CHEMISTRY

Theprogress of the science of chemistry forms one phase of the progress of human thought. While at first mankind was contented to observe certain phenomena, and to utilize them for industrial purposes, if they were found suitable, “philosophers,” as the thinking portion of our race loved to call themselves, have always attempted to assign some explanation for observed facts, and to group them into similars and dissimilars. It was for long imagined, following the doctrines of the Greeks and of their predecessors, that all matter consisted of four elements or principles, names which survive to this day in popular language. These were “fire,” “air,” “water,” and “earth.” It was not until the seventeenth century that Boyle in hisSceptical Chymist(1661) laid the foundations of the modern science, by pointing out that it was impossible to explain the existence of the fairly numerous chemical substances known in his day, or the changes which they can be made to undergo, by means of the ancient Greek hypotheses regarding the constitution of matter. He laid down the definition of the modern meaning of the word “element”; he declined to accept the current view that the properties of matter could be modified by its assimilating the qualities of fire, air, earth, or water, and he defined an element as theconstituentof a compound body. The first problem, then, to be solved, was to determine which of the numerous forms of matter were to be regarded as elementary, and which are compound, or composed of two or more elements in a state ofcombination; and to produce such compounds by causing the appropriate elements to unite with each other.

One of the first objects to excite curiosity and interest was the air which surrounds us, and in which we live and move and have our being. It was, however, endowed with a semi-spiritual and scarcely corporeal nature in the ideas of our ancestors, for it does not affect the senses of sight, smell, or taste, and though it can be felt, yet it eludes our grasp. The word “gas,” moreover, was not invented until Van Helmont devised it to designate various kinds of “airs” which he had observed. The important part which gases play in the constitution of many chemical compounds was accordingly overlooked; and, indeed, it appeared to be almost as striking a feat of necromancy to produce a quantity of a gas of great volume from a small pinch of solid powder as for a “Jinn” of enormous stature but of delicate texture to issue from a brass pot, as related in theArabian Nights Entertainments. Gradually, however, it came to be recognized, not merely that gases have corporeal existence, but that they even possess weight. This, though foreshadowed by Torricelli, Jean Rey, and others, was first clearly proved by Black, professor of chemistry in Edinburgh, in 1752, through his masterly researches, as carbonic acid.

The ignorance of the material nature of gases and of their weight lies at the bottom of the “Phlogistic Theory,” a theory devised by Stahl about the year 1690, to account for the phenomena of combustion and respiration and the recovery or “reduction” of metals from their “earths” by heating with charcoal or allied bodies. According to this inverted theory, a substance capable of burning was imagined to contain more or less phlogiston, a principle which it parted with on burning, leaving an earth deprived of phlogiston, or “dephlogisticated,” behind if a metal. This earth, when heated with substances richin phlogiston, such as coal, wood, flour, and similar bodies, recovered the phlogiston, which it had lost on burning, and, with the added phlogiston, its metallic character. Other substances, such as phosphorus and sulphur, gave solids or acid liquids, to which phlogiston was not so easy to add; but even they could be rephlogisticated. On this hypothesis, it was the earths, and such acid liquids as sulphuric or phosphoric acids, which were the elements; the metals and sulphur and phosphorus were their compounds with phlogiston.

The discovery of oxygen by Priestley and by Scheele in 1774, and the explanation of its functions by Lavoisier during the following ten years, gave their true meaning to these phenomena. It was then recognized that combustion was union with oxygen; that an “earth” or “calx” was to be regarded as the compound of a metal with oxygen; that when a metal becomes tarnished, and converted into such an earthy powder, it is being oxidized; that this oxide, on ignition with charcoal or carbon, or with compounds such as coal, flour, or wood, of which carbon is a constituent, gives up its oxygen to the carbon, forming an oxide of carbon, carbonic oxide on the one hand, or carbonic “acid” on the other, while the metal is reproduced in its “reguline” or metallic condition, and that the true elements are metals, carbon, sulphur, phosphorus, and similar bodies, and not the products of their oxidation.

The discovery that air is in the main a mixture of nitrogen, an inert gas, and oxygen, an active one, together with a small proportion of carbonic “acid” (or, as it is now termed, anhydride)—a discovery perfected by Rutherford, Black, and Cavendish—and that water is a compound with oxygen of hydrogen, previously known as inflammable air, by Cavendish and by Watt, finally overthrew the theory of phlogiston; but at the beginning of this century it still lingered on, and was defended byPriestley until his death in 1804. Such, in brief, was the condition of chemical thought in the year 1800. Scheele had died in 1786, at the early age of forty-four; Lavoisier was one of the victims of the French Revolution, having been guillotined in 1794; Cavendish had ceased to work at chemical problems, and was devoting his extraordinary abilities to physical problems of the highest importance, while living the life of an eccentric recluse, and Priestley, driven by religious persecution from England to the more tolerant shores of America, was enjoying a peaceful old age, enlivened by occasional incursions into the region of sectarian controversy.

The first striking discovery of our century was that of the compound nature of the alkalies and of the alkaline earths. This discovery was made by Humphry Davy. Born in Cornwall in 1778, he began the study of chemistry, self-taught, in 1796; and in 1799 he became director of the “Pneumatic Institution,” an undertaking founded by Dr. Beddoes, at Bristol, for the purpose of experiments on the curative effects of gases in general. Here he at once made his mark by the discovery of the remarkable properties of “laughing gas,” or nitrous oxide. At the same time he constructed a galvanic battery, and began to perform experiments with it in attempting to decompose chemical compounds by its means. In 1801 Davy was appointed professor of chemistry at the Royal Institution, a society or club which had been founded a few years previously by Benjamin Thompson, Count Rumford, for the purpose of instructing and amusing its members with recent discoveries in chemistry and natural philosophy. In 1807 Davy applied his galvanic battery to the decomposition of damp caustic potash and soda, using platinum poles. He was rewarded by seeing globules of metal, resembling mercury in appearance, at the negative pole; and he subsequently proved that these globules, when burned,reproduced the alkali from which they had been derived. They also combined with “oxymuriatic acid,” as chlorine (discovered by Scheele) was then termed, forming ordinary salt, if sodium be employed, and the analogous salt, “muriate of potash,” if the allied metal, potassium, were subjected to combustion. By using mercury as the negative pole, and passing a current through a strong solution of the chloride of calcium, strontium, or barium, Davy succeeded in procuring mixtures with mercury or “amalgams” of their metals, to which he gave the names calcium, strontium, and barium. Distillation removed most of the mercury, and the metal was left behind in a state of comparative purity. The alkali metals, potassium and sodium, were found to attack glass, liberating “the basis of the silex,” to which the name silicon has since been given.

Thus nearly the last of the “earths” had been decomposed. It was proved that not merely were the “calces” of iron, copper, lead, and other well-known metals compounds of the respective metals with oxygen, but Davy showed that lime, and its allies, strontia and baryta, and even silica or flint, were to be regarded as oxides of elements of metallic appearance. To complete our review of this part of the subject, suffice it to say that aluminum, a metal now produced on an industrial scale, was prepared for the first time in 1827 by Wöhler, professor of chemistry at Göttingen, by the action of potassium on its chloride, and alumina, the earthy basis of clay, was shown to be the oxide of the metal aluminum. Indeed, the preparation of this metal in quantity is now carried out at Schoffhausen-on-the-Rhine and at the Falls of Foyers, in Scotland, by electrolysis of the oxide dissolved in melted cryolite, a mineral consisting of the fluorides of sodium and aluminum, by a method differing only in scale from that by means of which Davy isolated sodium and potassium in 1806.

To Davy, too, belongs the merit of having dethroned oxygen from its central position among the elements. Lavoisier gave to this important gas the name “oxygen,” because he imagined it to be the constituent of all acids. He renamed the common compounds of oxygen in such a manner that the term oxygen was not even represented in the name—only inferred. Thus a “nitrate” is a compound of an oxide of nitrogen and an oxide of a metal; a “sulphate,” of the oxide of a metal with one of the oxides of sulphur, and so on. Davy, by discovering the elementary nature of chlorine, showed, first, that it is not an oxide of hydrochloric acid (or muriatic acid as it was then called); and, second, that the latter acid is the compound of the element chlorine with hydrogen. This he did by passing chlorine over white-hot carbon—a substance eminently suited to deprive oxy-compounds of their oxygen—and proving that no oxide of carbon is thereby produced; by acting on certain chlorides, such as those of tin or phosphorus with ammonia, and showing that no oxide of tin or phosphorus is formed; and, lastly, by decomposing “muriatic acid gas” (gaseous hydrogen chloride) with sodium, and showing that the only product besides common salt is hydrogen. Instead, therefore, of the former theory that a chloride was a compound of the unknown basis of oxymuriatic acid with oxygen and the oxide of a metal, he introduced the simpler and correct view that a chloride is merely a compound of the element chlorine with a metal. In 1813 he established the similar nature of fluorine, pointing out that on the analogy of the chlorides it was a fair deduction that the fluorides are compounds of an undiscovered element, fluorine, with metals; and that hydrofluoric acid is the true analogue of hydrochloric acid. The truth of this forecast has been established of recent years by Henri Moissan, who isolated gaseous fluorine by subjecting a mixture of hydrofluoricacid and hydrogen potassium fluoride contained in a platinum U tube to the action of a powerful electric current. He has recently found that the tube may be equally well constructed of copper; and this may soon lead to the industrial application of the process. The difficulty of isolating fluorine is due to its extraordinary chemical energy; for there are few substances, elementary or compound, which resist the action of this pale yellow, suffocating gas. In 1811 iodine, separated by Courtois from the ashes of sea-plants, was shown by Davy to be an element analogous to chlorine. Gay-Lussac subsequently investigated it and prepared many of its compounds; and in 1826 the last of these elements, bromine, was discovered in the mother-liquor of sea-salt by Balard. The elements of this group have been termed “halogens,” or “salt producers.”

While Davy was pouring his researches into the astonished ears of the scientific and dilettante world, John Dalton, a Manchester school-master, conceived a theory that has proved of the utmost service to the science of chemistry, and which bids fair to outlast our day. It had been noticed by Wenzel, by Richter, by Wollaston, and by Cavendish, towards the end of the last century, that the same compounds contain the same constituents in the same proportions, or, as the phrase runs, “possess constant composition.” Wollaston, indeed, had gone one step farther, and had shown that when the vegetable acid, oxalic acid, is combined with potash, it forms two compounds, in one of which the acid is contained in twice as great an amount relatively to the potash as in the other. The namesmonoxalateandbinoxalateof potash were applied to these compounds, to indicate the respective proportions of the ingredients. Dalton conceived the happy idea that by applying the ancient Greek conception of atoms to such facts the relative weights of the atoms could be determined. Illustrating his views withthe two compounds of carbon with hydrogen, marsh gas and olefiant gas, and with the two acids of carbon, carbonic oxide, carbonic “acid,” he regarded the former as a compound of one atom of carbon and one of hydrogen, and the second as a compound of one atom of carbon and two of hydrogen, and similarly for the two oxides of carbon. Knowing the relative weights in which these elements enter into combination, we can deduce the relative weights of the atoms. Placing the relative weight of an atom of hydrogen equal to unity, we have:

MarshGasOlefiantGasCarbonicOxideCarbonicAcidCarbon66Carbon56Hydrogen12Oxygen816

Thus the first compound, marsh gas, was regarded by Dalton as composed of an atom of carbon in union with an atom of hydrogen; or, to reproduce his symbols, as⊜◉; while the second, olefiant gas, on this hypothesis, was a compound of two atoms of hydrogen with one of carbon, or◉◍◉. Similarly the symbols◍○, and○◍○were given to the two compounds of carbon with oxygen. So water was assigned the symbol◉○, for Dalton imagined it to be a compound of one atom of hydrogen with one of oxygen. Compounds containing only two atoms were termed by him “binary”; those containing three, “ternary”; four, “quaternary,” and so on. The weight of an atom of oxygen was eight times that of an atom of hydrogen; while that of an atom of carbon was six times as great as the unit. By assigning symbols to the elements, consisting of the initial letters of their names, or of the first two letters, formulas were developed, indicating the composition of the compound, the atomic weights of the elements being assured. Thus, NaO signified a compound of an atom of sodium (natrium), weighing twenty-three times asmuch as a similar atom of hydrogen, with an atom of oxygen, possessing eight times the weight of an atom of hydrogen. Therefore, thirty-one pounds of soda should consist of twenty-three pounds of sodium in combination with eight pounds of oxygen, for, according to Dalton, each smallest particle of soda contains an atom of each element, and the proportion is not changed, however many particles be considered.

It has been pointed out by Judge Stallo, of Philadelphia, in hisConcepts of Physics, that such a hypothesis as that of Dalton is no explanation; that a fact of nature, as, for example, the fact of simple and multiple proportions, is notexplainedby being minified. Allowing the general truth of this statement, it is, nevertheless, undoubted that chemistry owes much to Dalton’s hypothesis—a lucky guess at first, it represents one of the fundamental truths of nature, although its form must be somewhat modified from that in which Dalton conceived it. Dalton’s work was first expounded by Thomas Thomson, professor at Glasgow, in hisSystem of Chemistry, published in 1805; and subsequently in Dalton’s ownNew System of Chemical Philosophy, the three volumes of which were published in 1808, in 1810, and in 1827.

The determination of these “Constants of Nature” was at once followed out by many chemists, Thomson among the first. But chief among the chemists who have pursued this branch of work was Jacob Berzelius, a Swede, who devoted his long life (1779–1848) to the manufacture of compounds, and to the determination of their composition, or, as it is still termed, the determination of the “atomic weights”—more correctly, “equivalents”—of the elements of which they are composed. It is to him that we owe most of our analytical methods, for, prior to his time, there were few, if any, accurate analyses. Although Lavoisier had devised a methodfor the analysis of compounds of carbon, viz., by burning the organic compounds in an atmosphere of oxygen contained in a bell-jar over mercury, and measuring the volume of carbon dioxide produced, as well as that of the residual oxygen, Berzelius achieved the same results more accurately and more expeditiously by heating the substance, mixed with chlorate of potassium and sodium chloride, and then estimating the hydrogen as well as the carbon; this process was afterwards perfected by Liebig. Berzelius, however, was able to show that compounds of carbon, like those of other elements, were instances of combination in constant and in multiple proportions.

In 1815 two papers were published in theAnnals of Philosophyby Dr. Prout, which have had much influence on the progress of chemistry. They dealt with the figures which were being obtained by Thomson, Berzelius, and others, at that time supposed to represent the “atomic weights” of the elements. Prout’s hypothesis, based on only a few numbers, was that the atomic weights of all elements were multiples of that of hydrogen, taken as unity. There was much dispute regarding this assertion at the time, but as it was contradicted by Berzelius’s numbers, the balance of opinion was against it. But about the year 1840 Dumas discovered an error in the number (12.12) given by Berzelius as the atomic weight of carbon; and with his collaborator, Stas, undertook the redetermination of the atomic weights of the commoner elements—for example, carbon, oxygen, chlorine, and calcium. This line of research was subsequently pursued alone by Stas, whose name will always be remembered for the precision and accuracy of his experiments. At first Dumas and Stas inclined to the view that Prout’s hypothesis was a just one, but it was completely disproved by Stas’s subsequent work, as well as by that of numerous other observers. It is, nevertheless,curious that a much larger proportion of the atomic weights approximate to whole numbers than would be foretold by the doctrine of chances, and perhaps the last has not been heard of Prout’s hypothesis, although in its original crude form it is no longer worthy of credence.

One of the most noteworthy of the discoveries of the century was made by Gay-Lussac (1778–1850) in the year 1808. In conjunction with Alexander von Humboldt, Gay-Lussac had rediscovered about three years before what had previously been established by Cavendish—namely, that, as nearly as possible, two volumes of hydrogen combine with one volume of oxygen to form water, the gases having been measured at the same temperature and pressure. Humboldt suggested to Gay-Lussac that it would be well to investigate whether similar simple relations exist between the volumes of other gaseous substances when they combine with each other. This turned out to be the case; it appeared that almost exactly two volumes of carbonic oxide unite with one volume of oxygen to form carbon dioxide; that equal volumes of chlorine and hydrogen unite to form hydrochloric acid gas; that two volumes of ammonia gas consist of three volumes of hydrogen in union with one volume of nitrogen, and so on. From such facts, Gay-Lussac was led to make the statement that: The weights of equal volumes of both simple and compound gases, and therefore their densities, are proportional to their empirically found combining weights, or to rational multiples of the latter. Gay-Lussac recognized this discovery of his to be a support for the atomic theory; but it did not accord with many of the then received atomic weights. The assumption that equal volumes of gases contain equal numbers of particles, or, as they were termed by him,molécules intégrantes, was made in 1811 by Avogadro, professor of physics at Turin (1776–1856). This theory, which has proved of the utmost importance to the sciencesboth of physics and of chemistry, had no doubt occurred to Gay-Lussac, and had been rejected by him for the following reasons: A certain volume of hydrogen, say one cubic inch, may be supposed to contain an equal number of particles (atoms) as an equal volume of chlorine. Now these two gases unite in equal volumes. The deduction appears so far quite legitimate that one atom of hydrogen has combined with one atom of chlorine. But the resulting gas occupies two cubic inches, and must therefore contain the same number of particles of hydrogen chloride, the compound of the two elements, as one cubic inch originally contained of hydrogen, or of chlorine. Thus we have two cubic inches containing, of uncombined gases, twice as many particles as is contained in that volume, after combination. Avogadro’s hypothesis solved the difficulty. By premising two different orders of particles, now termedatomsandmolecules, the solution was plain. According to him, each particle, or molecule, of hydrogen is a complex, and contains two atoms; the same is the case with chlorine. When these gases combine, or ratherreact, to form hydrogen chloride, the phenomenon is one of a change of partners; the molecule, the double atom, of hydrogen splits; the same is the case with the molecule of chlorine; and each liberated atom of hydrogen unites with a liberated atom of chlorine, forming a compound, hydrogen chloride, which equally consists of a molecule, or double atom. Thus two cubic inches of hydrogen chloride consist of a definite number of molecules, equal in number to those contained in a cubic inch of hydrogen, plus those contained in a cubic inch of chlorine. The case is precisely similar, if other compounds of gases be considered.

Berzelius was at first inclined to adopt this theory, and indeed went so far as to change many of his atomic weights to make them fit it. But later he somewhat withdrew from his position, for it appeared to him that itwas hazardous to extend to liquids and solids a theory which could be held only of gases. Avogadro’s suggestion, therefore, rested in abeyance until the publication, in 1858, by Cannizzaro, now professor of chemistry in Rome, of an essay in which all the arguments in favor of the hypothesis were collected and stated in a masterly manner. It will be advisable to revert to this hypothesis at a later point, and to consider other guides for the determination of atomic weights.

In 1819, Dulong (1785–1838), director of the Ecole Polytechnique at Paris, and Petit (1791–1820), professor of physics there, made the discovery that equal amounts of heat are required to raise equally the temperature of solid and liquid elements, provided quantities are taken proportional to their atomic weights. Thus, to raise the temperature of 56 grammes of iron through one degree requires approximately the same amount of heat as is required to raise through one degree 32 grammes of sulphur, 63.5 grammes of copper, and so on; these numbers representing the atomic weights of the elements named. In other words,equal numbers of atoms have equal capacity for heat. The number of heat units, or calories (one calory is the amount of heat required to raise the temperature of 1 gramme of water through 1° C.), which is necessary to raise the atomic weight expressed in grammes of any solid or liquid element through 1° C. is approximately 6.2; it varies between 5.7 and 6.6 in actual part. This affords a means of determining the true value of the atomic weight of an element, as the following example will show: The analysis of the only compound of zinc and chlorine shows that it contains 47.49 per cent. of zinc and 52.16 per cent. of chlorine. Now one grain of hydrogen combines with 35.5 grains of chlorine to form 36.5 grains of hydrogen chloride; and, as already remarked, one volume of hydrogen and one volume of chlorine combine, forming two volumes of hydrogen chloride.Applying Avogadro’s hypothesis, one molecule of hydrogen and one molecule of chlorine react to yield two molecules of hydrogen chloride; and as each molecule is supposed to consist in this case of two atoms, hydrogen chloride consists of one atom of each of its constituent elements. The amount of that element, therefore, which combines with 35.5 grains of chlorine may give the numerical value of the atomic weight of the element, if the compound contains one atom of each element; in that case the formula of the above compound would be zinc, and the atomic weight of zinc, 32.7; but if the formula is ZuCl3, the atomic weight of zinc would be 32.7 × 2; if ZuCl3, 32.7 × 3, and so on. The specific heat of metallic zinc enables this question to be solved. For it has been found, experimentally, to be about 0.095; and 6.2 ÷ 0.095 = 65.2, a close approximation to 32.7 × 2 = 65.4. The conclusion is therefore drawn that zinc chloride is composed of one atom of zinc in combination with two atoms of chlorine, that the atomic weight of zinc is 65.4, and that the molecular weight of zinc chloride is 65.4 + (35.5 × 2) = 136.4. Inasmuch as the relative weight of a molecule of hydrogen is 2 (that of an atom being 1), zinc chloride in the gaseous state should be 136.4 ÷ 2 = 68.2 times that of hydrogen, measured at the same temperature and pressure. This has been found, experimentally, to be the case.

The methods of determining the vapor densities, or relative weights of vapors, are three in number; the first method, due to Dumas (1827), consists in vaporizing the substance in question in a bulb of glass or of porcelain, at a known temperature, closing the bulb while still hot, and weighing it after it is cold. Knowing the capacity of the bulb, the weight of hydrogen necessary to fill it at the desired temperature can be calculated, and the density of the vapor thus arrived at. A second method was devised by Gay-Lussac and perfected by A. W. Hofmann(1868); and a third, preferable for its simplicity and ease of execution, is due to Victor Meyer (1881).

In 1858, as already remarked, Cannizzaro showed the connection between these known facts, and for the first time attention was called to the true atomic weights, which were, up to that time, confused with equivalents, or weights of elements required to replace one unit weight of hydrogen. These were generally regarded as atomic weights by Dalton and his contemporaries.

Some exceptions had been observed to the law of Dulong and Petit, viz., beryllium, or glucinium, an element occurring in emeralds; boron, of which borax is a compound; silicon, the component of quartz and flint, and carbon. It was found by Weber that at high temperatures the specific heats of these elements are higher, and the atomic heats approximate to the number of 6.2; but this behavior is not peculiar to these elements, for it appears that the specific heat of all elements increases with rise of temperature.

A certain number of exceptions have also been noticed to the law of Gay-Lussac, which may be formulated: the molecular weight of a compound in a gaseous state is twice its density referred to hydrogen. Thus equal volumes of ammonia and hydrogen chloride unite to form ammonium chloride. It was to be expected that the density should be half the molecular weight, thus:

NH3+ HCl = NH4Cl; and 53.5 ÷ 2 = 26.75 = density.(14+3) (1+35.5) 53.5

But the density actually found is only half that number, viz., 13.37; and for long this and similar cases were supposed to be exceptions to the law of Gay-Lussac, viz., that equal volumes of gases at the same pressure expand equally for equal rise of temperature. In other instances the gradual decrease in density with rise of temperaturecan be followed, as with chloral hydrate, the products of which are chloral and water.

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 namedissociation. 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 pressureis 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, andas 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 SO3as 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 constituentof 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 H2and 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 SO3and K2O, but if the opposite view be correct, then K2is 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-, andmeta-phosphoric acids. To understand the bearingof 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 CO2with 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, beit 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:

H. O. H.; CH3. O. H.; CH3. O. CH3; and NH3; NH2; H3; NH(CH3)2; and N(CH3)3; the first representing compounds following the water type, the latter the ammonia type. This suggestion had been previously made by Laurent, in 1846. But Williamson extended his views to inorganic compounds; thus, sulphuric acid was represented as constructed on the double water type—HO. SO2. OH, being derived from H. O. (H. H) O. H, the two hydrogen atoms enclosed in brackets being replaced by the radical SO2. To these types Gerhardt added the hydrogen and hydrogen chloride types, H.H. and H.Cl; and, later, Kekulé, professor in Bonn (1829), added the marsh gas type C(H)4. The next important step was taken by Frankland, professor in the Royal School of Mines, London; his work, however, had been anticipated by Cunn Brown, professor at Edinburgh University, in a pamphlet even yet little known. It was to attribute to elements one or more powers of combination. To these he gave the name “valency,” and the capacity of possessing valency was called “quantivalence.” Thus hydrogen was taken as a “monad,” or monovalent. Chlorine, because it unites with hydrogen atom to atom, is also a monad. Oxygen, having the power to combine with two atoms of hydrogen, was termed a dyad, or divalent; nitrogen a triad, or trivalent; carbon a tetrad, or tetravalent, and so on. This is evident from inspection of the formulas of their compounds with hydrogen, thus:

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