Physics and Chemistry are twin sisters—daughters of Natural Philosophy; like Juno’s swans, coupled and inseparable. Physics is concerned with the forms of energy which affect matter; chemistry with the study of matter so affected. Each, then, is complementary to the other. Philosophers of old drew no practical distinction between them, at least as regards their own studies. Men like Boyle, Black, Cavendish, Lavoisier, Dalton, Faraday, Graham, Bunsen, were pioneers “on a very broad gauge,” pushing their inquiries into territories common to the two branches as their genius or inclinations directed them. Accordingly, it has happened that many so-called physical laws have been discovered by men who were professed chemists. It has also happened that men who began their scientific career as chemists, like Dalton, Regnault, and Magnus, eventually gave the whole of their energies to physical measurements; or, like Black, Faraday, and Graham, devoted themselves to the elucidation of physical problems. As certain of these physical laws and problems have greatly influenced the progress of chemistry, it becomes necessary, in any historical treatment of the subject, to give some account of their origin, and to show how they affected the development of chemical theory.

The relations of heat to chemical phenomena are so obvious and so intimate that the study of their connection necessarily attracted attention in very early times. But it was only when this study became quantitative that any important generalisations became possible. Most quantitative estimations of heat depend eventually upon the thermometer; and thermometry is indebted to Englishmen in the first instance for attempts to render the instrument trustworthy.

In this connection may be mentioned the names of Newton and Shuckburgh. Brooke Taylor, in 1723, made a special study of the mercurial thermometer as a measurer of temperature. In other words, he sought to discover whether equal differences of expansion or contraction of mercury corresponded to equal additions or abstractions of heat. The results showed that the principle of the mercurial thermometer is valid within at least the limits of temperature between the boiling and freezing-points of water. These experiments were subsequently repeated and confirmed by Cavendish, and, independently, by Black.

The discovery of the phenomenon of latent heat by Black some time prior to 1760 marks an epoch in the history of science. It was then for the first time clearly recognised that the state of aggregation of a substance is associated with a definite thermal quantity, and that, in order to effect a change, a definite amount of energy, in the form of heat, must be employed. The quantitative connection that exists between work and energy was thus foreshadowed.

The doctrine of specific heat was taught by Black in his lectures at Glasgow between 1761 and 1765. The subject was subsequently investigated experimentally by Irvine between 1765 and 1770, and by Crawford in 1779. A series of determinations was published in 1781 by Wilcke, in the Transactions of the Swedish Academy. In these the term specific caloric, since changed to specific heat, was first used. About this time the determination of the amount of heat required to raise substances through a definite interval of temperature was made the subject of experiment by many observers, notably by Lavoisier and Laplace, who greatly improved the calorimetric arrangements. The values they obtained long remained the most trustworthy estimations of the specific heats of substances. Their joint research had a further influence on the development of thermo-chemistry by indicating the general experimental conditions which were needed to ensure accuracy in such determinations. Lavoisier and Laplace also measured, in 1782–1783, the heat disengaged by the combustion of substances, and that evolved during respiration. In 1819 Dulong and Petit pointed out that the specific heat of a number of substances, more particularly the metals, were inversely proportional to their atomic weights; or, in other words, the product of the specific heat into the atomic weight was a constant. The nature of the relation will be seen from the following table of certain of the results obtained by Dulong and Petit:—

It will be seen that these various elements have an uniform, or nearly uniform, atomic heat—approximately 6.2 on the average.

This would appear to prove that, as Dulong and Petit expressed it, “the atoms of simple substances have equal capacities for heat.” The variations from a constant value are due partly to errors of observation, but more particularly to the circumstance that the substances compared are not all in a strictly comparable condition—e.g., they are not all equally remote from their melting points. It was shown, moreover, that the amount of heat needed to raise a substance through a definite interval of temperature increased with the temperature. The range of temperature through which a determination was made in a particular instance affected, therefore, the value of the specific heat. The most noteworthy departures from a uniform value were observed to occur among the metalloids—e.g., carbon, the various modifications of which had different specific heats—and generally among elements of low atomic weight, in which the variation of specific heat with temperature was particularly rapid.

