The second half of the 17th century witnessed remarkable transitions and developments in all branches of natural science, and the facts accumulated by preceding generations Boyle. during their generally unordered researches were replaced by a co-ordination of experiment and deduction. From the mazy and incoherent alchemical and iatrochemical doctrines, the former based on false conceptions of matter, the latter on erroneous views of life processes and physiology, a new science arose—the study of the composition of substances. The formulation of this definition of chemistry was due to Robert Boyle. In his Sceptical Chemist (1662) he freely criticized the prevailing scientific views and methods, with the object of showing that true knowledge could only be gained by the logical application of the principles of experiment and deduction. Boyle’s masterly exposition of this method is his most important contribution to scientific progress. At the same time he clarified the conception of elements and compounds, rejecting the older notions, the four elements of the “vulgar Peripateticks” and the three principles of the “vulgar Stagyrists,” and defining an element as a substance incapable of decomposition, and a compound as composed of two or more elements. He explained chemical combination on the hypotheses that matter consisted of minute corpuscles, that by the coalescence of corpuscles of different substances distinctly new corpuscles of a compound were formed, and that each corpuscle had a certain affinity for other corpuscles.
Although Boyle practised the methods which he expounded, he was unable to gain general acceptance of his doctrine of elements; and, strangely enough, the theory which next dominated chemical thought was an alchemical Phlogistic theory. invention, and lacked the lucidity and perspicuity of Boyle’s views. This theory, named the phlogistic theory, was primarily based upon certain experiments on combustion and calcination, and in effect reduced the number of the alchemical principles, while setting up a new one, a principle of combustibility, named phlogiston (from φλοιστός, burnt). Much discussion had centred about fire or the “igneous principle.” On the one hand, it had been held that when a substance was burned or calcined, it combined with an “air”; on the other hand, the operation was supposed to be attended by the destruction or loss of the igneous principle. Georg Ernst Stahl, following in some measure the views held by Johann Joachim Becher, as, for instance, that all combustibles contain a “sulphur” (which notion is itself of older date than Becher’s terra pinguis), regarded all substances as capable of resolution into two components, the inflammable principle phlogiston, and another element—“water,” “acid” or “earth.” The violence or completeness of combustion was proportional to the amount of phlogiston present. Combustion meant the liberation of phlogiston. Metals on calcination gave calces from which the metals could be recovered by adding phlogiston, and experiment showed that this could generally be effected by the action of coal or carbon, which was therefore regarded as practically pure phlogiston; the other constituent being regarded as an acid. At the hands of Stahl and his school, the phlogistic theory, by exhibiting a fundamental similarity between all processes of combustion and by its remarkable flexibility, came to be a general theory of chemical action. The objections of the antiphlogistonists, such as the fact that calces weigh more than the original metals instead of less as the theory suggests, were answered by postulating that phlogiston was a principle of levity, or even completely ignored as an accident, the change of qualities being regarded as the only matter of importance. It is remarkable that this theory should have gained the esteem of the notable chemists who flourished in the 18th century. Henry Cavendish, a careful and accurate experimenter, was a phlogistonist, as were J. Black, K. W. Scheele, A. S. Marggraf, J. Priestley and many others who might be mentioned.
Descriptive chemistry was now assuming considerable proportions; the experimental inquiries suggested by Boyle were being assiduously developed; and a wealth of observations Lavoisier. was being accumulated, for the explanation of which the resources of the dominant theory were sorely taxed. To quote Antoine Laurent Lavoisier, “... chemists have turned phlogiston into a vague principle, ... which consequently adapts itself to all the explanations for which it may be required. Sometimes this principle has weight, and sometimes it has not; sometimes it is free fire and sometimes it is fire combined with the earthy element; sometimes it passes through the pores of vessels, sometimes these are impervious to it; it explains both causticity and non-causticity, transparency and opacity, colours and their absence; it is a veritable Proteus changing in form at each instant.” Lavoisier may be justly regarded as the founder of modern or quantitative chemistry. First and foremost, he demanded that the balance must be used in all investigations into chemical changes. He established as fundamental that combustion and calcination were attended by an increase of weight, and concluded, as did Jean Rey and John Mayow in the 17th century, that the increase was due to the combination of the metal with the air. The problem could obviously be completely solved only when the composition of the air, and the parts played by its components, had been determined. At all times the air had received attention, especially since van Helmont made his far-reaching investigations on gases. Mayow had suggested the existence of two components, a spiritus nitroaerus which supported combustion, and a spiritus nitri acidi which extinguished fire; J. Priestley and K. W. Scheele, although they isolated oxygen, were fogged by the phlogistic tenets; and H. Cavendish, who had isolated the nitrogen of the atmosphere, had failed to decide conclusively what had really happened to the air which disappeared during combustion.
Lavoisier adequately recognized and acknowledged how much he owed to the researches of others; to himself is due the co-ordination of these researches, and the welding of his results into a doctrine to which the phlogistic theory ultimately succumbed. He burned phosphorus in air standing over mercury, and showed that (1) there was a limit to the amount of phosphorus which could be burned in the confined air, (2) that when no more phosphorus could be burned, one-fifth of the air had disappeared, (3) that the weight of the air lost was nearly equal to the difference in the weights of the white solid produced and the phosphorus burned, (4) that the density of the residual air was less than that of ordinary air. The same results were obtained with lead and tin; and a more elaborate repetition indubitably established their correctness. He also showed that on heating mercury calx alone an “air” was liberated which differed from other “airs,” and was slightly heavier than ordinary air; moreover, the weight of the “air” set free from a given weight of the calx was equal to the weight taken up in forming the calx from mercury, and if the calx be heated with charcoal, the metal was recovered and a gas named “fixed air,” the modern carbon dioxide, was formed. The former experiment had been performed by Scheele and Priestley, who had named the gas “phlogisticated air”; Lavoisier subsequently named it oxygen, regarding it as the “acid producer” (ὀξύς, sour). The theory advocated by Lavoisier came to displace the phlogistic conception; but at first its acceptance was slow. Chemical literature was full of the phlogistic modes of expression—oxygen was “dephlogisticated air,” nitrogen “phlogisticated air,” &c.—and this tended to retard its promotion. Yet really the transition from the one theory to the other was simple, it being only necessary to change the “addition or loss of phlogiston” into the “loss or addition of oxygen.” By his insistence upon the use of the balance as a quantitative check upon the masses involved in all chemical reactions, Lavoisier was enabled to establish by his own investigations and the results achieved by others the principle now known as the “conservation of mass.” Matter can neither be created nor destroyed; however a chemical system be changed, the weights before and after are equal.[2] To him is also due a rigorous examination of the nature of elements and compounds; he held the same views that were laid down by Boyle, and with the same prophetic foresight predicted that some of the elements which he himself accepted might be eventually found to be compounds.
It is unnecessary in this place to recapitulate the many results which had accumulated by the end of the 18th century, or to discuss the labours and theories of individual workers since these receive attention under biographical headings; in this article only the salient features in the history of our science can be treated. The beginning of the 19th century was attended by far-reaching discoveries in the nature of the composition of compounds. Investigations proceeded in two directions:—(1) the nature of chemical affinity, (2) the laws Chemical Affinity. of chemical combination. The first question has not yet been solved, although it has been speculated upon from the earliest times. The alchemists explained chemical action by means of such phrases as “like attracts like,” substances being said to combine when one “loved” the other, and the reverse when it “hated” it. Boyle rejected this terminology, which was only strictly applicable to intelligent beings; and he used the word “affinity” as had been previously done by Stahl and others. The modern sense of the word, viz. the force which holds chemically dissimilar substances together (and also similar substances as is seen in di-, tri-, and poly-atomic molecules), was introduced by Hermann Boerhaave, and made more precise by Sir Isaac Newton. The laws of chemical combination were solved, in a measure, by John Dalton, and the solution expressed as Dalton’s “atomic theory.” Lavoisier appears to have assumed that the composition of every chemical compound was constant, and the same opinion was the basis of much experimental inquiry at the hands of Joseph Louis Proust during 1801 to 1809, who vigorously combated the doctrine of Claude Louis Berthollet (Essai de statique chimique, 1803), viz. that fixed proportions of elements and compounds combine only under exceptional conditions, the general rule being that the composition of a compound may vary continuously between certain limits.[3]
