CHEMISTRY
HISTORICAL INTRODUCTION.—The acquaintance of the ancients with the modes of extracting several of the metals from their ores, and also with the arts of dyeing, tanning, and glassmaking, and their recognition of various kinds of salts, earths, and inflammable substances show that they must have been possessed of a knowledge of a considerable number of chemical facts; but that this knowledge was to any extent gained by experimental research rather than by mere accident, or that when acquired it was applied to the classification of chemical phenomena, or to the establishment of any theory explanatory of them, there is no evidence to show. Until comparatively recent times the principles of metaphysical philosophy were not recognized as distinct from those of chemistry; men of learning gave themselves up to speculation upon the obvious physical characteristics of matter, whilst they neglected the indirect observation of their intrinsic and specific properties; analogies were a sufficient basis for the classification of bodies, and a consideration of their external peculiarities—“a decomposition of bodies,” to use the words of Whewell, “into adjectives, not into substantives ”—stood in the place of analysis. Thus we find that the qualities of the “elements” of the school of Aristotle are all physical, they are dry or humid, warm or cold, light or heavy; the idea of substances distinguished by special chemical properties was as yet no less foreign to men’s minds than a knowledge of their ultimate composition.
Ideas similar to those of Aristotle concerning the elementary constitution of the universe were early prevalent in the East, whence they appear to have found their way into Europe. The elements, according to the Hindus, were earth, air, fire, water, and ether , and in the fourth book of “Chow,” forming part of the Chinese historical records known as the Shoo King, there is a document supposed to date from 2000 B.C., in which is given an account of the five elements, namely, earth, fire, water, metal, and wood. Of what precisely was meant by an “element” in the language of the ancient philosophers and early chemists it is difficult to get any definite idea; the term could hardly, in fact, be used otherwise than in a vague sense before the exact processes of chemical analysis had shown that the properties of matter vary according to the presence or absence within it of definite quantities of certain substances, distinct in properties from one another, and un-resolvable into other substances.
To the doctrine of a plurality of elements, as opposed to the systems of Thales and Heraclitus, may be ascribed the origin of the conception that by the analysis and synthesis of bodies the various kinds of matter with all their diversity of physical features might be produced, a conception that took practical shape in the processes of alchemy, which, as Liebig has remarked, “was never at any time different from chemistry.” During the alchemistical period a knowledge of the properties of bodies was acquired; afterwards chemistry showed the relations, connections, and limits of these properties. The first mention of chemistry (χημeia) is found in the dictionary of Suidas, who flourished in the 11th century; he defines it as “the preparaticn of silver and gold,” and relates that Diocletian, lest the Egyptians should become rich and capable of resisting the Roman power, caused their books on chemistry to be burnt. He further asserts that the art was known as early as the period of the Argonautic expedition, the golden fleece being a treatise written on skins (δ∈pμασι) concerning the making of gold. The belief in the art of making gold and silver, held by the Greeks from the 5th to the 15tlι century, was by them communicated to the Arabs, possibly not long after the conquest of Egypt in 640; and from the 11th to the 15th century alchemy was diligently studied by the philosophers of Italy, France, Germany, and England.
That the claims of alchemy, notwithstanding repeated demonstrations of their futility, so long received the serious attention of mankind, is attributable to various causes. Not only did impostors find free scope in the credulity of an age of ignorance for the exercise of their arts; but men of talent and culture, relying on tradition, were led honestly to support the doctrine of the transmutation of metals. The existence of the philosopher’s stone having once been accepted as an ascertained fact, it is not extraordinary that Isaacus Hollandus is able to indicate the method of its preparation from “adamic” or “virgin” earth, and its action when medicinally employed; that Roger Bacon, Raymond Lully, Basil Valentin, and John Price know the exact quantities of it to be used in transmutation; and that George Rippel, in the 15th century, has grounds for regarding its action as similar to that of a ferment. In the view of some alchemists, the ultimate principles of matter were Aristotle’s four elements, the proximate constituents were sulphur and mercury, the father and mother of the metals; gold was supposed to have attained to the perfection of its nature by passing in succession through the forms of lead, brass, and silver; gold and silver were held to contain very pure red sulphur and white quicksilver, whereas in the other metals these materials were coarser and of a different colour. Geber, judging from an analogy instituted between the healthy human being and gold, the most perfect of the metals, regarded silver, mercury, copper, iron, lead, and tin in the light of lepers that required to be healed.
To the evidence of their imaginations the alchemists were able to add that of actual observation; the fact that many ores resembling metals were changed and decomposed by heat could not but offer support to theories formed at a time when the nature of chemical combination was not understood; and the apparent transition of many bodies into one another, as, for example, that of clouds into water, was not less wonderful to them than the transmutation of the lighter metals into gold.
It was in the 16th century that a new race of alchemists, or spagyrists, as they were termed, arose, who, abandoning the search for the philosopher’s stone, began to direct their energies to the discovery of chemical remedies for the various diseases of the body. “The true use of chemistry,” says Paracelsus (1493—1541) “is not to make gold, but to prepare medicines.” Rejecting the teaching of Galen, he admitted three or four elements, the star, the root, the element, and the sperm or true seed, which were originally confounded together in the chaos or yliados ; these elements he asserted were composed of the three principles sideric salt and sulphur and mercury, the cause respectively of the qualities of fixity, combustibility, and fluidity and volatility. The theories of Paracelsus found many advocates, amongst whom may be mentioned Thurneysser (1531-1596), Bodenstein, Taxites, Dorn, Sennert, and Duchesne; and with some modifications they were maintained in the 17th century by Dr Willis (1621-1675), the celebrated English anatomist and iatro-chemist, and by Lefebvre and Lémery in France, according to whose system matter consisted of the active principles mercury or spirit, sulphur or oil, and salt, and the passive principles water or phlegm and earth.
Among the contemporaries but not the followers of Paracelsus, the German metallurgist Agricola (1494-1555) deserves mention; his great work, De Re Metallica, is the most valuable contribution to practical chemistry that appeared in the 16th century. Libavius also, who diedin 1616, did much to forward chemical science at this period. From his writings, however, in which he puts forward the views both of Paracelsus and of Aristotle concerning the constitution of bodies, it does not appear that his notions of chemical combination were more definite than those of his predecessors.
J. B. Van Helmout (1577-1644), who, like Paracelsus, repudiated the doctrines of the Galenists, held opinions that in many respects were no advance upon those of the former. He looked upon water as the true principle of all existing things, inclusive of the three principles salt, sulphur, and mercury, which therefore were not elements; to air, however, he granted the rank of a true element. The archoeus — something without form, and independent of the elements—he imagined to draw all bodies from water, to which its generating spirit was attracted by the odour of a ferment or aura vitalis. The vapour produced by the fermentation of water was, according to Van Helmont, a gas, and the same term was by him for the first time applied to carbon dioxide, which he termed gas sylvestre, and to other bodies resembling air.
To Francis de la Boë Sylvius (1614-1672), who studied with care the works of Van Helmont and of Descartes, is due the foundation of the iatro-chemical sect among physicians. In his view the health of the human frame depends upon the relation of its fluids, which were acid and alkaline (acidum aud lixivum), and these by union produced a neutral and milder substance; two kinds of diseases were distinguished, the result either of alkaline or of acid acridity. The new doctrine served to explain many chemical facts, and led to the establishment by Lémery and Macquer of a distinction between acid and alkaline or, as they were afterwards called, basic compounds. This recognition of this chemical difference in bodies and their consequent disposition to unite prepared the way for the conception of chemical attraction or affinity.
In the works of Glauber (1604-1668), alchemy, the preparation of chemical medicines, and the processes employed by him for that end are treated of. His Miraculum Mundi has for its subject the virtues of the sal mirabile, sulphate of sodium, or Glauber’s salt, of which he was the discoverer; and in other of his works he describes various chlorides of the metals, the sulphates of iron and copper, aud sulphuric, nitric, and hydrochloric acids, but with respect to their ultimate constitution he advances no theory; he variously states in his different works that mercury and salt are the principles of all metals, that salt is the origin of all things, aud again that water and earth have produced all the minerals and metals.
The first to attempt to overthrow the doctrines of the iatro-chemists was Robert Boyle (1627-1691), who in 1661 published the first edition of The Sceptical Chemist, or Ghymico-Physical Doubts and Paradoxes touching the Experiments whereby vulgar Spagyrists are wont to endeavour to evince their Salt, Sulphur, and Mercury to be the true Principles of Things, a treatise in which he shows the doubtful character of the doctrine of the threefold constitution of matter, and lays stress upon the influence of heat in the formation of new bodies, not necessarily preexistent as such in the substances from which they are produced. If, as he tells us, he is somewhat too indulgent of suspicion against the hypotheses or arguments of other chemists, he is only acting in compliance with the advice of Aristotle, and bending a crooked stick the contrary way, to reduce it at length to straightness. Into the mouth of
“Themistius” he puts the complaint that “Aristotle’s hypothesis had not been called in question till in the last century Paracelsus and some few other sooty empiricks, having their eyes darkened and their brains troubled with the smoke of their furnaces, began to rail at the Peripatetick doctrine, which they were too illiterate to understand, and to tell the credulous world that they could see the three ingredients in mixed bodies, which, to gain themselves the repute of inventera, they endeavoured to disguise by calling them—instead of earth, and fire, and vapour—salt, sulphur, and mercury, to which they gave the canting title of hypostatical principles.” Boyle inclines to a belief in “but one universal matter of things, as it is known that the Aristotelians themselves acknowledge, who called it ĭnateria prima; the portions of this matter seem to differ from one another in certain qualities or accidents, fewer or more.” He thinks that elementary corpuscles are of various sizes, and of more sorts than three or four or five; and that the combination of two of these corpuscles may give rise to a new body as really one as either of the corpuscles before they were mingled or confounded, this concretion being endowed with distinct qualities, and no more by fire or any known way of analysis divisible into the corpuscles that had first concurred to make it, than either of them could by the same means be subdivided into other particles. He furthermore deduces from his arguments the corollary, “That it may as yet be doubted whether or no there be any determinate number of elements; or, if you please, whether or no all compound bodies do consist of the same number of elementary ingredients or material principles.” In another work, The Imperfections of the Chemical Doctrine of Qualities, Boyle points out the arbitrary nature of Sylvius’s classification of all substances as acids and alkalies, and the needlessness and unsatisfactory character of his hypotheses. Iatro-chemistry was opposed also by Conring (1606—1681), Sydenham (1624-1689), Pitcairne (1652-1713), and his pupil Boerhaave (1668-1738), the author of the excellent Elementa Chemiae; and though vigorously supported by De Blegny, Borrichius, Viridet, Vieussens, and others, it gradually lost repute, and was finally overthrown by F. Hoffmann (1660-1742).
Of the labours of Kunckel (1630-1703) in the cause of chemistry it is impossible to give an account within the com- 'pass of the present sketch; but whilst the science was enriched by means of his numerous researches, amongst which may be mentioned those on phosphorus, it received no assistance from his theoretical views concerning the constitution of bodies; thus, for instance, he rejected the belief in the three principles of Paracelsus, yet maintained that all metals contained common quicksilver; and though their increase in weight by calcination was not, according to him, due to the absorption of ponderable fiery material, the explanation he offered of this phenomenon was even less satisfactory.