Nevertheless, the significance of the generalisation discovered by Dulong and Petit, in spite of its limitations, was quickly appreciated, as it was perceived that a knowledge of the specific heat of an element might be of great value in determining its atomic weight. The immediate effect was that a certain number of the atomic weights fixed by Berzelius mainly on chemical considerations were required to be halved. Although subsequent experience has proved that the law of Dulong and Petit is not capable of the simple mathematical expression they gave it, it has shown itself to be of great value in fixing doubtful atomic weights.

Pierre Louis Dulong was born in 1785 at Rouen, and, after studying chemistry and physics at the Polytechnic School at Paris, became its Professor of Chemistry and subsequently its Professor of Physics. In 1830 he was made its Director of Studies; and in 1832 he became permanent Secretary of the Academy of Sciences. As a young man he worked with Berzelius, with whom he made the first approximately accurate determination of the gravimetric composition of water. In 1811 he discovered the highly explosive nitrogen chloride, in the investigation of which he was severely injured, losing an eye and several fingers. He died in 1838. His collaborator, Alexis Therese Petit, was born in 1791 at Vesoul, and died, when holding the position of Professor of Physics at the Lycée Bonaparte, in 1820.

The attempt made by Neumann to extend Dulong and Petit’s “law” to compound substances was only partially successful. Nor has any important generalisation followed from our knowledge of the specific heat of liquids. Almost simultaneously with the publication of Dulong and Petit’s “law,” Mitscherlich made known the fact that similarity in chemical constitution is frequently accompanied by identity of crystalline form. Boyle, as far back as the middle of the seventeenth century, had insisted upon the importance of the forms of crystals in throwing light upon the internal structure of bodies. Romé de l’Isle and Hauy had remarked that many different substances had the same crystalline form. It had been observed that a crystal of potash alum would continue to grow and preserve its shape in a solution of ammonia alum; and similar observations had been shown to occur in the case of vitriols. The invention of the reflecting goniometer by Wollaston greatly facilitated the investigation of such phenomena. Mitscherlich showed that the phosphates and arseniates of analogous composition had the same crystalline shape, or, in other words, were isomorphous. The same fact was observed to occur in the case of the analogously constituted sulphates and selenates, and in that of the oxides of magnesium and zinc, etc. The value of isomorphous relations in determining the group-relationships of the elements and in deducing the composition of salts was at once recognised by Berzelius, who styled the discovery of isomorphism by his pupil Mitscherlich as “the most important since the establishment of the doctrine of chemical proportions.” The quantities of the isomorphously replacing elements in a compound were regarded by him as a measure of their atomic weights; and the principle was subsequently constantly employed by him, whenever possible, as a criterion in fixing their values. Other investigators have followed his example in this respect; and isomorphism is still regarded as an important consideration in establishing the genetic relations of an element.

Eilhard Mitscherlich, the son of a minister, was born in 1794 at Neu Ende, near Jever, in Oldenburg, and, after studying philology and oriental languages at Heidelberg, went to Paris, and thence to Göttingen, where he occupied himself with natural science. In 1818 he repaired to Berlin and commenced to work on the arseniates and phosphates, the similarity in the crystal-forms of which he was the first to detect. His friend Gustav Rose, the mineralogist, thereupon instructed him in the methods of crystallography; to enable him to verify his discovery and to establish it by goniometric measurements. In 1821 he joined Berzelius at Stockholm, where he pursued his inquiries on the connection between crystal-form and chemical composition. It was at the suggestion of Berzelius that he adopted the term “isomorphy” to express this connection—the mechanical consequence of identity of atomic constitution. In the same year he was appointed Klaproth’s successor in Berlin, where he died in 1863.

Mitscherlich also worked on the manganates and permanganates, on selenic acid, on benzene and its derivatives, and on the artificial production of minerals.

The study of the physical phenomena of gases, initiated in 1660 by Boyle’s discovery of the law of gaseous pressure, has greatly contributed to our knowledge of their intrinsic nature. Boyle himself only proved his law in the case of atmospheric air; but the observation was subsequently (1676) generalised by Marriotte. Charles, Dalton, and Gay Lussac independently showed that gases have the same rate of thermal expansion.

That gases are made up of particles possessing an internal movement was surmised by the Greeks; but experimental evidence for such a view of their constitution was first presented by Thomas Graham in 1829–1831, when he discovered that gases move, or are diffused, at rates inversely proportional to the square roots of their densities. Observations of a like character, which found their explanation in Graham’s discovery, had previously been made by Priestley, Döbereiner, and Saussure. This interchange in the position of their particles is a property inherent in gases. Inequality of density is not essential to diffusion. Graham proved this by connecting together two vessels, one containing nitrogen and the other carbonic oxide, which have the same density. After the expiration of a certain time both gases were found to be uniformly diffused through the vessels.

