It was recognized by St. Claire Deville (1857) that the decrease in density of such mixtures of gases was due, not to their being exceptions to Avogadro’s law, but to the gradual decomposition of the compound body with rise of temperature. To this gradual decomposition he gave the name dissociation. This conception has proved of the utmost importance to the science, as will be seen in the sequel. To take the above instance of ammonium chloride, its abnormal density is due to its dissociation into ammonia and hydrogen chloride; and the gas which is obtained on raising its temperature consists, not of gaseous ammonium chloride, but of a mixture of ammonia and hydrogen chloride, which, as is easily seen, occupy, when separate, twice the volume that would be occupied by the gaseous compound. Of recent years it has been shown by Brereton Baker that, if perfectly free from moisture, ammonium chloride gasifies as such, and that its density in the state of vapor is, in fact, 26.75.
The molecular complexity of gases has thus gradually become comprehended, and the truth of Avogadro’s law has gained acceptance. And as a means of picturing the behavior of gaseous molecules, the “Kinetic Theory of Gases” has been devised by Joule, Clausius, Maxwell, Thomson (Lord Kelvin), and others. On the assumption that the pressure of a gas on the walls of the vessel which contains it is due to the continued impacts of its molecules, and that the temperature of a gas is represented by the product of the mass of the molecules, or the square of their velocity, it has been possible to offer a mechanical explanation of Boyle’s law, that at constant temperature the volume of a gas diminishes in proportion as the pressure increases; of Gay-Lussac’s law, that all gases expand equally for equal rise of temperature, provided pressure is kept constant; the condition being that equal volumes of gases contain equal numbers of molecules. A striking support is lent to this chain of reasoning by the facts discovered by Thomas Graham (1805–1869), professor at University College, London, and subsequently master of the Royal Mint. Graham discovered that the rate of diffusion of gases into each other is inversely as the square roots of their densities. For instance, the density of hydrogen being taken as unity, that of oxygen is sixteen times as great; if a vessel containing hydrogen be made to communicate with one containing oxygen, the hydrogen will pass into the oxygen and mix with it; and, conversely, the oxygen will pass into the hydrogen vessel. This is due to the intrinsic motion of the molecule of each gas. And Graham found, experimentally, that for each volume of oxygen which enters the hydrogen vessel four volumes of hydrogen will enter the oxygen vessel. Now, 4 = √16; and as these masses are relatively 1 and 16, and their temperatures are equal, the square of their velocities are respectively 1 and 16.
The question of the molecular complexity of gases being thus disposed of, it remains to be considered what are the relative complexity of liquid molecules. The answer is indicated by a study of the capillary phenomena of liquids, one method of measuring which is the height of their ascent in narrow or capillary tubes. We shall not enter here into detail as to the method and arguments necessary; suffice it to say that the Hungarian physicist Eötvös was the first to indicate the direction of research, and that Ramsay and Shields succeeded in proving that the complexity of the molecules of most liquids is not greater than that of the gases which they form on being vaporized; and also that certain liquids, e.g., water, the alcohols, and other liquids, are more or less “associated,” i.e., their molecules occur in couplices of two, three, four, or more, and as the temperature is raised the complexity of molecular structure diminishes.
As regards the molecular complexity of solids, nothing definite is known, and, moreover, there appears to be no method capable of revealing it.
