Given in the Royal University of Genoa
Il Nuovo Cimento, vol. vii. (1858), pp. 321-366
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I BELIEVE that the progress of science made in these last years has confirmed the hypothesis of Avogadro, of Ampère, and of Dumas on the similar constitution of substances in the gaseous state; that is, that equal volumes of these substances, whether simple or compound, contain an equal number of molecules: not however an equal number of atoms, since the molecules of the different substances, or those of the same substance in its different states, may contain a different number of atoms, whether of the same or of diverse nature.
In order to lead my students to the conviction which I have reached myself, I wish to place them on the same path as that by which I have arrived at it -- the path, that is, of the historical examination of chemical theories.
I commence, then, in the first lecture by showing how, from the examination of the physical properties of gaseous bodies, and from the law of Gay-Lussac on the volume relations between components and compounds, there arose almost spontaneously the hypothesis alluded to above, which was first of all enunciated by Avogadro, and shortly afterwards by Ampère. Analysing the conception of these two physicists, I show that it contains nothing contradictory to known facts, provided that we distinguish, as they did, molecules from atoms; provided that we do not confuse the criteria by which the number and the weight of the former are compared, with the criteria which serve to deduce the weight of the latter; provided that, finally, we have not fixed in our minds the prejudice that whilst the molecules of compound substances may consist of different numbers of atoms, the molecules of the various simple substances must all contain either one atom, or at least an equal number of atoms.
In the second lecture I set myself the task of investigating the reasons why this hypothesis of Avogadro and Ampère was not immediately accepted by the majority of chemists. I therefore expound rapidly the work and the ideas of those who examined the relationships of the reacting quantities of substances without concerning themselves with the volumes which these substances occupy in the gaseous state; and I pause to explain the ideas of Berzelius, by the influence of which the hypothesis above cited appeared to chemists out of harmony with the facts.
I examine the order of the ideas of Berzelius, and show how on the one hand he developed and completed the dualistic theory of Lavoisier by his own electro-chemical hypothesis, and how on the other hand, influenced by the atomic theory of Dalton (which had been confirmed by the experiments of Wollaston), he applied this theory and took it into agreement with the dualistic electro-chemical theory, whilst at the same time he extended the laws of Richter and tried to harmonise them with the results of Proust. I bring out clearly the reason why he was led to assume that the atoms, whilst separate in simple bodies, should unite to form the atoms of a compound of the first order, and these in turn, uniting in simple proportions, should form composite atoms of the second order, and why (since he could not admit that when two substances give a single compound, a molecule of the one and a molecule of the other, instead of uniting to form a single molecule, should change into two molecules of the same nature) he could not accept the hypothesis of Avogadro and of Ampère, which in many cases leads to the conclusion just indicated.
I then show how Berzelius, being unable to escape from his own dualistic ideas, and yet wishing to explain the simple relations discovered by Gay-Lussac between the volumes of gaseous compounds and their gaseous components, was led to formulate a hypothesis very different from that of Avogadro and of Ampère, namely, that equal volumes of simple substances in the gaseous state contain the same number of atoms, which in combination unite intact; how, later, the vapour densities of many simple substances having been determined, he had to restrict this hypothesis by saying that only simple substances which are permanent gases obey this law; how, not believing that composite atoms even of the same order could be equidistant in the gaseous state under the same conditions, he was led to suppose that in the molecules of hydrochloric, hydriodic, and hydrobromic acids, and in those of water and sulphuretted hydrogen, there was contained the same quantity of hydrogen, although the different behaviour of these compounds confirmed the deductions from the hypothesis of Avogadro and of Ampère.
I conclude this lecture by showing that we have only to distinguish atoms from molecules in order to reconcile all the experimental results known to Berzelius, and have no need to assume any difference in constitution between permanent and coercible, or between simple and compound gases, in contradiction to the physical properties of all elastic fluids.
In the third lecture I pass in review the various researches of physicists on gaseous bodies, and show that all the new researches from Gay-Lussac to Clausius confirm the hypothesis of Avogadro and of Ampère that the distances between the molecules, so long as they remain in the gaseous state, do not depend on their nature, nor on their mass, nor on the number of atoms they contain, but only on their temperature and on the pressure to which they are subjected.