To Becher (1635-1682) and to Stahl (1660-1734) chemistry owes the introduction of the first consistent theory of the constitution of compounds and of chemical action. Becher held that the primary ingredients of matter were water and earth, and that from these were produced three earths—the fusible or stony, the fatty, and the fluid earths,—improperly called salt, sulphur, and mercury. Stahl, who developed the doctrines of Becher, enumeratedfour elements—water, acid, earth, and phlogiston. Becher had explained the calcination of metals on the supposition that they consisted of an earth and a something of which they became deprived on ignition; the burning of brimstone was, in like manner, thought to be its resolution into an acid and true sulphur, or that combustible part which was dispelled by heat. It was this supposed combustible body to which the name phlogiston (φλoγiστov, combustible) was applied by Stahl — the matena aut principium ignis non ipse ignis. The phlogiston of Stahl answers in some measure to the souls and spirits assigned to metals and salts by the alchemists, or to what Geber called the “humidity,” and Cardan the “celestial heat” of metals. When by means of charcoal a metallic calx was reduced, or a compound containing sulphur was obtained from fused sodium sulphate, phlogiston was supposed to be absorbed from the charcoal, which with 4 amp black and other reducing agents came in time to be regarded as nearly pure phlogiston. Bodies that would not burn were thought to have already parted with their phlogiston. From a consideration of the insolubility of most combustible substances arose the idea that phlogiston was a dry and earthy body, capable of receiving a motion of great velocity—the motus verticillaris — manifested when ignition or flame was produced. John Rey had in 1630 remarked that metals grow heavier when calcined by the absorption of “thickened air,” but had given no general theory of combustion, or explanation why many substances become lighter or are lost sight of when heated. Boyle, too, bad noticed the increase of weight caused by the calcination of metals, and had attributed it to the combination of the latter with heat particles; and Stabel and others were not slow to object that this fact negatived the supposition that calcination consisted in a subtraction of phlogiston; the Stahlians, however, met the difficulty by declaring that substance to be the principle of levity or negative weight.
F. Hoffmann, who contributed greatly to the progress of analytical chemistry in Germany, held with Stahl that sulphur consisted of acid and phlogiston, and that combustible bodies contained something which might be described as phlogiston, but thought it possible that the calces of metals were formed, not by the subtraction of phlogiston, but by the combination of the metals with an acid material. Boerhaave, without directly attacking the phlogistic theory, casts doubts upon the assumption of the existence of a combustible principle and of earthy substances in the metals. The view of Homberg (1652-1715) was that the principle of combustibility in inflammable minerals and in vegetable substances was sulphur; and E. F. Geoffroy (1672-1731) regarded phlogiston as a sulphurous or oily principle. Amongst the most active supporters of the doctrines of Stahl were Neumann (1683-1737); J. II. Pott (1692-1777), distinguished for his researches on the behaviour of mineral substances at high temperatures; Marggraf (1709-1782); and Macquer (1718-1784), the discoverer of arsenic acid. Other celebrated chemists who flourished during the phlogistic period were Réaumur (1683- 1757), Hellot (1685—1766), and Duhamel du Monceau (1700-1782), who first proved the nature of the base of common salt.
The phlogistic theory of Stahl, though incorrect, was of no small assistance towards a true understanding of chemical phenomena. It was based upon experimental data, the interpretation of which served for the correlation of facts of which but vague and enigmatic explanations had formerly been given. The supposed subtraction of phlogiston in the calcination of metals, though equivalent in reality to the addħion of oxygen, was yet a loss of potential energy, by virtue of the combination of the metal with the gas; and the gain of phlogiston was an increase of potential energy, attendant on the removal of oxygen.
It was only in the latter part of the 18th century that the influence of the presence of air upon the formation of many chemical compounds was generally perceived, and that through the use of the balance the nature of gases began to be comprehended and such airy nothingness became commonly regarded as an intimate and necessary constituent of various solid and fluid bodies. The phlogistic theory gave to its adherents so plausible and moreover so circumstantial an account of the modes of chemical action, that facts and observations which caused at a later time a complete revolution in the theory of chemistry, such, for instance, as regarded the existence and properties of oxygen, remained without explanation, and almost unheeded. Robert Hooke, so early as 1665, in his Micrograplιia, foreshadowed the discoveries of Priestley and his contemporaries, when he advanced the opinion that in common air there existed a substance like, if not the same as, that fixed in saltpetre, and which at an elevated temperature dissolved combustibles such as sulphurous bodies with a rapidity sufficient to occasion the motion of fire, and to create light; this solvent he considered to he far less for a given bulk of air than of saltpetre. The investigations of Mayow (1645-1679) are particularly interesting. In treatises published at Oxford in 1668 and 1674 oxygen is actually described by him under the name of fire-air, aerial spirit, and nitre-air ; all acids are said to contain it, and it is necessary for combustion and respiration, processes which are therefore analogous; it is the nitre-air of the atmosphere that causes fermentation and the souring of wines, that produces sulphuric acid from sulphur, and effects the calcination of metals.
Early in the 18th century Newton in his Opticks indicated the nature and the modes of formation of gases. “Dense bodies,” he tells us, “by fermentation rarefy into several sorts of air; and this air by fermentation, and sometimes without it, returns into dense bodies;” and further on he remarks that the particles shaken off from bodies by heat or fermentation, so soon as they are beyond the reach of the attraction of the body, recede from it, and also from one another with great strength, so as sometimes to take up above a million of times more space than they did before in the form of a dense body. This vast contraction and expansion seem to him unintelligible by feigning the particles of air to be springy and ramous, or rolled up like hoops, or by any other means than a repulsive power; the particles of fluids which do not cohere strongly are most easily rarefied into vapour; but those which are grosser, or cohere by a stronger attraction, are not separated without a stronger heat, or perhaps not without fermentation; being rarefied by fermentation they become true permanent air, those particles receding from one another with the greatest force, and being most difficultly brought together, which upon contact adhere most firmly.— (Opticks, bk. iii., qu. 30 and 31, 1730.)
In 1727 Dr Stephen Hales (1677-1761), who had for some years been engaged in investigations similar to those of Mayow, gave to the world in his Statical Essays the collective results of his observations. The atmosphere he describes in this work as a fine elastic fluid, with particles of very different natures floating in it, whereby it is fitted to be the breath of life of vegetables as well as of animals. The effect of respiration and of the burning of sulphur in air is to deprive it of its elasticity; and country air is cleaner and more elastic than that of towns. Elasticity, we read, is not an essential immutable property of air particles; “they are easily changed from an elastic to a fixed state by the strong attraction of the acid, sulphureous, and saline particles, which abound in the air. Whence it is reasonable to conclude that our atmosphere is a chaos, consisting not only of elastic, but also of unelastic air particles, which in great plenty float in it.”— (Stat. Ess., vol. i. 4th ed., 1769.) Hales did not, however, attempt to determine the distinctive properties of the various gaseous constituents of the atmosphere, and of the substances on which he experimented: all are indiscriminately designated “air.” It is air that is generated by fermentation, and that contributes to the briskness of Pyrmont and other mineral waters; 108 cubic inches of air are procured from a cubic inch of iron filings and the same quantity of oil of vitriol; and 33 cubic inches of air are the result of distilling a cubic inch of dog’s blood. Hales determined also the volume of air to be obtained by distilling certain quantities of amber, chalk, coal, grey pyrites, aqua-fortis, antimony, tobacco, and other materials, but apparently with no other end in view than the establishment of the fact that air is contained in a great number of substances. He had learned to interrogate, but not to cross-examine nature.
The first important step towards a knowledge of the specific properties of the various gaseous bodies was that made in the middle of the 18th century by Dr Black, who experimentally proved that the causticity acquired on ignition by mild magnesia and lime was attributable not to the entrance into them of ponderable caloric, but to the expulsion of a peculiar kind of air, which occurred fixed, or in a state of combination, in the unburnt or mild earths, and caused them to be heavier before than after exposure to heat. He found it possible, in fact, to impart to these substances a large amount of heat, which became latent, whilst at the same time their weight was lessened by the loss of “fixed air” (carbon dioxide). It was discovered by Black that alkalies in contact with quicklime became caustic by giving up their fixed air to the lime, which was thereby increased in weight and rendered mild. It was thus, by employing the balance as an experimental test of the composition of bodies, that Black laid the foundation of quantitative chemistry, and in so doing gave the first occasion to the strife that twenty years later began to rage between the followers of Stahl and the antiphlogistians.
Foremost in the number of those who after Black distinguished themselves as pneumatic chemists, was Dr J. Priestley (1733-1804). His first discovery, made in 1772, was nitric oxide gas, which he soon employed in the analysis of air. Boyle, more than a century before Priestley began his experiments, had stated in The Sceptical Chemist, that “without the addition of any extraneous body, quicksilver may by fire alone, and that in glass vessels, be deprived of its silver-like colour, and be turned into a red body; from this red body, without addition, likewise may be obtained a mercury bright and specular as it was before.”— (Boylds Works, ed. Birch, p. 352, Lond. 1744.) On the 1st of August, 1774, Priestley discovered that the red oxide of mercury evolved a gas when heated. This gas (oxgyen) being superior even to the air as a supporter of combustion was regarded by him as dephlogisticated air ; the incombustible part of the atmosphere he supposed to be saturated with phlogiston, on the assumption that a gas was so much the better adapted for supporting combustion as it contained within itself a smaller quantity of that body. Common air, by drawing phlogiston from burning substances, became, as he thought, phlogisticated air, and on that account had no longer any attraction for phlogiston, or, in other words, any power of supporting combustion. The phlogiston evolved in the burning of combustibles and in the calcination of metals was supposed to unite with the atmosphere or the dephlogisticated air contained therein, and that which was produced by the action on the atmosphere of the phlogiston lost by the metals was the cause of the increased weight of their calces. The opinion that the air given off during the solution of metals in acid was their combustible constituent had been advanced in 1700 by Lémery; and Priestley, guided apparently by the notion suggested by Cavendish’s experiments, that to unite with acids metals must part with their phlogiston, considered inflammable air (hydrogen) either as identical with or at least very rich in that principle. In 1788, in order to explain the formation of water from a mixture of inflammable and dephlogisticated air, he put forward the hypothesis that water entered into the composition of these and of fixed and other airs; inflammable air, he thought, might be the principle of alkalinity, dephlogisticated air, as Lavoisier had shown, being the principle of acidity. To the last Priestley was an advocate of the phlogistic philosophy, and though unquestionably one of the fathers of modern chemistry was always, to quote the words of Cuvier, “un père qui ne voulut jamais reconnoitre sa fille.” Besides nitric oxide and nitrogen, Priestley first made known sulphurous acid gas, gaseous ammonia and hydrochloric acid, and carbon monoxide; and he it was who, by showing that the condition of ammoniacal gas and of common air is altered by the transmission of electric sparks, led to Berthollet’s analysis of ammonia, and Cavendish’s discovery of the composition of nitric acid.
Henry Cavendish (1731-1810), who, like Priestley, was of the phlogistic school, contributed by his discoveries and carefully conducted investigations, especially as regards gases, scarcely less than that experimenter to the advance in chemical knowledge which before the beginning of the 19th century effected the subversion of the Stahlian philosophy. To him chemists are indebted for the invention of the pneumatio trough, and to him is due the first recognition of the importance of determining the specific gravities of the various gases. He established the radical difference between hydrogen and nitrogen, and discovered in 1781 that hydrogen and dephlogisticated air (oxygen), when exploded in a close vessel in proportions sufficient almost entirely to phlogisticate the burnt air, produced pure water; and that water was also formed when a mixture of common air and inflammable air was exploded, a reduction of one-fifth of the bulk of the former air being then observable. According to Cavendish, water consisted of phlogiston and dephlogisticated air; inflammable air, of phlogiston and water; the action of dephlogisticated upon inflammable air when exploded with it was to unite with its phlogiston to form water, and consequently to set free the water of the inflammable air; thus both airs became water.— (Trans. Roy. Soc., 1784 and 1785.) Lavoisier had shown in 1770 the incorrectness of the notion prevalent among chemists that water by continued boiling and redistillation could be transformed into an earth. Cavendish’s discovery deprived it of the rank of an element, to which, according to the vague Aristotelian doctrines of the time, it was entitled, and thus prepared the way for the acceptation of correct and definite views concerning the elementary bodies. Lavoisier, availing himself of the facts ascertained by Cavendish, taught that oxygen, the so-calleddephlogisticated air, was an element, and that combined with it was imponderable caloric; inflammable air, or hydrogen, as he termed it, was another element, which had the power of disengaging from caloric a weight equal to its own of oxygen, with which it united to form water. The new doctrine did not, however, meet with very ready acceptance from the phlogistians. “It is inconceivable,” writes one of them, “how water, which is absolutely incombustible, should have so combustible a body as inflammable gas is for one of its component parts; whereas, by admitting pure air in its whole substance to be one of the component parts of water, and the other to consist of the base only of inflammable gas, which being burnt by the passage of the electric spark through it, its phlogiston is converted into light and heat, the whole doctrine of the generation of water becomes plain and easy.”—(Hopson, Chemistry, 1789.)