While the researches of which a short account has now been given have led to knowledge regarding the nature of molecules, the structure of the molecule has excited interest since the early years of the century, and its investigation has led to important results. The fact of the decomposition of acidified water by an electric current, discovered by Nicholson and Carlisle, and of salts into “bases” and “acids” by Berzelius and Hisinger in 1803, led to the belief that a close connection exists between electric energy, or, as it was then termed, “electric force,” and the affinity which holds the constituents of chemical compounds in combination. In 1807 Davy propounded the theory that all compounds consist of two portions, one electro-positive and the other electro-negative. This idea was the result of experiments on the behavior of substances, such, for example, as copper and sulphur—if portions of these elements be insulated and then brought into contact they become oppositely electrified. The degree of electrification is intensified by rise of temperature until, when combination ensues, the electrification vanishes. Combination, therefore, according to Davy, is concurrent with the equalization of potentials. In 1812 Berzelius brought forward an electro-chemical theory which for the following twenty years was generally accepted. His primary assumption was that the atoms of elements, or, in certain cases, groups of atoms, are themselves electrified; that each atom, or group of atoms, possesses two poles, one positive, the other negative; that the electrification of one of these poles predominates over that of the other, so that the atom or group is itself, as a whole, electro-positive, or electro-negative; that combination ensued between such oppositely electrified bodies by the neutralization, partial or complete, of their electric charges; and, lastly, that the polarity of an element or group could be determined by noting whether the element or group separated at the positive or at the negative pole of the galvanic battery, or electrolysis. For Berzelius, oxygen was the most electro-negative and potassium the most electro-positive of the elements, the bridge between the “non-metals” and the “metals” being hydrogen, which, with nitrogen, forms a basic, or electro-positive, group, while with chlorine, etc., it forms electro-negative groups. The fact that an electric current splits compounds in solution into two portions led Berzelius to devise his “dualistic” system, which involved the assumption that all compounds consist of two portions, one electro-positive, the other electro-negative. Thus sulphate of magnesium and potassium was to be regarded as composed of electro-positive potassium sulphate in combination with electro-negative magnesium sulphate; the former in its turn consisted of electro-negative sulphur trioxide (SO3) in combination with electro-positive oxide of potassium (K2O); while each of these proximate constituents of potassium sulphate were themselves composed of the electro-negative oxygen in combination with electro-positive sulphur, or potassium. On contrasting sulphur with potassium, however, the former was considered more electro-negative than the latter; so that the group SO3 as a whole was electro-negative, while K2O was electro-positive. The symbols given above, which are still in universal use, were also devised by Berzelius for the purpose of illustrating and emphasizing his views. These views, however, met with little acceptance at the time in England.
Lavoisier’s idea, that oxygen was the necessary constituent of all acids, began about this time to lose ground. For Davy had proved the elementary nature of chlorine; and hydrochloric acid, one of the strongest, was thus seen to contain no oxygen, and Davy expressed the view, founded on his observation, that iodic “acid,” I2O5, was devoid of acid properties until dissolved in water, and that the essential constituent of all acids was hydrogen, not oxygen. The bearing of this theory on the dualistic theory is, that while, e.g., sulphuric acid was regarded by Berzelius as SO3, containing no hydrogen, and was supposed to be separated as such at the positive pole of a battery, Davy’s suggestion led to the opposite conclusion that the formula of sulphuric acid is H2SO4, and that by the current it is resolved into H2 and SO4. Faraday’s electrolytic law, that when a current is passed through electrolytes in solution the elements are liberated in quantities proportional to their equivalents, led to the abandonment of the dualistic theory. For when a current is passed in succession through acidified water, fused lead chloride, and a solution of potassium sulphate, the quantities of hydrogen and oxygen from the water, of lead and chlorine from the lead chloride, and the potassium of the sulphate are in accordance with Faraday’s law. But in addition to the potassium there is liberated at the same pole an equivalent of hydrogen. Now, if Berzelius’s theory be true, the products should be SO3 and K2O, but if the opposite view be correct, then K2 is liberated first and by its subsequent action on water it yields potash and its equivalent of hydrogen. This was pointed out first by Daniell, professor at King’s College, London, and it was regarded as a powerful argument against Berzelius’s system. In 1833, too, Graham investigated the phosphoric acids, and prepared the salts of three, to which he gave the names, ortho-, pyro-, and meta- phosphoric acids. To understand the bearing of this on the doctrine of dualism it must be remembered that P2O5, pentoxide of phosphorus, was at that date named phosphoric acid. When dissolved in water it reacts with bases, forming salts—the phosphates. But the quantity of water necessary was not then considered essential; Graham, however, showed that there exist three series of salts—one set derived from P2O5,3H2O, one from P2O5,2H2O, and a third from P2O5,H2O. His way of stating the fact was that water could play the part of a base; for example, the ordinary phosphate of commerce possessed, according to him, the formula P2O5,2Na2O,H2O, two-thirds of the “water of constitution” being replaced by oxide of sodium. Liebig, then professor at Giessen (1803–1873), founded on these and on similar observations of his own the doctrine of poly-basic acids—acids in which one, two, three, or more atoms of hydrogen were replaceable by metals. Thus, instead of writing, as Graham did, P2O5,2Na2O,H2O, he wrote, PO4Na2H; and for orthophosphoric acid PO4H3. The group of atoms (PO4), therefore, existed throughout the whole series of orthophosphates, and could exist in combination with hydrogen, with hydrogen and metals, or with metals alone. Similarly the group (P2O7) was characteristic of pyrophosphates and (PO3) of metaphosphates, for P2O5,2H2O=(P2O7)H4; and P2O5,H2O=2(PO3)H.