How these laws were found to be interdependent and mutually connected, and how they led up to a molecular theory of gases which serves to explain them, as well as certain other gaseous phenomena to be subsequently noted, will be shown in the second part of this work.

By the end of the period with which we are concerned—that is, the middle of the nineteenth century—a considerable body of information had been accumulated as to the conditions which determine the different states of aggregation of matter—that is, the conditions which allow of the passage of the gaseous state into that of the liquid, and of the liquid into that of the solid. That the same substance was capable of existence in the three states of gas, liquid, and solid was of course evident from the case of water. Even the most primitive races must have realised that steam, dew, rain, snow, hail, and ice were only modifications of one and the same substance. As knowledge increased, other substances came to be known which resembled water in their capacity for existence in various physical states. It was but natural to assume that this was a general attribute, and that all substances would, sooner or later, be found capable of existence in each of the different conditions of aggregation.

Attempts were made during the first quarter of the last century to prove that all the æriform bodies then known were simply vapours more or less remote from their point of liquefaction, and still further removed from their point of congelation. Monge and Clouet condensed sulphur dioxide some time before 1800; and Northmore, in 1805, liquefied chlorine. But these observations attracted little attention until Faraday, in 1823, independently effected the liquefaction of chlorine, and Davy that of hydrochloric acid. Faraday almost immediately afterwards liquefied sulphur dioxide, sulphuretted hydrogen, carbon dioxide, euchlorine, nitrous oxide, cyanogen, and ammonia.

Other experimenters, among whom may be mentioned Thilorier and Natterer, greatly improved the mechanical appliances for liquefying these gases; liquid carbonic acid and nitrous oxide were obtained in considerable quantities, and employed in the production of cold. Certain of the gases—hydrogen, oxygen, nitrogen, nitric oxide, carbonic oxide, etc.—resisted all attempts to liquefy them; and hence gaseous substances came to be classified as permanent and non-permanent, depending upon whether they could or could not be liquefied. The division was felt to be irrational even at the time it was made. There seemed no à priori reason why carbon dioxide and nitrous oxide should be liquefiable, while carbonic oxide and nitric oxide should resist all attempts to coerce them into changing their state. The real clue to the conditions required to effect the liquefaction of a gas was not discovered until nearly half a century later, when, as will be shown subsequently, the arbitrary division of gases into permanent and non-permanent was swept away.

The discovery of the law of gaseous combination by Gay Lussac, and the recognition by Ampère and Avogadro of the relation between the density of a gas or a vapour and its atomic weight, early led to improvements in the methods of determining the absolute weights of gases and vapours, especially by French chemists. Both Gay Lussac and Dumas devised processes for determining vapour densities which were in use until late in the century, and which, although now superseded by more convenient and more rapid modifications afforded valuable information concerning the molecular weights of substances and the phenomena of gaseous dissociation.

During the first decade of the nineteenth century Dalton and Henry discovered the simple law which connects pressure with the solubility of a gas in any solvent upon which it exerts no specific action. Dalton further developed the law so as to include the absorption by a solvent of the several constituents of a gaseous mixture.

Attempts were made by Schröder, Kopp, and others, to discover relations between the weights of unit volumes of liquids and solids and their chemical nature; but such attempts were only partially successful, owing to the difficulty of finding valid conditions of comparison. By comparing the specific gravities of liquids at their boiling-points Kopp succeeded in detecting a number of regularities among their specific volumes which seem to indicate that a comprehensive generalisation connecting them may yet be discovered. Kopp has also shown that regularities exist among the boiling-points of correlated substances, and that there is an interdependence between the temperature of their ebullition and the chemical characters of compounds.

This short summary will suffice to show that attempts to discover relations between the physical attributes of substances and their chemical nature were made more or less sporadically from the time that chemistry was pursued in the spirit of science. But it is only in recent times that any great accession to knowledge has resulted from such efforts. The science of physical chemistry is practically a creation of our own period. Its systematic study may be said to date only from the last quarter of the nineteenth century, since which time it has made extraordinary progress. Its broad features will be dealt with in the second volume of this work.