In the forth lecture I pass under review the chemical theories since Berzelius: I pause to examine how Dumas, inclining to the idea of Ampère, had habituated chemists who busied themselves with organic substances to apply this idea in determining the molecular weights of compounds; and what were the reasons which had stopped him half way in the application of this theory. I then expound, in continuation of this, two different methods -- the one due to Berzelius, the other to Ampère and Dumas -- which were used to determine formulae in inorganic and in organic chemistry respectively until Laurent and Gerhardt sought to bring both parts of the science into harmony. I explain clearly how the discoveries made by Gerhardt, Williamson, Hofmann, Wurtz, Berthelot, Frankland, and others, on the constitution of organic compounds confirm the hypothesis of Avogadro and Ampère, and how that part of Gerhardt's theory which corresponds best with facts and best explains their connection, is nothing but the extension of Ampère's theory, that is, its complete application, already begun by Dumas.
I draw attention, however, to the fact that Gerhardt did not always consistently follow the theory which had given him such fertile results; since he assumed that equal volumes of gaseous bodies contain the same number of molecules, only in the majority of cases, but not always.
I show how he was constrained by a prejudice, the reverse of that of Berzelius, frequently to distort the facts. Whilst Berzelius, on the one hand, did not admit that the molecules of simple substance could be divided in the act of combination, Gerhardt supposes that all the molecules of simple substances could be divided in the act of combination, Gerhardt supposes that all the molecules of simple substances are divisible in chemical action. This prejudice forces him to suppose that the molecule of mercury and all the metals consists of two atoms, like that of hydrogen, and therefore that the compounds of all the metals are of the same type as those of hydrogen. This error even persists in the minds of chemists, and has prevented them from discovering amongst the metals the existence of biatomic radicals perfectly analogous to those lately discovered by Wurtz in organic chemistry.
From the historical examination of chemical theories as well as from physical researches, I draw the conclusion that to bring harmony all the branches of chemistry we must have recourse to the complete application of the theory of Avogadro and Ampère in order to compare the weights and the numbers of the molecules; and I propose in the sequel to show that the conclusions drawn from it are invariably in accordance with all physical and chemical laws hitherto discovered.
I begin in the fifth lecture by applying the hypothesis of Avogadro and Ampère to determine the weights of molecules even before their composition in known.
On the basis of the hypothesis cited above, the weights of the molecules are proportional to the densities of vapours to express the weights of the molecules, it is expedient to refer them all to the density of a simple gas taken as unity, rather than to the weight of a mixture of two gases such as air.
Hydrogen being the lightest gas, we may take it as the unit to which we refer the densities of other gaseous bodies, which in such a case express the weights of the molecules compared to the weight of the molecule of hydrogen = 1.
Since I prefer to take as common unit for the weights of the molecules and for their fractions, the weight of a half and not of the whole molecule of hydrogen, I therefore refer the densities of the various gaseous bodies to that of hydrogen = 2. If the densities are referred to air = 1, it is sufficient to multiply by 14.438 to change them to those referred to that of hydrogen = 1; and by 28.87 to refer them to the density of hydrogen = 2.
I write the two series of number, expressing these weights in the following manner:
|Names of Substances||Densities of weights|
of one volume, the
volume of Hydrogen
i.e., weights of the
molecules referred to
the weight of a whole
molecule of Hydrogen
taken as unity.
to that of Hydrogen
= 2, i.e., weights of
referred to the weight
of half a molecule of
Hydrogen taken as
|Sulphur below 1000°||96||192|
|Sulphur* above 1000°||32||64|
* This determination was made by Bineau, but I believe it requires confirmation.
** The numbers expressing the densities are approximate: we arrive at a closer approximation by comparing them with those derived from chemical dta, and bringing the two into harmony.
Whoever wishes to refer the densities to hydrogen = 1 and the weights of the molecules to the weight of half a molecule of hydrogen, can say that the weights of the molecule are all represented by the weight of two volumes.
I myself, however, for simplicity of exposition, prefer to refer the densities to that of hydrogen = 2, and so the weights of the molecules are all represented by the weight of one volume.
From the few examples contained in the table, I show that the same substance in its different allotropic states can have different molecular weights, without concealing the fact that the experimental data on which this conclusion is founded still require confirmation.
I assume that the study of the various compounds has been begun by determining the weights of the molecules, i.e., their densities in the gaseous state, without enquiring if they are simple or compound.
I then come to the examination of the composition of these molecules. If the substance is undecomposable, we are forced to admit that its molecule is entirely made up by the weight of one and the same kind of matter. If the body is composite, its elementary analysis is made, and thus we discover the constant relations between the weights of its components: then the weight of the molecule s divided into parts proportional to the numbers expressing the relative weights of the components, and thus we obtain the quantities of these components contained in the molecule of the compound referred to the same unit as that to which we refer the weights of all the molecules. By this method I have constructed the following table:-
|Name of Substance||Weight of one|
i.e., weight of
referred to the
weight of half
a molecule of
Hydrogen = 1
|Component weights of one volume,|
i.e., components weights of the
molecule, all referred to the weight
of half a molecule of Hydrogen
|Oxygen, ordinary||32||32 Oxygen|
|" electrised||128||128 "|
|Sulphur below 1000°||192||192 Sulphur|
|Sulphur above 1000° (?)||64||64 "|
|Hydrochloric Acid||36.5||35.5 Chlorine 1 Hydrogen|
|Hydrobromic Acid||81||80 Bromine 1 Hydrogen|
|Hydriodic Acid||128||127 Iodine 1 Hydrogen|
|Water||18||16 Oxygen 2 Hydrogen|
|Ammonia||17||14 Nitrogen 3 Hydrogen|
|Arseniuretted Hyd.||78||75 Arsenic 3 Hydrogen|
|Phosphuretted Hyd.||35||32 Phosphorus 3 Hydrogen|
|Calomel||235.6||35.5 Chlorine 200 Mercury|
|Corrosive Sublimate||271||71  " 200  "|
|Arsenic Trichloride||181.5||106.5  " 75 Arsenic|
|Protochloride of Phosphorus||138.5||106.5  " 32 Phosphorus|
|Perchloride of Iron||325||213  " 112 Iron|
|Protoxide of Nitrogen||44||16 Oxygen 28 Nitrogen|
|Binoxide of Nitrogen||30||16  " 14  "|
|Carbonic Acid||28||16  " 12 Carbon|
| " Acid||44||32  " 12  "|
|Ethylene||28||4 Hydrogen 24  "|
|Propylene||42||6  " 36  "|
|Acetic Acid, hydrated||60||4 "|
|" anhydrous||102||6 Hydrogen|
All the numbers contained in the preceding table are comparable amongst themselves, being referred to the same units. And to fix this well in the minds of pupils, I have recourse to a very simple artifice: I say to them, namely, "Suppose it to be shown that half molecule of hydrogen weighs a millionth of a milligram, then all the numbers of the preceding table become concrete numbers, expressing in millionth of a milligram the concrete weights of the molecules and their components: the same thing would follow if the common unit had any other concrete value,"and so I lead them to gain a clear conception of the comparability of these numbers, whatever be the concrete value of the common unit.
Once this artifice has served its purpose, I hasten to destroy it by explaining how it is not possible in reality to know the concrete value of this unit; but the clear ideas remain in the minds of my pupils whatever may be their degree of mathematical knowledge. I proceed pretty much as engineers do when they destroy the wooden scaffolding which has served them to construct their bridges, as soon as these can support themselves. But I fear that you will say, "Is it worth the trouble and the waste of time and ink to tell me of this very common artifice?" I am, however, constrained to tell you that I have paused to do so because I have become attached to this pedagogic expedient, having had such great success with it amongst my pupils, and thus I recommend it to all those who, like myself, must teach chemistry to youths not well accustomed to the comparison of quantities.
Once my students have become familiar with the importance of the numbers as they are exhibited the preceding table, it is easy to lead them to discover the law which results from their comparison. "Compare," I say to them, "the various quantities of the same element contained in the molecule of the free substance and in those of all its different compounds, and you will not be able to escape the following law: The different quantities of the same element contained in different molecules are all whole multiples of on and the same quantity, which, always being entire, as the right to be called an atom."
|One molecule||of free hydrogen||contains||2||of hydrogen||= 2 x 1|
|"||of hydrochloric acid||"||1||"||= 1 x 1|
|"||of hydrobromic acid||"||1||"||= 1 x 1|
|"||of hydriodic acid||"||1||"||= 1 x 1|
|"||of hydrocyanic acid||"||1||"||= 1 x 1|
|"||of water||"||2||"||= 2 x 1|
|"||of sulphuretted hy-|
|"||2||"||= 3 x 1|
|"||of formic acid||"||2||"||= 2 x 1|
|"||of ammonia||"||3||"||= 3 x 1|
|"||of gaseous phosphur-|
|"||3||"||= 3 x 1|
|"||of acetic acid||"||4||"||= 4 x 1|
|"||of ethylene||"||4||"||= 4 x 1|
|"||of alcohol||"||6||"||= 6 x 1|
|"||of ether||"||10||"||= 10 x 1|
Thus all the various weights of hydrogen contained in the different molecules are integral multiples of the weights contained in the molecule of hydrochloric acid, which justifies our having taken it as common unit of the weights of the atoms and of the molecules. The atom of hydrogen is contained twice in the molecule of free hydrogen.
In the same way it is shown that the various quantities of chlorine existing in different molecules are all whole multiples of the quantity contained in the molecule of hydrochloric acid, that is, of 35.5, and that the quantities of oxygen existing in the different molecules are all whole multiples of the quantity contained in the molecule of water, that is, of 16, which quantity is half of that contained in the molecule of free oxygen, and eighth part of that contained on the molecule of electrised oxygen (ozone).
|One molecule||of free oxygen||contains||32||of oxygen||= 2 x 16|
|"||of ozone||"||128||"||= 8 x 16|
|"||of water||"||16||"||= 1 x 16|
|"||of ether||"||16||"||= 1 x 16|
|"||of acetic acid||"||32||"||= 2 x 16|
|One molecule||of free chlorine||contains||71||of chlorine||= 2 x 35.5|
|"||of hydrochloric acid||"||35.5||of chlorine||= 1 x 35.5|
|"||of corrosive sublimate||"||71||of chlorine||= 2 x 35.5|
|"||of chloride of arsenic||"||106.5||of chlorine||= 3 x 35.5|
|"||of chloride of tin||"||142||of chlorine||= 4 x 35.5|
In a similar way may be found the smallest quantity of each element which enters as a whole into the molecules which contain it, and to which may be given with reason the name of atom. In order, then, to find the atomic weight of each element, it is necessary first of all to know the weights of all or of the greater part of the molecules in which it is contained and their compostion.
If it should appear to any one that this method of finding the weights of the molecules is too hypothetical, then let him compare the composition of equal volumes of substances in the gaseous state under the same conditions. He will not be able to escape the following law: The various quantities of the same element contained in equal volumes either of the free element or of its compounds are all whole multiples of one and the same quantity; that is, each element has a special numerical value by means of which and of integral coefficients the composition by weight of equal volumes of the different substances in which it is contained may be expressed. Now, since all chemical reactions take place between equal volumes, or integral multiples of them, it is possible to express all chemical reactions by means of the same numerical values and integral coefficients. The law enunciated in the form just indicated is a direct deduction from the facts: but who is not led to assume from this same law that the weights of equal volumes represent the molecular weights, although other proofs are wanting? I thus prefer to substitute in the expression of the law the word molecule instead of volume. This is advantageous for teaching, because, when the vapour densities cannot be determined, recourse is had to other means for deducing the weights of the molecules of compounds. The whole substance of my cures consists in this: to prove the exactness of these latter methods by showing that they lead to the same results as the vapour density when both kind of method can be adopted at the same time for determining molecular weights.
The law above enunciated, called by me the law of atoms, contains in itself that of multiple proportions and that of simple relations between the volumes; which I demonstrate amply in my lecture. After this I easily succeed in explaining how, expressing by symbols the different atomic weights of the various elements, it is possible to express by means of formulae the composition of their molecules and of those of their compounds, and I pause a little to make my pupils familiar with the passage from gaseous volume to molecule, the first directly expressing the fact and the second interpreting it. Above all, I study to implant in their minds thoroughly the difference between molecule and atom. It is possible indeed to know the atomic weight of an element without knowing its molecular weight; this is seen in the case of carbon. A great number of the compounds of this substance being volatile, the weights of the molecules and their composition may be compared, and it is seen that the quantities of carbon which they contain are all integral multiples of 12, which quantity is thus the atom of carbon and expressed by the symbol C ; but since we cannot determine the vapour density of free carbon we have no means of knowing the weight of its molecule, and thus we cannot know how many times the atom is contained it. Analogy does not in any way help us, because we observe that the molecules of the most closely analogous substances (such as sulphur and oxygen) , and even the molecules in the same substance in its allotropic states, are composed of different numbers of atoms. We have no means of prediction the vapour density of carbon: the only thing that we can say is that it will be either 12 or an integral multiple 12 (in my system of numbers). The number which is given in different treatises on chemistry as the theoretical density of carbon is quite arbitrary, and a useless datum in chemical calculations: it is useless for calculation and verifying the weights of the molecules of the various compounds of carbon, because the weight of the molecule of free carbon may be ignored it we know the weights of the molecules of all its compounds; it is useless for determining the weight of the atom of carbon, because this is deduced by comparing the composition of a certain number of molecules containing carbon, and the knowledge of the weight of the molecule of this last would scarcely add a datum more to those which are already sufficient for the solution of the problem. Any one will easily convince himself of this by placing in the following manner the numbers expressing the molecular weights derived from the densities and the weights of the components contained in them: --
|Names of Compounds|
|Weights of the components|
of the molecules referred to
the weight of the atom of
Hydrogen taken as unity
H = 1
C = 12
O = 16
S = 32
|Carbonic Oxide||28||12 Carbon 16 Oxygen||CO|
|" Acid||44||12 " 32 "||CO2|
|Sulphide of Carbon||76||12 " 64 Sulphur||CS2|
|Marsh Gas||16||12 " 4 Hydrogen||CH4|
|Ethylene||28||24 " 4 "||C2H4|
|Propylene||42||36 " 6 "||C3H6|
|Ether||74||48 " 10 "|
In the list of molecules containing carbon there might be placed also that of free carbon if the weight of it were known; but this would not have any greater utility than what we would derive by writing in the list one more compound of carbon; this is, it would do nothing but verify once more that the quantity of carbon contained in any molecules, whether, of the element itself of of its compounds, is 12 or n x 12 = Cn, n being an integral number.
I then discuss whether it is better to express the composition of the molecules of compounds as a function of the molecules of the components, or if, on the other hand, it is better, as I commenced by doing, to express the composition of both in terms of those constant quantities which always enter by whole numbers into both, that is, by means of the atoms. Thus for example, is it better to indicate in the formula that one molecule of hydrochloric acid contains the weight of half a molecule of hydrogen and half a molecule of of chlorine, or that it contains on atom of the other, pointing out at the same time that the molecules of both of these substances consist of two atoms?
Should we adopt the formulae made with symbols indicating the molecules of the elements, then many coefficients of these symbols would be fractional, and the formula of a compound would indicate directly the ratio of the volumes occupied by the components and by the compounds in the gaseous state. This was proposed by Dumas in his classical memoir, Sur quelques points de la Theorie atomique (Annales de Chimie et de physique, tom. 33,1826).
To discuss the question proposed, I give to the molecules of the elements symbols of a different kind from those employed to represent the atoms, and in this way I compare the formula made with the two kinds of symbols.
|Atoms or Molecules||Symbols of the|
the atoms of
|Atom of Hydrogen||H 1/2||= H =||1|
|Molecule of Hydrogen||H||= H2 =||2|
|Atom of Oxygen||O 1/2 = Oz 1/8||= O =||16|
|Molecule of ordinary Oxygen||O||= O2 =||32|
|Molecule of electrised oxygen|
|Oz||= O8 =||128|
|Atom of Sulphur||S 1/2 = Sa 1/6||= S =||32|
|Molecule of Sulphur above 1000°|
( Bineau )
|S||= S2 =||64|
|Molecule of Sulphur below 1000°||Sa||= S6 =||192|
|" Water||HO1/2 = HOz1/8||= H2O =||18|
|" Sulphuretted Hydrogen||HS 1/2 = HSa 1/6||= H2S =||34|
These few examples are sufficient to demonstrate the inconveniences associated with the formula indicating the composition of compound molecules as a function of the entire component molecules, which may be summed up as follows:--
1° It is not possible to determine the weight of the molecules of many elements the density of which in the gaseous state cannot be ascertained.
2° If it is true that oxygen and sulphur have different densities in their different allotropic states, that is, it they have different molecular weight, then their compounds would have two or more formulae according as the quantities of their components were referred to the molecules of one or the other allotropic state.
3° The molecules of analogous substances (such as sulphur and oxygen) being composed of different numbers of atoms, the formulae of analogous compounds would be dissimilar. If we indicate, instead, the composition of the molecules by means of the atoms, it is seen that analogous compounds contain in their molecules an equal number of atoms.
It is true that when we employ in the formulae the symbols expressing the weights of the molecules, i.e., of equal volumes, the relationship between the volumes of the components and those of the compounds follows directly ; but this relationship between the volumes of the components and those of the compounds follows directly ; but this relationship is also indicated in the formulae expressing the number of atoms ; it is sufficient to bear in mind that the atom represented by a symbol is either the entire molecule of the free substance or a fraction of it, that is, it is sufficient to know the atomic formula of the free molecule. Thus, to take an example, it is sufficient to know that the atom of oxygen, O, is one-half of the molecule of ordinary oxygen and an eighth part of the molecule of electrised oxygen-to know that the weight of the atom of oxygen is represented by 1/2 volume of free oxygen and 1/8 of electrised oxygen. In short, it is easy to accustom students to consider the weights of the atoms as being represented either by a whole volume or by a fraction of a volume, according as the atom is equal to the whole molecule or to a fraction of it. In this system of formulae, whose molecule or to a fraction of it. In this system of formulae, those which represent the weights and the composition of the molecules, whether of elements or of compounds, represent the weight and the composition of equal gaseous volumes under the same conditions. The atom of each element is represented by the quantity of it which constantly enters as a whole into equal volumes of the free substance or of its compounds ; it may be either the entire quantity contained in one volume of the free substance or a simple sub-multiple of this quantity.
This foundation of the atomic theory having been laid, I begin in the following lecture -- the sixth -- to examine the constitution of the molecules of the chlorides, bromides, and iodides. Since the greater part of these are volatile, and since we know their densities in the gaseous state, there cannot remain any doubt as to the approximate weights of the molecules, and so of the quantities of chlorine, bromine, and iodine contained in them. These quantities being always integral multiples of the weights of chlorine, always integral multiples of the weights of chlorine, bromine, and iodine contained in hydrochloric, hydrobromic, and hydriodic acids, i.e., of the weights of the half molecules, there can remain no doubt as to the atomic weights of these substance, and thus as to the number of atoms existing in the molecules of their compounds, whose weights and composition are known.
A difficulty sometimes appears in deciding whether the quantity of the other element combined with one atom of these halogens is 1, 2, 3, or n atoms in the molecule ; to decide this, it is necessary to compare the composition of all the other molecules containing the same element and find out the weight as this element which constantly enters as a whole. When we cannot determine, it is necessary then to have recourse to other criteria to know the weights of their molecules and to deduce the weight of the atom of the element. What I am to expound in the sequel serves to teach my pupils the method of employing these other criteria to verify or to determine atomic weights and the composition of molecules. I begin by making them study the following tables of some chlorides, bromides, and iodides whose vapour densities are know ; I write their formulae, certain of justifying later the value assigned to the atomic weights of some elements existing in the compounds indicated. I do not omit to draw their attention once more to the atomic weights of hydrogen, chlorine, bromine, and iodine being all equal to the weights of half a molecule, and represented by the weight of half a volume, which I indicate in the following table:--
|Weight of the atom of Hydrogen or half a mole-|
cule represented by the weight of 1/2 volume
|Weight of the atom of Chlorine or half a mole-|
cule represented by the weight of 1/2 volume
|Weight of the atom of Bromine or half a mole-|
cule represented by the weight of 1/2 volume
|Weight of the atom of Iodine or half a mole-|
cule represented by the weight of 1/2 volume
These data being given, there follows the table of some compounds of the halogens :--
|Names of the Chlorides||Weight of equal|
volumes in the gaseous
state, under the same
conditions, referred to
the weight of 1/2 volume
of Hydrogen = 1:
i.e., weights of the
molecules referred to
the weight of the atom
of Hydrogen = 1
|Composition of equal volume in the gaseous state,|
under the same conditions, i.e., composition of the
molecules, the weights of the components being all
refered to the weight of the atom of Hydrogen taken
as unity, i.e., the common unit adopted for the weights
of atoms and of molecules.
or of equal
volumes in the
under the same
|Free Chlorine||71||71 of chlorine||Cl2|
|Hydrochloric acid||36.5||35.5 " 1 of Hydrogen||HCl|
|Protochloride of Mercury or Calomel||235.5||35.5 " 200 of Mercury||HgCl|
|Bichloride of Mercury or Corrosive Sublimate||271||71 " 200 "||HgCl2|
|Chloride of Ethyl||64.5||35.5 " 5 of Hydrogen 24 of Carbon||C2H5Cl|
|" Acetyl||78.5||35.5 " 3 " |
24 of Carbon 16 of Oxygen
|" Ethylene||99||71 of Chlorine 4 of Hydrogen 24 of Carbon||C2H4Cl2|
|" Arsenic||181.5||106.5 " 75 of Arsenic||AsCl3|
|Protochloride of Phosphorus||138.5||106.5 " 32 of Phosphorus||PCl3|
|Chloride of Boron||117.5||106.5 " 11 of Boron||BCl3|
|Bichloride of Tin||256.6||142 " 117.6 of Tin||SnCl4|
|" Titainum||198||142 " 56 of Titainum||TiCl4|
|Chloride of Silicon||170||142 " 28 of Silicon||SiCl4|
|" Zirconium||231||142 " 89 of Zirconium||ZrCl4|
|" Aluminium||267||213 " 54 of Aluminium||Al2Cl6|
|Perchloride of Iron||325||213 " 112 of Iron||Fe2Cl6|
|Sesquichloride of Chromium||319||213 " 106 of Chromium||Cr2Cl6|
I stop to examine the composition of the molecules of the two chlorides and the two iodides of mercury. There can remain no doubt that the protochloride contains in its molecules the same quantity of chlorine as hydrochloric acid, that the bichloride contains twice as much, and that the quantity of mercury contained in the molecules of both is the same. The supposition made by some chemists that the quantities of chlorine contained in the two molecules are equal, and on the other hand that the quantities of mercury are different, is supported by no valid reason. The vapour densities of the two chlorides having been determined, and it having been observed that equal mercury, and that the quantity of chlorine contained in one volume of the vapour of calomel is equal to acid gas under the same conditions, whilst the quantity of chlorine contained in one volume of corrosive sublimate is twice that contained in an equal volume of calomel or of hydrochloric acid gas, the relative molecular composition of the two chlorides cannot be doubtful. The same may be said of the two iodides. Does the constant quantity of mercury existing in the molecules of these compounds, and represented by the number 200, correspond to one or more atoms? The observation that in these compounds the same quantity of mercury is combined with one or two atoms of chlorine or of iodine, would itself incline us to believe that this quantity is that which enters always as a whole into all the molecules containing mercury, namely, the atom; whence Hg = 200.
To verify this, it would be necessary to compare the various quantities of mercury contained in all the molecules of its compounds whose weights and composition are known with certainty. Few other compounds of mercury besides those indicated above lend themselves to this; still there are some in organic chemistry the formulae of which express well the molecular composition; in these formulae we always find Hg2 = 200, chemists having made Hg = 100 and H = 1. This is a confirmation that the atom of mercury is 200 and not 100, no compound of mercury existing whose molecule contains less than this quantity of it. For verification I refer to the law of the specific heats of elements and of compounds.
I call the quantity of heat consumed by the atoms or the molecules the product of their weights into their specific heats. I compare the heat consumed by the atom of mercury with that consumed by the atom of mercury with that consumed by the atoms of iodine and of bromine in the same physical state, and find them almost equal, which confirms the accuracy of the relation between the atomic weight of mercury and that of each of the two halogens, and thus also, indirectly, between the atomic weight of mercury and that of hydrogen, whose specific heats cannot be directly compared.
Thus we have --
| Name of|
| Specific heat,|
i.e., heat required
to heat unit
| Products of specific|
heats by atomic
weights, i.e., heat
required to heat
the atom 1°
The same thing is shown by comparing the specific heats of the different compounds of mercury. Woestyn and Garnier have shown that the state of combination does not notably change the calorific capacity of the atoms; and since this is almost equal in the various elements, the molecules would require, to heat them 1° quantities of heat proportional to the number of atoms which they contain. If Hg = 200, that is, if the formulae of the two chlorides and iodides of mercury are HgCl, HgI, HgCl2, HgI2, it will be necessary that the molecules of the first pair should consume twice as much heat as each separate atom, and those of the second pair three times as much; and this is so in fact, as may be seen in the following table:-
= p x c
= p x c / n