The first clear ideas of the structure of the molecule were, however, gained from the study of the compounds of carbon. It was difficult to apply the dualistic theory to them. For few of them are electrolytes, and therefore their products of electrolysis, being non-existent, could not be classified. Nevertheless, Gay-Lussac regarded alcohol, C2H6O, as a compound of C2H4, ethylene, and H2O, water; and oxalic acid (anhydrous), C2O3, as one of CO2 with CO. The discovery of “isomeric compounds,” i.e., of compounds which possess the same ultimate formula and yet differ entirely in their properties, forced upon chemists the necessity of attending to the structure of the molecule; for only by such a supposition could the difference between two isomeric bodies be explained. In 1823 Liebig discovered that silver fulminate and silver cyanate both possessed the empirical formula AgCNO; in 1825 this was followed by the discovery by Faraday that oil gas contains a hydrocarbon identical in composition with ethylene, C2H4, yet differing from it in properties; and in 1829 Wöhler, professor in Göttingen (1800–1882), discovered that urea, a constituent of urine, could be produced by heating ammonium cyanate, NH4CNO, a substance of the same formula. It therefore became clear that the identity of a compound must depend on some other cause than its ultimate composition.
In 1833 Liebig and Wöhler took an important step in elucidating this question by their investigations on benzoic acid and acid obtainable by distilling a resin named gum benzoin. They showed that this acid, C7H6O2, could be conceived as consisting of the group C7H5O, to which they gave the name “benzoyl,” in combination with OH; that benzoic aldehyde, C7H6O, might be regarded as its compound with hydrogen; that it also formed compounds with chlorine, and bromine, and sulphur, and replaced hydrogen in ammonia (C7H6O,NH2). They termed this group, benzoyl, a “compound element” or a “radical.” This research was followed by one by Robert Bunsen, professor at Heidelberg, born in 1811, and recently (1899) dead, which bore reference to cacodyl, a compound of arsenic, carbon and hydrogen, in which the idea of a radical was confirmed and amplified.
The idea of a radical having thus become established, Jean Baptiste Andrée Dumas, professor in Paris (1800–1884), propounded the theory of “substitution,” i.e., that an element such as chlorine or oxygen (which, be it noticed, is electro-negative on Berzelius’s scale) could replace hydrogen in carbon compounds, atom for atom, the resulting compound belonging to the same “type” as the one from which it was derived. And Laurent, warden of the mint at Paris (1807–1853), and Gerhardt, professor at Montpelier and at Strasburg (1816–1856), emphasized the fact that one element, be it what it may, can replace another without fundamentally altering its chemical character, and also that an atom of hydrogen can be replaced by a group of atoms or radical, behaving for the occasion like the atom of an element. It is to Laurent and Gerhardt that we owe the definition of an atom—the smallest quantity of an element which can be present in a compound; an equivalent—that weight of an element which combines with or replaces one part by weight of hydrogen; and a molecule—the smallest quantity which can exist in a free state, whether of an element or a compound. They recognized, too, that a molecule of hydrogen, chlorine, etc., consists of two atoms.
In 1849 Wurtz, professor in Paris (1817–1884), and Hofmann, then professor in the College of Chemistry in London, afterwards at Berlin (1818–1892), discovered a series of compounds allied to ammonia, NH3, in which one or more atoms of hydrogen were replaced by a group or radical, such as methyl (CH3), ethyl (C2H5), or phenyl (C6H5). Wurtz referred such compounds to the ammonia “type.” They all resemble ammonia in their physical properties—smell, taste, etc.—as well as in their power of uniting with acids to form salts resembling ammonium chloride (NH4Cl), and other ammonium compounds. Shortly afterwards Williamson, professor at University College, London, added the “water type,” in consequence of his researches on “mixed ethers”—bodies in which the hydrogen of water might be regarded as replaced by organic radicals. Thus we have the series: