecules " which we call atoms. According to him the inte- 

 grant molecules of gases are all equally distant one from the 

 other, and these distances are so great in proportion to the 

 dimensions of the molecules, that the mutual attraction between 

 the latter is reduced to nothing. 



These integrant molecules are composed of a greater or 

 smaller number of elementary molecules, not only m com- 

 pound, but Ukewise in simple bodies ; the integrant molecules 

 of chlorine, for example, are composed of four elementary 

 molecules, and the same is the case with the integrant mole- 

 cules of hydrogen. What happens, then, when chlorine and 

 hydrogen combine together ? The integrant molecules of these 

 two bodies are then resolved into elementary molecules, which 

 combine, two by two, to form hydrochloric acid. 



Ideas analogous to those of the Italian chemist were enun- 

 ciated in 1814 by Ampere, and thus there has been intro- 

 duced into chemical science the notion that there exist two 

 kinds of small particles, namely, molecules and atoms, 

 the former being diffused in equal numbers through equal 

 volumes of gases. 



But this notion, so clearly enunciated more than sixty years 

 ago, was afterwards destined to be obscured. Berzelius, taking 

 up Ampere's proposition, altered it by substituting atoms for 

 molecules, and saying that ' ' equal volumes of gases contain 

 equal numbers of atoms." This proposition, which has given 

 rise to long discussions, must now be rejected, for it is inexact. 

 It is to Gerhardt, and more recently to Cannizzaro, that is due 

 the honour of having restored the thesis of Avogadro and 

 Ampere, and pointed out its importance in connection with 

 chemical theory. This I must explain in conclusion. 



In the first place Gerhardt simplified the rule of Avo- 

 gadro. The latter supposed that a molecule of chlorine or 

 of hydrogen contains four atoms, whereas Gerhardt regards it as 

 consisting of two. Avogadro's proposition thus modified, 

 assumes a very simple form, and may be enunciated in the 

 following tenns. Suppose that a volume, or the unit of volume, 

 of hydrogen contains one atom ; then the molecules of all gases 

 and vapours will occupy two volumes. Thus, a molecule of 

 hydrogen formed of two atoms will occupy two volumes, and a 

 molecule'of chlorine formed of two atom> will likewise occupy 

 two volumes. What now will happen when chlorine combines 

 with hydrogen? The molecules will be cut in two, and each of 

 the two chlorine-atoms uniting itself to an atom of hydrogen, 

 two molecules of hydrochloric acid will be formed, each occu- 

 pying two volumes. Thus if an atom of hydrogen occupies one 

 volume, a molecule of hydrochloric acid will occupy two 

 volumes. The same is the case -v^ ith the molecules of all other 

 gases and vapours. 



A molecule of water formed of 2 at. H and i at. O occupies 2 volumes. 

 ,, ammonia ,, 3 at. H and i at. N „ „ 



,, marsh gas „ 4 at. H and i at. C „ „ 



This list might be prolonged by taking as examples a large 

 number of gaseous or volatile bodies belonging both to mineral 

 and to organic chemistry, and including chlorinated, brominated, 

 and oxygenated compounds of the metalloids, and of a large 

 number of metals. The countless volatile compounds of organic 

 chemistry, hydrocarbons, alcohols, chlorides, bromides, organo- 

 metallic compounds, compound ammonias, aldehydes, ketones — 

 all this legion of various compounds — conform to • the law of 

 Avogadro and Ampere, their molecules occupying two volumes 

 if an atom of hydrogen occupies one volume. Hence it follows 

 that the relative weights of two volumes represent the relative 

 weights of the molecules, or the molecular weigh ts. To 

 find these latter, therefore, it is sufficient to double the numbers 

 which express the weights of a single volume, or of the unit 

 of volume, that is to say the densities. The densities of^ases 

 may be referred to that of hydrogen as unity, and the atomic 

 weights to that of hydrogen. The unit being then the same, it 

 follows that the numbers which express the double densities 

 referred to hydrogen will also represent the molecular weights. 



Chemists represent the constitution of molecules by formulae, 

 each of which shows the number of atoms condensed within the 

 molecule. Now the molecular weights being known, it is very 

 easy to deduce the formulae from them, as the se formula; must 

 represent the number of atoms comprised in two volumes. Such 

 is the relation which exists between the Law of Volumes and 

 Chemical Notation. The rule of Avogadro and Ampere has, in 

 fact, become one of the bases of this notation. There are, how- 

 ever, certain exceptions to its generality, but they are probably 

 more apparent than real. Sal-ammoniac, ammonium sulph- 

 hydrate, phosphorus pentachloride, iodine trichloride, sulphuric 

 acid, calomel, amylene hydrobromide, and chloral hydrate, have 

 vapour-densities such that their molecules appear to occupy four 

 volumes. Such, however, is not the case ; and it may be shown 

 that the bodies in question do not volatilise without decompo- 

 sition, but that, when they are heated, their molecules split up 

 into two, each of which occupies two volumes. Being unable to 

 analyse all the cases above-mentioned, I will confine myself to 

 the last, viz., chloral hydrate, which has given rise to a 

 long discussion. 



The question to be decided is, whether this compound is or 

 is not decomposed by conversion into vapour? If it really 

 suffers decomposition, it should be resolved into anhydrous 

 chloral and water, lliat this decomposition really takes place 

 may be shown by a method based on the theory of dissociation 

 developed by M. H. Sainte-Claire Deville. 



Here is the case in a few words. We have here in a tube a 

 certain volume of the vapour of chloral hydrate under a certain 

 pressure ; it is required to show that this vapour contains vapovu: 

 of water. For this purpose we are about to introduce into it 

 a body capable of emitting vapour of water, crystallised potas- 

 sium oxalate, for example. If the atmosphere is dry, this salt 

 will give off vapour of water just as it would in dry air or in vapour 

 of chloroform at the same temperature, and it will continue to 

 emit this vapour until the atmosphere shall have taken up a 

 degree of humidity corresponding with that which is designated 

 by M. H. Sainte-Claire Deville the dissociation tension of the 

 hydrated salt in question. If, on the other hand, the chloral 

 atmosphere is moist, and exhibits exactly the degree of humidity 

 just defined, the crystallised oxalate will not emit any water. In 

 this first tube, then, we have the vapour of chloral hydrate ; the 

 second contains vapour of chloroform. This latter is dry, and 

 I am about to prove to you that the former is moist. In fact, 

 the crystallised potassium oxalate which we are introducing into 

 the chloroform tube will rapidly depress the level of the 

 mercury by emitting vapour of water, whereas in the atmosphere 

 of chloral hydrate it will not emit vapour of water, and con- 

 sequently will not depress the level of the mercury. This shows 

 that chloral hydrate undergoes decomposition when converted 

 into vapour, and this supposed exception to the rule of Avo- 

 gadro and Ampere vanishes, like all the rest, when submitted to 

 the test of experiment. This rule appears, then, like a grand 

 law of nature, as simple in its enunciation as it is important in 

 its consequences. 



Such are the considerations which I wished to lay before you 

 on the physical and chemical constitution of gases. Does not 

 this exposition seem to show that, of all the states which matter 

 can assume, the gaseous state is the most accessible to our 

 researches, and the best knoAvn — not, indeed, that we can affirm 

 the certainty of the theoretical considerations which I have 

 brought before you, for they are but probable. In the physical 

 sciences nothing is certain but well-observed facts and their im- 

 mediate consequences ; and, whenever we attempt to make these 

 facts the basis of any general theory, hypothetical data are apt 

 to mix themselves up with our deductions. In the present case 

 the hypothesis consists in assuming that gases, and matter in 

 general, are formed of molecules, and these latter of atoms. [No 

 one has ever seen these molecules and atoms, and it is certain 

 that nobody ever ^vill see them. Does it follow then that we 

 ought to reject or disdain this hypothesis ? By no means. Our 

 theories may be verified in their consequences, and may thereby 

 acquire a certain degree of probability. The theory under con- 

 sideration has been subjected to this ordeal, and nothing has 

 hitherto been found to contradict it. It is probable, indeed, that 

 gases are composed of small particles moving freely in space, 

 with immense velocities, and capable of communicating their 

 motion by collision or by friction. It is probable that these 

 molecules are diffused in space in numbers so enormous that the 

 most rarefied spaces still contain legions of them ; and it is this 

 circumstance which explains the possibility of the movements of 

 the radiometer. 