Among the most eminent of the contemporaries of Priestley and Cavendish that cherished a belief in the existence of phlogiston was the Swedish chemist Scheele (1742-1786). In experiments made to ascertain the nature of heat and. fire, he found that measured quantities of common air, when kept in contact with certain substances, e.y., solution of potassium sulphide or moist iron-filings, contracted in volume, and became incapable of supporting combustion. As the specific gravity of the air had not augmented, the decrease of bulk, Scheele concluded, could not be due, as he had at first conjectured, to the absorption of phlogiston; the atmosphere must, therefore, consist of two distinct bodies. One of these, the residual air, he assumed to be incapable of combining with phlogiston; the other, having a strong attraction for that substance, had united with it, forming heat, which had penetrated through the walls of the vessel containing it— hence the diminution of the original volume of air. Heat, Scheele considered, was decomposed by means of bodies which had a strong attraction for its phlogiston, such as the calces of gold, silver, and mercury, and oil of vitriol mixed with black manganese ore, and consequently the other constituent of heat, empyreal or fire-air (oxygen), became isolated. Heat could be synthesized, for it was produced by the union of the phlogiston of coals with fireair. Light, like heat, was a compound of fire-air and phlogiston, but was richer in the latter constituent, to the varying proportions of which it owed its differences of colour. Subsequently, when it became impossible for Scheele to ignore the consideration of the increase observable in many substances after burning or calcination, he so far modified his views as to regard fire-air as a compound containing, with a very little phlogiston, a saline principle (principium salinum) and water, which last gave to fire-air the greater part of its weight. When fire-air formed heat oy combining with phlogiston, it gave up its water to the materials it dephlogisticated, and thus it was that they were rendered heavier by ignition. Such, in brief, were the theoretical conceptions of Scheele; it is upon his work as a practical chemist that his fame must rest. Tartaric acid was isolated by him in 1769, and he made the discovery in 1774 of baryta and of dephlogisticated muriatic acid (chlorine), and in 1779 of glycerine, the properties of which he pointed out in 1784; in 1781 he demonstrated the nature of hydrofluoric acid, first obtained by him in 1771, and prepared tungstic acid, before unknown; and between the years 1776 and 1786 he discovered benzoic, molybdic, lactic, mucic, oxalic, malic, and gallic acids, and made important observations on the compounds of arsenious acid.
The advance made during the last part of the 18th century in analytical chemistry is attributable in great measure to the labours of Torbern Bergman of Upsala (1735-1784), who devised systematic methods of examining compounds by the wet way, and by means of the blow-pipe, and first rendered it possible to analyze minerals insoluble in acids by fusing them with an alkali or alkaline carbonate. In 1718 E. F. Geoffroy had published tables in which he exhibited the reciprocal chemical affinities of various substances, and these tables had been improved upon by Gellert in his Metallurgic Chemistry, 1751, and by Limbourg, 1761. Bergman, in 1775, gave in a dissertation on elective attractions, as he named affinity, an arrangement in 59 columns of all the chemical substances known at the time, in which was shown the order of their decomposition when in solution, and when exposed to a strong heat. The nature of the compounds formed by the mixture of reagents depended, in Bergman’s estimation, on the sum of their attractions. Bergman contributed also in some measure to the determination of the constitution of neutral salts—a subject treated of by Homberg in 1699, and, after Bergman and Kirwan, investigated by Wenzel in his Vorlesungun über die chemische Verwandtschaft der Körper, published in 1777. From Wenzel’s observations the idea of equivalency took its rise. He showed that the products of the mutual decomposition of two neutral salts were themselves neutral, or, in other words, that the same weight of base satisfies definite quantities of two different acids. Thus, according to his experiments, 123 parts of lime and 222 of potash must be considered equivalent to each other, being both competent to neutralize 181 ’ 5 parts of sulphuric, or 240 parts of nitric acid.
In England, ten years previously to the publication of Wenzel’s treatise, Cavendish described certain quantities of fixed alkali and marble as “equivalent;” and in 1788 he stated that a quantity of oil of vitriol sufficient to produce 100 parts of plumbum ponderosum with sugar of lead would dissolve 33 of marble, since he found by experiment that so much oil of vitriol would saturate as much fixed alkali as a quantity of nitrous acid sufficient to dissolve 33 of marble.— (Phil. Trans, 1767, p. 102; 1788, p. 178.)
In 1792, J. B. Richter (1762-1807) published a work on Stöchiometrie, or the Art of Measuring Chemical Elements, in which he gave in two series of tables the weights of different bodies which neutralized 1000 parts of various acids, and the weights of acids that similarly corresponded to 1000 parts in the case of the bases, and pointed out the proportionality that existed between the weights of the bases, as also of the acids, in each series. There was, he remarked, a constant ratio between the quantity of an acid and the quantity of oxygen in the weights of the bases needed for its complete saturation,—a fact afterwards restated by Berzelius, who show ed that a simple and uniform relation was observable between the amounts of oxygen in the acid and basic portions of salts of the same class.
Although neither the science of chemistry nor yet a change in its objects can be said to have originated with Lavoisier (1743-1794), the means he introduced of attaining those objects, the ideas he put forth concerning the constitution of bodies, and the explanations he gave of various phenomena were new, and gave to the science in the twenty years preceding his death a completely altered aspect. The mine of chemistry had yielded rich returns long before Lavoisier came to labour in it; he availed himself of the old workings and, extending them, opened the main lode. “He discovered,” says Liebig, “no new body, no new property, no natural phenomenon previously unknown; but all the facts established by him were the necessary consequences of the labours of those who preceded him. His merit, his immortal glory consisted in this—that he infused into the body of the science a new spirit; but all the members of that body were already in existence, and rightly joined together.”— (Letters on Chemistry, ii.) In the first of his papers, in 1765, Lavoisier indicates no doubt of the existence of phlogiston. In 1775 he still spoke of it, but in the following year he expressed his conviction that, for the elucidation of certain phenomena, one must ascribe to phlogiston other qualities than those assigned to it by Stahl. It was in 1775 that he presented to the French Academy a memoir On the Nature of the Principle which combines with the Metals during their Calcination, and which augments their Weight, in which he describes the formation of fixed air from charcoal in the reduction of calces, and the preparation and properties of oxygen; but he makes no allusion to the fact that Priestley had in November of the foregoing year made him acquainted with that gas, or to the observations of Bayen in the Journal de Physique, February and April 1774, to the effect that the change of a metal into a calx is attributable to the fixation of an aerial fluid, and that the red oxide of mercury is reduced by heat without the addition of anything, and loses in weight by reduction.
In 1776 Lavoisier, by experiments with phosphorus and sulphur, confirmed results already arrived at by him— that the atmosphere contained two gases, azote or nitrogen, and a highly respirable air, the absorption of which by burning substances accounts for their increase in weight. In 1777 he combated the assertion of Priestley that air is rendered irrespirable by becoming loaded with phlogiston, and showed that air in which candles were burnt furnished about two-fifths of its volume of fixed air, and that pure or dephlogisticated air under the same conditions became almost completely transformed into that gas. In a paper received by the Academy in 1778 Lavoisier broached his theory that the dephlogisticated air of Priestley was the universal acidifying or oxygenizing principle, which by combination with charcoal, sulphur, nitrogen, and phosphorus formed carbonic, vitriolic, nitric, and phosphoric acids, and with metals, calces; all the phenomena of combustion, Lavoisier contended, were explicable without the supposition of the existence of phlogiston, of which there was no evidence. In 1783, when he had determined to discover by carefully conducted experiments the nature of the product of the combustion of hydrogen, Lavoisier learned that Cavendish had ascertained that body to be water. He therefore repeated Cavendish’s experiments, and with the aid of Meusnier proved that water contains hydrogen, by passing steam through a red-hot porcelain tube containing iron wire. Lavoisier had long been unable to hold the doctrine that hydrogen was the phlogiston of metals, seeing that the calces were actually heavier than the metals supposed to contain them, and that the hydrogen supposed to be evolved by the metals, though light, certainly had weight. Lavoisier now explained the production of hydrogen during the solution of metals in acids on the assumption that water was decomposed, its oxygen uniting with the metals, whilst its hydrogen escaped. In the case, however, of nitric acid, oxygen was supposed to be supplied by the acid. The effect of Lavoisier’s reasonings upon his contemporaries is illustrated by Priestley’s remark (Phil. Trans., 1788, p. 155) that, “the doctrine of the decomposition of water being set aside, that of phlogiston (which in consequence of the late experiments on water has been almost universally abandoned) will much better stand its ground.” But the fate of the Stahlian philosophy was sealed; and in 1792 Klaproth and the Berlin Academy gave in their adhesion to the new doctrines. Gren and Wiegleb in Germany, Delametherie in France, and Kirwan in England endeavoured but in vain, to support the phlogistic theory; and ere long Lavoisier’s innovations could be regarded not simply as the anti-phlogistic, but as the universally accepted system of chemistry. “It was the glory of Lavoisier,” wrote Davy in 1814, “to lay the foundation of a sound logic in chemistry by showing that the existence of this principle (phlogiston), or of other principles, should not be assumed where they could not be detected.” In 1789 was published Lavoisier’s Traite élémentaire de Chimie, in which the new chemical doctrines were set forth with remarkable clearness and ability. A list of 33 “simple substances” is given in the 2d part of the 1st vol. of this work, wherein light and caloric are included with oxygen, azote, and hydrogen as elements of bodies; but as to whether light was a modification of caloric, or caloric a modification of light, it was impossible, Lavoisier considered, to decide (Traité élément., tom. 1. part ii. chap. 1). Under the head of “oxidable or acidifiable substances” he placed sulphur, phosphorus, carbon, and the muriatic, fluoric, and boracic radicals; the “oxidable and acidifiable metals” are antimony, silver, arsenic, bismuth, cobalt, copper, tin, iron, manganese, mercury, molybdenum, nickel, gold, platinum, lead, tungsten, and zinc; and the “simple earthy and salifiable substances” are lime, magnesia, baryta, alumina, and silica. Of metallic oxides, he says (Traité élément., tom. i. part i. chap, vil), “they ought not to be considered as completely saturated with oxygen, because their action on that principle is balanced by the force of attraction exercised upon it by caloric. Oxygen, then, in the calcination of metals, obeys really two forces —that exercised by the caloric and that exercised by the metal.” Davy, who by his experiments on the effects of friction on ice did much to dispel the belief in the materiality of heat, but who regarded light as a body in a peculiar state of existence, and consisting of minute particles, held the view that light in phosoxygen (oxygen gas) was intimately combined with oxygen.— (Works, vol. ii. pp. 11-32.)
The following table will serve to show the progress that has been made in the knowledge of the elementary bodies from the time of the publication of Lavoisier’s Traité élémentaire; the second and third columns give the authorities by whom and the dates when these bodies were either first isolated, or were recognized as simple substances :— [table]
Lavoisier was assisted in the establishment of his system by Fourcroy (1755-1809), Monge (1746-1818), Guyton de Morveau (1737-1816), and Berthollet (1748-1822); Lavoisier’s theory that oxygen was the principle of acidity was not, however, accepted by Berthollet. In 1803 appeared Berthollet’s Chemical Statics, in which Bergman’s conclusions with respect to the laws of the combination of acids and bases were disputed. Affinity, it was argued, could not be simply an attraction, for then no decomposition, but only an addition of constituents would take place when solutions of different compounds were mixed together; affinity might be regarded as an attraction between combining substances probably like that existing between the planetary bodies—an attraction dependent on mass, not on elective force, so that combination between the various kinds of matter could take place in all sorts of proportions. The complete removal of a constituent of a compound by means of a decomposing agent was attributed by Berthollet to its elasticity or insolubility in the free state. Thus, he considered that barium sulphate was precipitated by solution of potassium sulphate because it was insoluble in water, but that a triple compound was formed when solutions of the sulphates of sodium and potassium were mingled. The decomposition of ammonium carbonate by sulphuric acid he asserted to be due to the elasticity of the carbonic acid gas that escaped. Berthollet’s hypotheses found an opponent in Proust (1755-1826), who had already written on the constancy of the relations by weight of the component parts of bodies, and had shown that tin and antimony unite with oxygen, and iron with sulphur, in two definite proportions respectively. In a controversy with Berthollet that extended over some years he successfully demonstrated that all compounds are definite, and contain only certain determinate proportions of their constituents.
In August 1804, Dalton (1766-1844) communicated to Dr Thomson his theory of the atomic composition of bodies. Dalton’s ideas had been vaguely foreshadowed by W. Higgins of Pembroke College, Oxford, in a publication entitled A Comparative View of the Phlogistic and Antiphlogistic Theories, where we read that—
“In volatile vitriolic acid a single ultimate particle of sulphur is intimately united only to a single particle of dephlogisticated air; and in perfect vitriolic acid every single particle of sulphur is united to two of dephlogisticated air, being the quantity necessary to saturation. As two cubic inches of light inflammable air require but one of dephlogisticated air to condense them, we must suppose that they contain equal number of divisions, and that the difference of their specific gravity depends chiefly on the size of their ultimate particles; or we must suppose that the ultimate particles of light inflammable air require two or three or more of dephlogisticated air to saturate them. If this latter were the case, we might produce water in an intermediate state, as well as the vitriolic or the nitrous acid, which appears to be impossible; for in whatever proportion we mix our acids, or under whatever circumstance we combine them, the result is invariably the same.”—(2d nd., 1791.)
The atomic theory first suggested itself to Dalton during his investigations on marsh-gas, or light carburetted hydrogen, and olefiant gas. He calculated that if the weight of carbon in each of these compounds were reckoned to be the same, then marsh-gas contained exactly twice the weight of hydrogen present in olefiant gas. He further observed that the quantity of oxygen in carbonic acid gas was twice as much as in carbonic oxide gas. These and similar facts he conceived might be explained by assuming the ultimate particles of matter to be incapable of further division, or, in other words, atoms, possessing definite weights, the ratios of which could be denoted by numbers, the weight of an atom of hydrogen being taken as unity. The combination of these atoms with one another would account then for the definite proportions in which the elements united. For clearness the atoms were represented by symbols, thus :— [table]
Binary compounds with their weights were denoted thus :— [table]
and ternary compounds in a similar manner :— [table]
The weight of the smallest particle of a compound was therefore, according to Dalton’s theory, to be obtained by adding together the weights of its constituent atoms; the laws discovered by Wenzel and Richter were thus included within the law of atomic weights.
The new theory was promulgated in Dr Thomson’s System of Chemistry, 3d edit., 1807. In the same year Thomson showed that the amount of oxalic acid required to form strontium binoxalate was twice as great as that jn the case of the oxalate, supposing the weight of strontium in each salt to be the same; and Dr WoUaston proved that three potassium oxalates were obtainable, to form which a definite amount of potash needed weights of oxalic acid in the proportion of I: 2: 4.
In 1808, the year of the publication of Dalton’s New System of Chemical Philosophy, Gay-Lussac made known to the world the laws of the combinations of gases by volume— to which his attention had been directed by the discovery which he and A. Von Humboldt had made, that a definite volume of oxygen combined with exactly twice its bulk of hydrogen. He pointed out that there is a simple relation between the volumes of two gases which unite together, and also between their collective volume in the uncombined and in the combined condition. Thus, three volumes of hydrogen combine with one of nitrogen to form two volumes of ammonia; one volume of chlorine with one of hydrogen produces two volumes of hydrochloric acid gas; and two volumes of nitrogen and one of oxygen give two volumes of protoxide of nitrogen. The law of definite proportions was shown to hold good with respect to the volumes as well as to the weights of combining bodies.
In 1811 Avogadro, remarking that equal variations of temperature and pressure produce in all gases and vapours the same changes of volume, enunciated the hypothesis that equal volumes of any gas or vapour contain the same number of atoms, and the same doctrine was brought forward in 1814 by Ampère. On this hypothesis the density of equal volumes of gases was shown to represent the relative weights of their atoms; and thus, as it had been discovered that gases frequently do not unite volume for volume, a distinction came to be drawn between atomic weights and equivalents. If we say atom for volume, wrote Berzelius, we find in Gay-Lussac’s discovery one of the most direct arguments in favour of Dalton’s hypothesis.
Berzelius (1779-1848), from considerations based on the law of combination by volume, accounted the atoms of elements distinct from their equivalents. Thus two volumes of hydrogen were recognized as the equivalents of one volume of oxygen, the relative weights of equal volumes of the two gases being those of their atoms. Berzelius adopted 100 parts of oxygen as his standard of atomic weight, the atomic weight of hydrogen was therefore 6 - 24, its equivalent, 12 ’ 48. He considered that the atoms of aluminium, arsenic, antimony, bromine, chlorine, fluorine, hydrogen, iodine, nitrogen, and some other elements had a weight equal to only half that of their equivalents, which latter were double and inseparable atoms. In place of the symbols used by Dalton to denote the proportions in which the elements combine by weight, he employed a notation in which letters were used to indicate the names of the elements. He introduced also an abridged notation, in which the equivalents or double atoms were represented by drawing a bar through the symbol of the element. A dot being used to signify an atom of oxygen, the formula of water was written H: and HC4 denoted hydrochloric acid, which was viewed as consisting of a double atom of hydrogen united to a double atom of chlorine,—an hypothesis which left unexplained the fact that the combination of the so-called double atoms of hydrogen and chlorine resulted always in the formation of two molecules, instead of one, of hydrochloric acid gas. Berzelius constructed a table of atomic weights and equivalents, which the discovery by Dulong and Petit in 1819 of the connection between the specific heats and the atomic weights of the elements, and that of the law of isomorphism by Mitscherlich in 1820, enabled him to modify and improve. The equivalent notation of Berzelius was adopted by Gay-Lussac, and displaced in time that founded upon the atomic weights; but it was not generally applied with strictness to all compounds, molecular and not equivalent formulae being employed in some cases. It had this objection, that it masked the relative atom-fixing powers of the various acids; thus the formula AgO.JPO 5 did not express the tribasic character of a molecule of phosphoric acid, and its relations to the molecules of acids of less basicity. Observing the ratio between the oxygen of bases and acids, Berzelius was led to the conception of the dualistic theory, according to which all compounds consist of paired constituents or groups of constituents. This theory was applied to both organic and inorganic substances, and received considerable support from the development of the doctrine of compound radicles, of which Berzelius was the chief supporter.
De Morveau, in a memoir On the Development of the Principles of Methodical Nomenclature, had in 1787 given the name of radicles to the “acidifiable bases” of acids, and Lavoisier in his Traité élémentaire had spoken of the “hydrocarbon radicles” in oils, starch, sugar, and gum. In 1817' Berzelius, following Lavoisier, held the opinion that all inorganic oxides contained simple radicles, and organic oxides compound radicles. In 1832 Liebig and Wöhler discovered the composition of bitter-almond oil, which they showed, on the assumption of the existence of a radicle C 7 H 5 O, might be compared with the compounds of potassium and other metals. Berzelius at first accepted their explanation of the constitution of benzoyl-compounds, but afterwards rejected it, as his electro-chemical theory did not support the view that oxygen could be a constituent of a radicle. Benzoic acid was represented by him as a compound containing the radicle C 44 H 5 , thus :—C 14 H 5 .O 3 + HO, and alcohol as an oxide of the radicle CH 3 , or 2(CH 3 ).O.
In 1834 Dumas made known his observations on the substitution of hydrogen by other elements; and Laurent subsequently concluded from numerous experiments that, when hydrogen is substituted by an equivalent of chlorine or bromine, these elements take the place occupied by the hydrogen, performing to a certain extent the functions of the latter in the new compound, which is therefore analogous to that from which it has been produced. Thus was given the first blow to the dualistic theory. The discovery of trichloracetic acid by Dumas gave considerable aid in the establishment of the new doctrine,- but Berzelius and others, who were unable to reconcile with fhe electrochemical theory the substitution of an electro-positive by an electro negative element, sought to explain the facts of substitution in accordance with dualistic notions. Berzelius considered that a compound in which - oxygen was a fourth element was at once an oxide and a chloride; trichloracetic acid was, in his opinion, a copulated compound of sesquichlυride of carbon with sesquioxide of carbon (oxalic acid) C 2 C1 3 + C 2 O 3 + HO; acetic acid, on the other hand, was a trioxide of acetyl (C 4 H 3 ) with the formula C 4 H 3 .O 3 + HO. The substitution-compounds discovered by Malaguti and Régnault were in like manner represented dualistically; dichloroformic ether, for example, was written 2C 2 H.O 3 + C 2 H.Cl 3 + 2C 4 H 3 .O 3 + C 4 H q .CL; such complex formulae, however, which implied in most cases a widely different constitution for bodies obviously related in properties, were received with but little favour by chemists. At length Meisens, by converting trichloracetic acid into acetic acid, gave decisive evidence as to the truth of the principle of the substitution theory.
Gerhardt, who regarded all compounds as simple molecules, certain atoms of which were displaceable by double decomposition, discarded entirely the conception of radicles. He classified organic substances according to the number of carbon atoms contained in their molecules; hence the recognition of homologous series of organic compounds. What he termed residues — molecules deprived of certain of their constituents—were in many instances identical with the radicles employed by Berzelius, but were not held necessarily to pre-exist in compounds. Gerhardt was the means of re-introducing, in a modified form, the atomic notation of Berzelius. Observing that the smallest quantities by weight of carbonic acid gas and water produced in reactions were expressed by the formulae C 2 O 4 and H 2 O 2 , he concluded that these must represent the molecules of the two bodies, which might be more conveniently written CO 2 and H 2 O, the atomic weights of hydrogen, oxygen, and carbon being taken as 1, 16, and 12 respectively. He made the atomic weights of bromine, chlorine, fluorine, hydrogen, iodine, nitrogen, and other of the elements equal to those given by Berzelius, thus enabling the formulae of water and most volatile compounds to be expressed in agreement with the law of Gay-Lussac; but he halved the old atomic weights of most of the metals, supposing that all metallic oxides were similar in constitution to water, or contained two atoms of metal to one of oxygen. Régnault afterwards proved that, according to the law of Dulong and Petit, this alteration ought to be made in the case only of the atomic weights of the metals lithium, potassium, sodium, and silver; many metals were accordingly to be regarded as having oxides of the general formula RO. Such metals, being compared with the diatomic radicles of organic chemistry, were called diatomic by Cannizzaro in 1858. Thus originated the idea of the polyatomicity of the metals.
The theory of types dates from the time of the discovery of trichloracetic acid by Dumas, who observed that this body and others similarly obtained must be of the same chemical type with the hydrogenated substances from which they are derived. The discovery by Wurtz of the compound ammonias, and by Hofmann of diethylamine and triethylamine, led to the creation of the ammonia type ; and Williamson, by the discovery in 1850 of the mixed ethers, was enabled to refer ether, alcohol, and acids to the water type, and to predict the existence of acetic anhydride, which, as well as benzoic anhydride, was discovered in 1852 by Gerhardt. To these types Gerhardt added two others, those of hydrogen and hydrochloric acid, and with the former associated the aldehydes, ketones, and many hydrocarbons, e.y., the radicles discovered by Frankland and Kolbe. The theory of types was still further’ extended by Williamson, to whom the conception of condensed types is due; by Odling, who first suggested the idea of representing the relations of compounds by the use of mixed types; and by Kekulé and numerous other investigators.
The foundation of electro-chemical science may be said to have been laid by Nicholson and Carlisle, who in 1800 discovered the decomposition of water by the agency of the voltaic pile; but the earliest electro-chemical experiments were those made by Priestley in 1775 upon ammonia gas, and by Deiman and Van Troostwijk in 1789 upon water, by means of frictional electricity. Cruickshank by experiments with the chlorides of magnesium, sodium, and ammonium demonstrated that when those salts are decomposed by the electric current, alkali always appears at the negative, and acid at the positive pole.
In 1803 Berzelius and Hisinger published their observations on the electrical decomposition of salts and some of their bases; oxygen, acids, and oxidized bases, they said, appear at the positive pole; combustible bodies, alkalies, and earths at the negative. Later experimenters, however, showed that it is the metal, not the oxide, that appears at the negative pole when salts are electrolyzed, and that oxides cannot be supposed to exist ready formed in salts. Davy, whose electrical experiments were commenced in 1800, undertook in 1806 a course of investigations which led him to the discovery of the metals of the alkalies and alkaline earths. In 1807 he expressed the opinion that bodies having an affinity for one another are in different states of electricity, and “that chemical and electrical attractions depend upon the same cause, acting in the one case on particles and in the other on masses of matter.”
In 1834: Faraday discovered that the decompositions effected by the voltaic current indicate the quantity by weight in which the elements combine, or the weights of the atoms of the atomic theory,. thus adding to the probability of the correctness of the supposition that the operations of the same agent are exhibited in both chemical and electrical phenomena. Latterly, the discovery of the action of the copper-zinc couple by Gladstone and Tribe has opened out a new field of electro-chemical research, aIready productive of important results. To mention, however, the numerous advances that have been made in recent times, whether in chemical physics or in chemistry proper—especially as regards the constitution and synthesis of the compounds of carbon—would be to exceed the scope of the present introductory notice. The reader must, therefore, be referred to the treatises and original memoirs of those whose labours have effected the modern development of chemistry, and have raised it to the high position which it occupies as a science at the present time. (f. h. b.) The Elements.
The examination of all the various substances met with in nature has led to the discovery of sixty-four different bodies, from which it is impossible by any means now at our disposal to separate simpler substances; they are consequently termed “elementary or simple bodies,” or “elements.” It is not asserted that such substances are absolutely simple, or that they may not be found hereafter to yield more than one kind of matter, but merely that so far as our knowledge extends it is so; indeed, recent spectroscopic researches favour the impression that some at least of the elements are, perhaps, compounds of simpler bodies.
It is the study of the laws which regulate the combination of the elements with one another, and to which their compounds are subject in their mutual actions, and of the properties of the elements and of the compounds formed by their union, which constitutes the science of chemistry.
The majority of the elements are of comparatively rare occurrence, whilst many of them are extremely rare, being met with only in certain localities. Gold and a few others almost always occur as such, and our atmosphere, we know, consists chiefly of the two elements, oxygen and nitrogen, in the free state; also some others, such as carbon, copper, silver, and sulphur, are occasionally met with in the uncombined state, but in general two or more elements are found united. The number of elements of which the materials are made up which principally compose the earth and the plants and animals living upon it is, however, exceedingly small. Thus, plants consist chiefly of carbon, hydrogen, and oxygen, and animals chiefly of these three elements together with nitrogen; water consists entirely of oxygen and hydrogen. The solid earth is mainly composed of substances such as quartz or silica, chalk or limestone, and various silicates like felspar and clay,—the elements which are the chief constituents of these substances being oxygen, silicon, carbon, calcium, magnesium, aluminium, iron, and potassium. The only elements which appear to be absolutely essential to vegetable or animal life are carbon, oxygen, hydrogen, nitrogen, sulphur, phosphorus, calcium, iron, potassium, sodium, magnesium, silicon, and chlorine; and the remainder are to us, so to speak, merely of artificial value.
Four of the elements—chlorine, hydrogen, nitrogen, and oxygen—are gases; and fluorine, which is not known with certainty in the free state, is probably also a gas; two are liquid at ordinary atmospheric temperatures, viz., mercury and bromine; and the element gallium, recently discovered in certain zinc ores, but as yet obtained in only very small quantities, is also said to be a liquid. The remaining elements are solids.
The following is a list of the elements now known, arranged in alphabetical order. The most important elements are distinguished in the table by capitals; whilst those which at present are of slight importance, on account of their rare occurrence, or of our insufficient knowledge of their properties, are printed in italics. [table]
Laws of Combination by Weight and Volume — Their explanation by the Atomic Theory — Determination of the relative Weights of Atoms.
The proportions in which elements unite together are definite and constant, a given compound always consisting of the same elements united in the same proportions.
Chloride of silver, for example, in whatever manner it may be prepared, invariably consists of chlorine and silver in the proportions by weight of 35’36 parts of the former and 107’66 of the latter.
But it is often the case that elements combine together in several proportions; whenever this occurs the several proportions in which the one element unites with the other invariably bear a simple relation to one another. Thus, 1 part by weight of hydrogen unites with 8 parts by weight of oxygen, forming water, and with 16 or 8 × 2 parts of oxygen, forming peroxide of hydrogen. Again, in nitrous oxide we have a compound of 8 parts by weight of oxygen and 14 of nitrogen; in nitric oxide a compound of 16 or 8 × 2 parts of oxygen and 14 of nitrogen; in nitrous anhydride a compound of 24 or 8 × 3 parts of oxygen and 14 of nitrogen; in nitric peroxide a compound of 32 or 8 × 4 parts of oxygen and 14 of nitrogen; and lastly, in nitric anhydride a compound of 40 or 8 × 5 parts of oxygen and 14 of nitrogen. This law is known as the law of combination in multiple proportions.
The proportions in which two elements combine with a third also represent the proportions in which, or in some simple multiple of which, they will themselves combine. For instance, 35 ’ 36 parts of chlorine and 79’75 parts of bromine combine with 107’66 parts of silver; and when chlorine and bromine unite it is in the proportion of 35 36 parts of the former to 79’75 parts of the latter. Iodine unites with silver in the proportion of 126’53 parts to 107’66 parts of the latter, but it combines with chlorine in two proportions, viz., in the proportion of 126’53 parts either to 35’36 or to three times 35’36 parts of chlorine. This is known as the law of combination in reciprocal proportions.
In explanation of these three laws deduced entirely from experimental observations, chemists have adopted the atomic or molecular theory which was first introduced into the science by Dalton at the commencement of this century. According to this theory the exceedingly small masses or molecules of which it is supposed matter consists are composite, being made up of indivisible particles or atoms (see the article Atom, vol. iii. p. 36). The molecules of the elements are assumed to consist of similar atoms, whereas those of compounds are congeries of dissimilar atonis; and the molecules which constitute a given kind of matter, it is supposed, are alike in weight and general properties, but differ from those of which all other kinds of matter are composed, so that every molecule belongs to one of a definite number of species. The study of the alterations which take place in the composition of molecules under the influence of various forces, and which result from their action upon one another, is the work of the chemist; whilst it is the province of the physicist to study the influences of those forces upon matter which affect entire molecules without in any way altering their composition.
The chemist has no means of ascertaining, nor does he attempt to ascertain, the absolute weights of the atoms or of the molecules of the various elements and their compounds; he concerns himself merely with their relative weights, hydrogen being adopted as the standard of reference since it is the lightest of all known elements. The relative weight of the atoms of the various elements referred to that of hydrogen regarded as 1 are given in the third column of the table on page 467. The determination of the exact atomic weight of an element is an operation of extreme difficulty, and one requiring the greatest analytical skill, so that as yet the atomic weights of only a limited number of elements have been ascertained with more than approximate accuracy. The most accurately determined atomic weights are those of hydrogen, oxygen, nitrogen, chlorine, bromine, iodine, lithium, potassium, sodium, silver, and thallium. Apparently the numbers obtained for these elements are practically perfect.
The manner in which the relative weights of the atoms of the elements are determined will be evident from the following considerations.
If, instead of comparing together the relative weights of the elements which enter into combination, the volumes which they occupy in the state of gas (at the same temperature and under the same pressure) before and after combination are compared, it is found that gases always unite together in very simple proportions, viz., either in equal volumes, or in volumes which bear some simple relation such as 1: 2, 1: 3, 1: 4, 2: 3, <fcc. Moreover, whatever the number of volumes before combination, it always is reduced to two on combination. Thus, equal volumes of hydrogen and chlorine gases unite without condensation to form hydrochloric acid gas; in the production of water 2 volumes of hydrogen and 1 of oxygen combine, but form only 2 volumes of water-gas or steam; and if ammonia gas be decomposed by heat or a series of electric sparks, 2 volumes of the gas yield 3 volumes of hydrogen and 1 of nitrogen.
Now, according to the law of Boyle and Mariotte, the volume of a given mass of any gas varies inversely as the pressure, provided that the temperature remains the same; for instance, the quantity of air which is contained in a vessel of the capacity of 1 pint under the pressure of 1 atmosphere, or 15 lb upon the square inch, may be contained in a vessel of half a pint capacity if the pressure be doubled.
According to the law of Charles and Gay-Lussac, on the other hand, all gases expand equally by heat, provided the pressure remains constant,—the rate of expansion being τ ∣ 7 of the volume at 0° C. for each rise of l 0 C. in temperature; or in other words, the volume of a gas varies directly as the absolute temperature.
A gas which strictly conforms to these two laws is said to be a perfect gas, but none of the gases with which we are acquainted are perfect in this sense. Thus, Andrews’s experiments show that carbonic anhydride, which under a pressure of 36 atmospheres at 0° C. is reduced to the liquid state, condenses more than it should according to Boyle’s law. Again, the density of chlorine gas referred to air, according to Stas’s determination of the atomic weight of this element, should be 2’4501. The following table exhibits its density at various temperatures from 20° to 200° C., and it is evident that it is higher than it should be at all temperatures below 200° C. :— [table]
From the few accurate observations which have been made on this subject it appears that, in general, the departure from the laws of Boyle and Charles is greater the more the temperature of the gas approaches to that at which it becomes liquid; and chlorine affords an instructive illustration of this, since it is readily condensed to a liquid under the pressure of 4 atmospheres at 15 o, 5 C., or by cooling in a bath of solid carbonic anhydride and ether.
The general resemblance in the behaviour of gases under the influence of pressure and heat is very great, however, although not in absolute accordance with the laws of Boyle and Charles; by this we are led to the assumption that their physical constitution must be similar, and, therefore, to the acceptance of the proposition, originally stated by Avogadro in 1811, that equal volumes of different gases contain equal numbers of molecules. Obviously, therefore, if the relative weights of equal volumes of different gases are determined under the same conditions as to temperature and pressure, the temperature and pressure chosen being that at which the gases most closely approximate to the requirements of the laws above stated, an estimate of the relative weights of their molecules is obtained. For example, the density of nitrogen referred to hydrogen is 14, since a given volume of nitrogen is found to weigh 14 times as much as an equal volume of hydrogen at the same temperature and under the same pressure; hence, according to Avogadro’s hypothesis, the molecules of nitrogen are fourteen times as heavy as the hydrogen molecules.
In the formation of hydrochloric acid gas equal volumes of chlorine and hydrogen unite without condensation. The density of chlorine gas referred to hydrogen is 35'36, and the simplest possible hypothesis of the composition of hydrochloric acid is that it consists of an atom of hydrogen weighing 1, and an atom of chlorine weighing 35'36, so that its molecule, therefore, must weigh 33'36. But since the density of hydrochloric acid gas is ascertained by experiment to be only 18*18 as compared with that of hydrogen, and, according to Avogadro’s hypothesis, equal volumes of hydrogen and hydrochloric acid gas contain equal numbers of molecules, it follows that the weight of the hydrogen molecule as compared with that of hydrochloric acid must be 2, or in other words, that the hydrogen molecule consists of two atoms. The chlorine molecule in like manner must consist of two atoms, each weighing 35'36, and in the formation of hydrochloric acid from hydrogen and chlorine two molecules of hydrochloric acid are produced from a molecule of hydrogen and a molecule of chlorine: in the one molecule half the hydrogen is displaced by chlorine, in the other half the chlorine is displaced by hydrogen. It may be proved that the assumption is correct that the molecule of hydrochloric acid contains only a single atom of chlorine, weighing 35'36, and that it does not consist, for example, of two atoms of chlorine each weighing 17'68, by comparing the various volatile compounds containing chlorine. In the first place their densities in the state of gas are determined, and a knowledge is thus obtained of the relative weights of their molecules as compared with that of the hydrogen molecule; the percentage of chlorine they contain is then ascertained by careful analysis. The density referred to hydrogen as unity multiplied by 2 gives the molecular weight of the compound; and the percentage of chlorine being known, the amount contained in the quantity expressed by the molecular weight is ascertained by a simple calculation. For example, the density of sulphur chloride is found to be 57 - 36, and its molecular weight is therefore 57'36 × 2 or 114'72; it contains 61'64 per cent, of chlorine, so that in 114'72 parts there are 70'72 of chlorine When the numbers thus deduced are compared it is seen that the lowest amongst them is 35'36, and that all the higher numbers are simple multiples of this; 35'36 is accordingly adopted as the number which expresses the weight of the atom of chlorine relatively to that of the hydrogen atom. A number of volatile chlorine compounds are thus compared in the following table :— [table]
In like manner, on comparing the various volatile compounds which contain oxygen, it is found that the number 16 represents the least weight of oxygen contained in the molecular weight of any of its compounds; 16 is therefore taken as the atomic weight of oxygen.
In all cases in which it is possible to obtain volatile compounds, the atomic weights of elements may be deduced in this manner; unfortunately, however, many of the elements do not furnish stable volatile compounds, so that hitherto the atomic weights of the following elements only have been ascertained by the application of Avogadro’s hypothesis :— [table]
The determination of the density of bodies in the state of gas is thus an operation of fundamental importance. The precise manner in which the determination is effected is described in most of the text-books on chemistry. The methods ordinarily employed in the case of liquids and solids which by the application of heat can be converted into vapour or gas without undergoing decomposition are known respectively as Dumas’s and Gay-Lussac’s, a modification of the latter method of great value has recently been introduced by Hofmann. By Dumas’s method the weight of substance is ascertained which will furnish a certain volume of gas at a certain temperature and pressure; by Gay-Lussac’s method, however, and by Hofmann’s modification of it, the volume of gas is measured which is furnished by a given weight of the substance at a certain temperature and pressure. By either method we arrive finally at a knowledge of the weight (w) of a certain volume (f) of the gas at a temperature t and pressure p ; and its density (D) referred to hydrogen Is then found by dividing the weight w by the weight (w) of an equal volume of hydrogen at the same temperature t and pressure p —
D = ^ w ’
Or the density referred to air may be calculated in a similar manner, and then converted into the density referred to hydrogen by multiplication by 14'43,— the number which expresses the density of air referred to hydrogen. Both methods require that the substance be heated to the temperature at which its vapour most closely approximates to the laws of Boyle and Charles, which is readily ascertained by experiment; this temperature, however, is often very considerably above the boiling point of the substance, and acetic acid may be cited as an illustration of this. Thus, although this acid boils at 119° C., its vapour does not exhibit the required density until it is heated to 250° C., as will be evident from the following table—the theoretical vapour density of acetic acid vapour referred to hydrogen being about 30 :— [table]
Owing to unavoidable experimental errors, and, in many cases, probably to the circumstance that the vapours of solid and liquid bodies are very imperfect gases at temperatures not much above their boiling points, the determination of the vapour density of a substance does not, as a rule, furnish a result of more than approximate accuracy,—the result being the more accurate, however, the more rarefied the vapour and the higher the temperature at which the density is determined, provided always that the temperature be not so high as to cause decomposition. But the relative proportions in which the elements combine, and the composition of compounds may usually be ascertained with very great—in many cases with almost absolute—accuracy by chemical analysis; and the determination of the density in the state of gas simply serves to prove which of the several multiple proportions, in which it is found that the elements combine together, is the true atomic weight. For instance, analysis shows that marsh gas is a compound of carbon and hydrogen exactly in the proportions of 1 part of the latter with 3 of the former; that carbonic anhydride consists of carbon and oxygen in the proportion of 16 parts or an atom of the latter and 6 or 3 × 2 parts of carbon; and carbonic oxide of the same elements, in the proportion, however, of 16 of oxygen and 12 or 3 × 4 of carbon. The question, therefore, is, Which of the numbers, 3, 6, or 12, represents the relative weight of the carbon atom; that is to say, is marsh gas a compound of an atom of hydrogen and an atom of carbon weighing 3, of two atoms of hydrogen and an atom of carbon weighing 6, or of four atoms of hydrogen and an atom of carbon weighing 12 î The molecular weights of three such compounds would be respectively 3 + 1 or 4, 6 + 2 or 8, and 12 + 4 or 16; and the corresponding theoretical densities referred to hydrogen 2, 4, and 8. Finding, however, by experiment that the density of marsh gas is, say, 7'5, we at once conclude that the atomic weight of carbon is 12 and not 6 or 3, because the observed density of marsh gas most closely accords with that required on this assumption. The difference between theory and experiment is practically seldom, if ever, so large as in this case, which is merely given as an illustration of the principle involved.
The equivalent of an element—that is to say, the amount of it which is capable of combining with or displacing one part by weight, or one atom, of hydrogen—being known, its atomic weight is not absolutely fixed by the determination of the density in the state of gas of its compounds with other elements; we are at most enabled to say from this that the atomic weight cannot exceed a certain value, —for instance, that the atomic weight of chlorine cannot exceed 35 , 36, because all its compounds contain either this amount or some simple multiple of it in their molecules. It is nevertheless possible that 35 ’ 36 is not the weight of one but of several atoms of chlorine; the probability that 35 ’ 36 is the true atomic weight is enormously increased, however, as compound after compound is examined and found to contain 35 ’ 36 or some simple multiple of 35 36 parts of chlorine in its molecule.
In the case of those elements of which stable volatile compounds have not been obtained, the study of their specific heats is of great importance, and moreover furnishes most important confirmation of the atomic weights deduced by the aid of Avogadro’s hypothesis. To raise the temperature of equal weights of different substances the same number of degrees, from 0° to 1° C. for instance, very different amounts of heat are required; and on the other hand very different amounts of heat are given out when equal weights of different substances are cooled from the initial temperature t° to a lower temperature t' ο . Of all bodies except hydrogen water has the greatest capacity for heat, and is, therefore, adopted as the standard of reference,—the number which expresses the amount of heat necessary to raise the temperature of a given weight of a body a certain number of degrees, or which is given out by it in cooling through a certain number of degrees, as compared with that required to raise the temperature of an equal weight of water the same number of degrees, being termed its specific heat. Thus, the specific heat of lithium is ’ 940b; that is to say, to raise the temperature of a given weight of lithium—1 gramme, for example—from 0° to 1° C. only requires - 9408 of the heat necessary to raise the temperature of 1 gramme of water from 0° to l o C. The specific heats in the solid state of the various elements of which the atomic weights have been determined by Avogadro’s hypothesis are given in the second column of the following table :— [table]
On comparing the numbers in the fourth column of this table it will be seen that they vary within comparatively narrow limits; and if certain of the elements are excepted, viz., boron, carbon, phosphorus, sulphur, silicon, and selenium, the agreement becomes much closer, the average product obtained by multiplying specific heat into atomic weight being about 6 ’ 3. From this it would appear that the specific heats of these elements are, at least approximately, inversely proportional to their atomic weights. From the observation of this relation in the case of only a small number of elements Dulong and Petit, in 1811, were led to infer that the atoms of all simple bodies have the same capacity for heat. The specific heat of a body varies, however, with the temperature; an extreme instance of this is afforded by the elements carbon, boron, and silicon, as will be evident on inspecting the following table of the specific heat of carbon in the form of diamond at various temperatures :— [table]
Hence, owing to the circumstance that the determinations of specific heat have not been made at temperatures which are comparable for the different elements, there is no doubt that many of the results which have been obtained are defective; but from Weber’s recent researches it appears that in the case of the solid elements there is a point for each element, after which the increase in specific heat with increase of temperature is insignificant, and when this point is reached the product of specific heat into atomic weight—the so-called atomic heat—varies within comparatively narrow limits. These limits, according to Weber, are from 5’5 to 6 ’ 5, but it appears probable that the superior limit is slightly greater than this, and as a matter of fact the atomic heats of nearly all the elements are nearer 6 ’ 5 than 5 ’ 5,—the latter number being characteristic of the so-called non-metallic elements, the atomic weights of which can be determined by the aid of Avogadro’s hypothesis.
In consequence of this relation between the specific heat of an element and its atomic weight, we can readily determine the atomic weights of elements which do not furnish volatile compounds, it being merely necessary to ascertain the equivalent of the element by analyzing its compounds, and to multiply the equivalent by such a number that the product when multiplied by the specific heat of the element will furnish a number not less than 5 ’ 5, and not much greater than 6 - 5. As an instance of the application of the method, the metal indium, one of the most recently discovered elements, may be cited. Analysis showed that the equivalent of indium was 37'8, but this number was doubled on account of the many points in which indium resembles zinc,—the latter metal having the equivalent 32 ’ 45, but the atomic weight 64 ’ 9. For some time, until Bunsen determined its specific heat, 75’6 was accepted as the atomic weight of indium, but he found that it was necessary to multiply the equivalent by 3 in order that the product of specific heat into atomic weight should correspond to that ordinarily observed, thus raising the atomic weight to 113 ’ 4.
The following table is a list of the elements whose atomic weights have been determined from the observation of their specific heats alone :— [table]
The only elements of which at present the atomic weights have not been ascertained, either by the vapourdensity or specific heat method, are erbium, thorium, uranium, and yttrium. Even the equivalent of gallium is not yet known.
Often, also, information which is of service in determining the atomic weights of elements may be obtained by observing the manner in which their compounds crystallize, for it is found that in many instances a given element in a compound may be displaced by another without altering the crystalline form; for example, in ordinary alum, which is a double sulphate of potassium and aluminium, the potassium may readily be displaced by sodium, or the aluminium by chromium, and yet the resulting compounds crystallize in precisely the same manner as the alum. Substances which thus agree in crystalline form , are said to be isomorphous. A great number of isomorphous substances have been examined by chemists; and they have been led to infer that, as a rule, when two compounds containing similar elements agree in crystalline form they contain the same number of atoms; and hence the quantity of an element which is capable of displacing an atom of another element in a compound without altering its crystalline form is regarded as its atomic weight. This is by no means always the case, however, and the occurrence of isomorphism cannot alone be taken as a proof of the atomic weight of an element; it is of service rather as a check on the determinations made in other ways, and as a stimulus to investigation in the case of elements which have not been sufficiently examined. For instance, certain phosphorous and arsenic minerals of similar composition were known to be isomorphous with a vanadium mineral, vanadinite, the composition of which, according to the received atomic weight of vanadium, exhibited no analogy with them. Roscoe was led by this to investigate anew the compounds of vanadium and to determine its atomic weight, and he found that what had hitherto been regarded as vanadium was really an oxide of the element, and that when the composition of the mineral vanadinite was calculated from the true atomic weight of vanadium, it was precisely similar to that of the minerals with which it was known to be isomorphous.
Molecular Weights.
We have pointed out (p. 469) that it is to be supposed that the molecule of hydrogen consists of two atoms; hence, if the atomic weight of hydrogen be taken as .1 its molecular weight is 2. In order to ascertain the molecular weights of other elements—that is to say, the relative weights of their molecules referred to that of hydrogen— it is merely necessary to determine their densities referred to hydrogen as unity, and then to multiply the densities by 2. Unfortunately, owing to the high temperature at which most of the elements are converted into vapour, the densities of only very few of them are as yet known.
When, however, the molecular weights of the elements are compared with their atomic weights it is found that they are not always, as in the case of hydrogen, double their atomic weights; hence it is inferred that the molecules of elements do not all contain two atoms. In a few cases the atomic weight and molecular weight agree, which necessitates the conclusion that the molecules are monatomic or consist of a single atom ; in a few other cases the molecular weight is either 4 or 6 times the atomic weight, and the molecules are therefore regarded as tetratomic or hexatomic.
The following table includes all the elements of which the molecular weights have been determined :— [table]
It will be seen that two numbers are given for sulphur This is because at temperatures above 80.0° C. the density of sulphur vapour is such as to indicate that the sulphur molecule consists of 2 atoms, whereas its density at about 500° C. is three times as great, and, consequently, it is to be supposed that the molecules are hexatomic. Selenium, which is closely allied to sulphur, exhibits a very similar behaviour, its vapour at about 1400° C. containing only diatomic molecules, but as the temperature sinks its density rapidly increases. These two elements, in fact, afford striking examples of the resolution of complex molecules into simpler molecules by heat; and it is by no means improbable that the tetratomic phosphorus and arsenic molecules will be found to behave similarly if sufficiently heated.
No method is known by which it is possible to determine the molecular weights of elements or compounds in any other state than that of gas, and the behaviour of sulphur is alone sufficient to prove that we are not justified in assuming that the molecular weights of liquid or solid elements or compounds are identical with their molecular weights as gases. Indeed it is in the highest degree probable that the molecules of solid and liquid bodies are very frequently far more complex than the molecules of the same bodies in the state of gas.
Chemical Notation — Valency of Elements — Rational Formulae.
Opposite the name of each element in the second column of the table on p. 467, the symbol is given which is always employed to represent it. This symbol, however, not only represents the particular element, but a certain definite quantity of it. Thus, the letter H always stands for 1 atom or 1 part by weight of hydrogen, the letter N for I atom or 14 parts of nitrogen, and the symbol Cl for 1 atom or 35 ’ 36 parts of chlorine. Compounds are in like manner represented by writing the symbols of their constituent elements side by side, and if more than one atom of each element be present, the number is indicated by a numeral placed on the right of the symbol of the element either below or above the line. Thus, hydrochloric acid is represented by the formula HC1, that is to say, it is a compound of an atom of hydrogen with an atom of chlorine, or of 1 part by weight of hydrogen with 3536 parts by weight of chlorine; again, sulphuric acid is represented by the formula H 2 SO 4 , which is a statement that it consists of 2 atoms of hydrogen, 1 of sulphur, and 4 of oxygen, and consequently of certain relative weights of these elements. A figure placed on the right of a symbol only affects the symbol to which it is attached, but when figures are placed in front of several symbols all are affected by it, thus 2H 2 SO 4 means H 2 SO 4 taken twice.
The distribution of weight in chemical change is readily expressed in the form of equations by the aid of these symbols; the equation
2HC1 + Zn = ZnC ’ ∣ 2 + H 2 , for example, is to be read as meaning that from certain definite weights of hydrochloric acid and zinc certain definite weights of two different bodies, zinc chloride and hydrogen, are produced. The + sign is invariably employed in this way either to express combination or action upon, the meaning usually attached to the use of the sign = being that from such and such bodies such and such other bodies are formed.
Usually, when the symbols of the elements are written or printed with a figure to the right, it is understood that this indicates a molecule of the element, the symbol alone representing an atom. Thus, the symbols H 2 and P 4 indicate that the molecules of hydrogen and phosphorus respectively contain 2 and 4 atoms. Since, according to the molecular theory, in all cases of chemical change the action is between molecules, such symbols as these ought always to be employed. Thus, the formation of hydrochloric acid from hydrogen and chlorine is correctly represented by the equation
H 2 + Cl 2 = 2HCb that is to say, a molecule of hydrogen and a molecule of chlorine give rise to two molecules of hydrochloric acid; whilst the following equation merely represents the relative weights of the elements which enter into reaction, and is not a complete expression of what is supposed to take place :—
H + C1 = HC1
As the molecular weights of comparatively few of the elements have been determined, however, it is possible only in a limited number of cases to employ such symbols. The molecular weights of the larger number of compounds are also unknown, but in all cases it is usual to represent them by formulae which to the best of our knowledge express their molecular composition in the state of gas, and not merely the relative number of atoms which they contain; thus, acetic acid consists of carbon, hydrogen, and oxygen in the proportion of one atom of carbon, two of hydrogen, and one of oxygen, but the determination of its vapour density shows that it has a molecular weight corresponding to the formula C 2 H 4 O 2 , which therefore is always employed to represent acetic acid. When chemical change is expressed with the aid of molecular formulae, not only is the distribution of weight represented, but from mere inspection of the symbols it is possible to deduce the relative volumes which the agentsand resultants occupy in the state of gas if measured at the same temperature and under the same pressure. Thus, the equation
2H 2 + O 2 = 2H 2 O not only represents that certain definite weights of hydrogen and oxygen furnish a certain definite weight of the compound which we term water, but that if the water in the state of gas, the hydrogen, and the oxygen are all measured at the same temperature and pressure, the volume occupied by the oxygen is only half that occupied by the hydrogen, whilst the resulting water-gas will only occupy the same volume as the hydrogen. In other words, 2 volumes of oxygen and 4 volumes of hydrogen furnish 4 volumes of water gas. A simple equation like this, therefore, when properly interpreted, affords a large amount of information. One other instance may be given; the equation
2NH 3 = N 2 + 3H 2 , represents the decomposition of ammonia gas into nitrogen and hydrogen gases by the electric spark, and it not only conveys the information that a certain relative weight of ammonia consisting of certain relative weights of hydrogen and nitrogen is broken up into certain relative weights of hydrogen and nitrogen, but also that the nitrogen will be contained in half the space which contained the ammonia, and that the volume of the hydrogen will be one and a half times as great as that of the original ammonia, so that in the decomposition of ammonia the volume becomes doubled.
Formulae which merely express the relative number of atoms of the different elements present in a compound are termed empirical formulae, and the formulae of all compounds whose molecular weights are undetermined are necessarily empirical. The molecular formula of a compound, however, is always a simple multiple of the empirical formula, if not identical with it; thus, the empirical formula of acetic acid is CH 2 O, and its molecular formula is C 2 H 4 O 2 , or twice CH 2 O. If the vapour density of a compound cannot be ascertained, its molecular formula can only be ascertained, with .more or less approximation to truth, by considering its properties, and especially its relations to allied compounds of known molecular weight. For example, chromic anhydride is represented as CrO 3 , and, although it cannot be vaporized, this is held to be.its molecular formula chiefly on account of the formation from chromic anhydride of a volatile chromium oxychloride, the molecular weight of which is found to accord with the formula CrO 2 Cl 2 .
But in addition *o empirical and molecular formulae, chemists are in the habit of employing various kinds of rational formulae, called constitutional formulae, graphic formulae, &c., which not only express the molecular composition of the compounds to which they apply, but also embody certain assumptions as to the manner in which the constituent atoms are arranged, and convey more or less information with regard to the nature of the compound itself, viz., the class to which it belongs, the manner in which it is formed, and the behaviour it will exhibit under various circumstances. Before explaining these formulae it will be necessary, however, to consider the differences in combining power exhibited by the various elements.
It is found that the number of atoms of a given element, of chlorine, for example, which unite with an atom of each of the other elements is very variable. Thus, hydrogen unites with but a single atom of chlorine, zinc with two, boron with three, silicon with four, and phosphorus with five. Those elements which are equivalent in combining or displacing power to a single atom of hydrogen are said to be univalent or monad elements; whilst those which are equivalent to two atoms of hydrogen are termed bivalent or dyad elements; and those equivalent to three, four, five, or six atoms of hydrogen triad, tetrad, pentad, or hexad elements. But not only is the combining power or valency (atomicity) of the elements different, it is also observed that one element may combine with another in several proportions, or that its valency may vary; for example, phosphorus forms two chlorides represented by the formulae PCl 3 and PC1 5 , and nitrogen the series of oxides represented by the formulae
N 2 O, NO, N 2 O 3 , N 2 O 4 , N 2 O 5 .
In explanation of these facts it is supposed that each element has a certain number of “units of affinity,” which may be entirely, or only in part, engaged when it enters into combination with other elements; and in those cases in which the entire number of units of affinity are not engaged by other elements, it is supposed that those which are thus disengaged neutralize each other, as it were. For example, in pentachloride of phosphorus the five units of affinity possessed by the phosphorus atom are satisfied by the five monad atoms of chlorine, but in the trichloride two are disengaged, and, it may be supposed, satisfy each other. Compounds in which all the units of affinity of the contained elements are engaged are said to be saturated, whilst those in which the affinities of the contained elements are not all engaged by other elements are said to be unsaturated. According to this view, it is necessary to assume that, in all unsaturated compounds, two, or some even number of affinities are disengaged; and also that all elements which combine with an even number of monad atoms cannot combine with an odd number, and vice versa, —in other words, that the number of units of affinity active in the case of any given element must be always either an even or an odd number, and that it cannot be at one time an even and at another an odd number. There are, however, a few remarkable exceptions to this “law.” Thus, it must be supposed that in nitric oxide, NO, an odd number of affinities are disengaged, since a single atom of dyad oxygen is united with a single atom of nitrogen, which in all its compounds with other elements acts either as a triad or pentad. When nitric peroxide, N 2 O 4 , is converted into gas, it decomposes, and at about 180° C. its vapour entirely consists of molecules of the composition NO 2 ; while at temperatures between this and 0° C. it consists of a mixture in different proportions of the two kinds of molecules, N 2 O 4 and NO 2 . The oxide NO 2 must be regarded as another instance of a compound in which an odd number of affinities of one of the contained elements are disengaged, since it contains two atoms of dyad oxygen united with a single atom of triad or pentad nitrogen. Again, when hexachloride of tungsten is converted into vapour it is decomposed into chlorine and a pentachloride, having a normal vapour density, but as in the majority of its compounds tungsten acts as a hexad, we apparently must regard its pentachloride as a compound in which an odd number of free affinities are disengaged. Hitherto no explanation has been given of these exceptions to what appears to be a law of almost universal application, viz., that the sum of the units of affinity of all the atoms in a compound is an even number.