This controversy was unfinished when Dalton published the first part of his New System of Chemical Philosophy in 1808, although the per saltum theory was the most popular. Dalton. Led thereto by speculations on gases, Dalton assumed that matter was composed of atoms, that in the elements the atoms were simple, and in compounds complex, being composed of elementary atoms. Dalton furthermore perceived that the same two elements or substances may combine in different proportions, and showed that these proportions had always a simple ratio to one another. This is the “law of multiple proportions.” He laid down the following arbitrary rules for determining the number of atoms in a compound:—if only one compound of two elements exists, it is a binary compound and its atom is composed of one atom of each element; if two compounds exist one is binary (say A + B) and the other ternary (say A + 2B); if three, then one is binary and the others may be ternary (A + 2B, and 2A + B), and so on. More important is his deduction of equivalent weights, i.e. the relative weights of atoms. He took hydrogen, the lightest substance known, to be the standard. From analyses of water, which he regarded as composed of one atom of hydrogen and one of oxygen, he deduced the relative weight of the oxygen atom to be 6.5; from marsh gas and olefiant gas he deduced carbon = 5, there being one atom of carbon and two of hydrogen in the former and one atom of hydrogen to one of carbon in the latter; nitrogen had an equivalent of 5, and so on.[4]
The value of Dalton’s generalizations can hardly be overestimated, notwithstanding the fact that in several cases they needed correction. The first step in this direction was effected by the co-ordination of Gay Lussac’s observations on the combining volumes of gases. He discovered that gases always combined in volumes having simple ratios, and that the volume of the product had a simple ratio to the volumes of the reacting gases. For example, one volume of oxygen combined with two of hydrogen to form two volumes of steam, three volumes of hydrogen combined with one of nitrogen to give two volumes of ammonia, one volume of hydrogen combined with one of chlorine to give two volumes of hydrochloric acid. An immediate inference was that the Daltonian “atom” must have parts which enter into combination with parts of other atoms; in other words, there must exist two orders of particles, viz. (1) particles derived by limiting mechanical subdivision, the modern molecule, and (2) particles derived from the first class by chemical subdivision, i.e. particles which are incapable of existing alone, but may exist in combination. Additional evidence as to the structure of the molecule was discussed by Avogadro in 1811, and by Ampere in 1814. From the gas-laws of Boyle and J.A.C. Charles—viz. equal changes in temperature and pressure occasion equal changes in equal volumes of all gases and vapours—Avogadro deduced the law:—Under the same conditions of temperature and pressure, equal volumes of gases contain equal numbers of molecules; and he showed that the relative weights of the molecules are determined as the ratios of the weights of equal volumes, or densities. He established the existence of molecules and atoms as we have defined above, and stated that the number of atoms in the molecule is generally 2, but may be 4, 8, &c. We cannot tell whether his choice of the powers of 2 is accident or design.
Notwithstanding Avogadro’s perspicuous investigation, and a similar exposition of the atom and molecule by A. M. Ampere, the views therein expressed were ignored both by Berzelius. their own and the succeeding generation. In place of the relative molecular weights, attention was concentrated on relative atomic or equivalent weights. This may be due in some measure to the small number of gaseous and easily volatile substances then known, to the attention which the study of the organic compounds received, and especially to the energetic investigations of J. J. Berzelius, who, fired with enthusiasm by the original theory of Dalton and the law of multiple proportions, determined the equivalents of combining ratios of many elements in an enormous number of compounds.[5] He prosecuted his labours in this field for thirty years; as proof of his industry it may be mentioned that as early as 1818 he had determined the combining ratios of about two thousand simple and compound substances.
We may here notice the important chemical symbolism or notation introduced by Berzelius, which greatly contributed to the definite and convenient representation of chemical composition Chemical notation. and the tracing of chemical reactions. The denotation of elements by symbols had been practised by the alchemists, and it is interesting to note that the symbols allotted to the well-known elements are identical with the astrological symbols of the sun and the other members of the solar system. Gold, the most perfect metal, had the symbol of the Sun, ☉; silver, the semiperfect metal, had the symbol of the Moon, ☽; copper, iron and antimony, the imperfect metals of the gold class, had the symbols of Venus ♀, Mars ♂, and the Earth ♁; tin and lead, the imperfect metals of the silver class, had the symbols of Jupiter ♃, and Saturn ♄; while mercury, the imperfect metal of both the gold and silver class, had the symbol of the planet, ☿. Torbern Olof Bergman used an elaborate system in his Opuscula physica et chemica (1783); the elements received symbols composed of circles, arcs of circles, and lines, while certain class symbols, such as