The number of units of affinity active in the case of any particular element is largely dependent, however, upon the nature of the element or elements with which it is associated. Thus, an atom of iodine only combines with one of hydrogen, but may unite with three of chlorine, which never combines with more than a single atom of hydrogen; an atom of phosphorus unites with only three atoms of hydrogen, but with five of chlorine, or with four of hydrogen and one of iodine; and the chlorides corresponding to the higher oxides of lead, nickel, manganese, and arsenic, PbO 2 , Ni 2 O 3 , MnO 2 , and As 2 O 5 do not exist as stable compounds, but only the lower chlorides, PbCl 2 , NiCl 2 , MnCl 2 , and AsC1 3 , are known.
It is difficult, therefore, to classify the elements according to their valencies; indeed, an absolute classification is scarcely possible. In the following table a number of the elements are arranged mostly according to their apparent maximum valencies :— [table]
The valency of an element is usually expressed by dashes or Roman numerals placed on the right of its symbol, thus: H', 0", B"', C ,τ , P τ , Mo’ 1 ; but in constructing graphic formulae the symbols of the elements are written with as many lines attached to each symbol as the element which it represents has units of affinity.
Graphic formulae are employed to express the manner in which it is assumed that the constituent atoms of compounds are associated together; for example, the trioxidβ of sulphur is usually regarded as a compound of an atom of hexad sulphur with three atoms of dyad oxygen, and this hypothesis is illustrated by the graphic formula
O=S=O
II ο
When this oxide is brought into contact with water it combines with it forming sulphuric acid, H 2 SO 4 , and it is supposed that in this compound only two of the oxygen atoms are wholly associated with the sulphur atom, each of the remaining oxygen atoms being united by one of its affinities to the sulphur atom, and by the remaining affinity to an atom of hydrogen; thus—
0
II
H—O—S-O—H
II ο
Again, the reactions of acetic acid, C 2 H 4 O 2 , show that the four atoms of hydrogen which it contains have not all the same function, and also that the two atoms of oxygen have different functions; the graphic formula which we are led to assign to acetic acid, viz.,—
H
I
H—C—H
I
O=C—0—H
serves in a measure to express this, three of the atoms of hydrogen being represented as associated with one of the atoms of carbon, whilst the fourth atom is associated with an atom of oxygen which is united by a single affinity to the second atom of carbon, to which, however, the second atom of oxygen is united by both of its affinities. It is not to be supposed that there are any actual bonds of union between the atoms; graphic formulae such as these merely express the hypothesis that certain of the atoms in a compound come directly within the sphere of attraction of certain other atoms, and only indirectly within the sphere of attraction of others,—an hypothesis to which chemists are led by observing that it is often possible to separate a group of elements from a compound, and to displace it by other elements or groups of elements.
Rational formulae of a much simpler description than these graphic formulae are generally employed. For instance, sulphuric acid is usually represented by the formula SO 2 (OH) 2 , which indicates that it may be regarded as a compound of the group SO 2 with twice the group OH. Each of these OH groups is equivalent in combining or displacing power to a monad element, since it consists of an atom of dyad oxygen associated with a single atom of monad hydrogen, so that in this case the SO 2 group is equivalent to an atom of a dyad element. This formula for sulphuric acid, however, merely represents such facts as that it is possible to displace an atom of hydrogen and an atom of oxygen in sulphuric acid by a single atom of chlorine, thus forming the compound SO 3 HC1; and that by the action of water on the compound SO 2 Cl 2 twice the group OH, or water minus an atom of hydrogen, is introduced in place of the two monad atoms of chlorine—
SO 2 C1 2 + 2HOH = SO 2 (OH) 2 + 2HC1.
Water. Sulphuric acid.
Constitutional formulae like these, in fact, are nothing more than symbolic expressions of the character of the compounds which they represent, the arrangement of symbols in a certain definite manner being understood to convey certain information with regard to the compounds represented.
Groups of two or more atoms like SO 2 and OH, which are capable of playing the part of elementary atoms (that is to say, which can be transferred from compound to compound), are termed compound radicles, the elementary atoms being simple radicles. Thus, the atom of hydrogen is a monad simple radicle, the atom of oxygen a dyad pimple radicle, whilst the group OH is a monad compound radicle.
It is often convenient to regard compounds as formed upon certain types; alcohol, for example, may be said to be a compound formed upon the water type, that is to say, a compound formed from water by displacing one of the atoms of hydrogen by the group of elements C 2 H 6 , thus 0 í H 0 I H 115 ( H (H
Water Alcohol.
Chemwcd Action.
Chemical change or chemical action may be said to take place whenever changes occur which involve an alteration in the composition of molecules, and may be the result of the action of agents such as heat, electricity, or light, or of two or more elements or compounds upon each other.
Three kinds of changes are to be distinguished, viz., changes which involve combination, changes which involve decomposition or separation, and changes which involve at the same time both decomposition and combination. Changes of the first and second kind, according to our present views of the constitution of molecules, are probably of very rare occurrence; in fact, chemical action appears almost always to involve the occurrence of both these kinds of change, for, as already pointed out, we must assume that the molecules of hydrogen, oxygen, and several other elements are diatomic, or that they consist of two atoms. Indeed, it appears probable that with few exceptions the elements are all compounds of similar atoms united together by one or more units of affinity, according to their valencies. If this be the case, however, it is evident that there is no real distinction between the reactions which take place when two elements combine together and when an element in a compound is displaced by another. The combination, as it is ordinarily termed, of chlorine with hydrogen, and the displacement of iodine in potassium iodide by the action of chlorine, may be cited as examples; if these reactions are represented, as such reactions very commonly are, by equations which merely express the relative weights of the bodies which enter into reaction, and of the products, thus—
H + Cl = HC1
Hydrogen, chlorine. Hydrochloric acid.
KI + Cl = KC1 + I
Potassium iodide. Chlorine. Potassium chloride. Iodine.
they appear to differ in character; but if they are correctly represented by molecular equations, or equations which express the relative .number of molecules which enter into reaction and which result from the reaction, it will be obvious that the character of the reaction is substantially the same in both cases, and that both are instances of the occurrence of what is ordinarily termed double decomposition—
H 2 + Cl 2 = 2HC1 Hydrogen. Chlorine. Hydrochloric acid.
2KI + Cl 2 = 2KC1 + I 2 .
Potassic iodide. Chlorine. Potassic chloride. Iodine.
For chemical action to take place between two bodies it is necessary that they should be in contact, and, therefore, generally speaking, that one of them should be in the state of liquid or gas.
In all cases of chemical change energy in the form of heat is either developed or absorbed, and the amount of heat developed or absorbed in a given reaction is as definite as are the weights of the substance engaged in the reaction. Thus, in the production of hydrochloric acid from hydrogen and chlorine 22,000 units of heat are developed; in the production of hydrobromic acid from hydrogen and bromine, however, only 8440 units of heat are developed; and in the formation of hydriodic acid from hydrogen and iodine 6040 units of heat are absorbed.
This difference in behaviour of the three elements, chlorine, bromine, and iodine, which in many respects exhibit considerable resemblance, may be explained in the following manner. We may suppose that in the formation of gaseous hydrochloric acid from gaseous chlorine and hydrogen, according to the equation
H 2 + C1 2 = HC1 + HC1, a certain amount of energy is expended in separating the atoms of hydrogen in the hydrogen molecule, and the atoms of chlorine in the chlorine molecule, from each other; but that heat is developed by the combination of the hydrogen atoms with the chlorine atoms, and that, as more energy is developed by the union of the atoms of hydrogen and chlorine than is expended in separating the hydrogen atoms from each other and the chlorine atoms from one another, the result of the action of the two elements upon each other is the development of heat,—the amount finally developed in the reaction being the difference between that absorbed in decomposing the elementary molecules and that developed by the combination of the atoms of chlorine and hydrogen. In the formation of gaseous hydrobromic acid from liquid bromine and gaseous hydrogen-H 2 + Br 2 = HBr+HBr, in addition to the energy expended in decomposing the hydrogen and bromine molecules, energy is also expended in converting the liquid bromine into the gaseous condition, and probably less heat is developed by the combination of bromine and hydrogen than by the combination of chlorine and hydrogen, so that the amount of heat finally developed is much less than is developed in the formation of hydrochloric acid. Lastly, in the production of gaseous hydriodic acid from hydrogen and solid iodine—
H 2 + I 2 = HI + HI, so much energy is expended in the decomposition of the hydrogen and iodine molecules and in the conversion of the iodine into the gaseous condition, that the heat which it may be supposed is developed by the combination of the hydrogen and iodine atoms is insufficient to balance the expenditure, and the final result is therefore negative; hence it is necessary in forming hydriodic acid from its elements to apply heat continuously.
These compounds also afford examples of the fact that, generally speaking, those compounds are most readily formed, and are most stable, in the formation of which the most heat is developed. Thus, chlorine enters into reaction with hydrogen, and removes hydrogen from hydrogenized bodies, far more readily than bromine; and hydrochloric acid is a far more stable substance than hydrobromic acid, hydriodic acid being greatly inferior even to hydrobromic acid in stability.
When two substances which by their action upon each other develop much heat enter into reaction, the reaction is usually complete without the employment of an excess of either; for example, when a mixture of hydrogen and oxygen, in the proportions to form water—
2H 2 + O 2 = 2OH 2 , is exploded, it is entirely converted into water. This is also the case if two substances are brought together in solution, by the action of which upon each other a third body is formed which is insoluble in the solvent employed, and which also does not tend to react upon any of the substances present; for instance, when a solution of a chloride is added to a solution of a silver salt, insoluble silver chloride is precipitated, and almost the whole of the silver is removed from solution, even if the amount of the chloride employed be not in excess of that theoretically required.
But if there be no tendency to form an insoluble compound, or one which is not liable to react upon any of the other substances present, this is no longer the case. For example, when a solution of a per-salt of iron is added to a solution of potassium thiocyanate, a deep red coloration is produced, owing to the formation of thiocyanate of iron. Theoretically the reaction takes place in the case of the pernitrate of iron in the manner represented by the equation Fe 2 (NO 3 ) 6 + 6KCNS = Fe 2 (CN τ S) 6 + 6KNO 3 ,∙ Ferric nitrate. Potassium thiocyanate. Ferric thiocyanate. Potassium nitrate, but it is found that even when more than sixty times the amount of potassium thiocyanate required by this equation is added, a portion of the ferric nitrate still remains unconverted, doubtless owing to the occurrence of the reverse change—
Fe 2 (CNS) 6 + 6KNO 3 = Fe 2 (NO 3 ) 6 + 6KCNS.
In this, as in most other cases in which substances act upon one another under such circumstances that the resulting compounds are free to react, the extent to which the different kinds of action which may occur take place is dependent upon the mass of the substances present in the mixture. As another instance of this kind, the decomposition of bismuth chloride by water may be cited. If a very large quantity of water be added, the chloride is entirely decomposed in the manner represented by the equation—
BiCl 3 + OH 2 = BiOCl + 2HC1,
Bismuth chloride. Bismuth oxychloride.
the oxychloride being precipitated; but if smaller quantities of water be added the decomposition is incomplete, and it is found that the extent to which decomposition takes place is proportional to the quantity of water employed, the decomposition being incomplete, except in presence of large quantities of water, because of the occurrence of the reverse action—
BiOCl + 2HC1 = BiCT 3 + OH 2 .
Chemical change which merely involves simple decomposition is also influenced by the presence of the products of decomposition. Thus when calcium carbonate is strongly heated in an open vessel, it is entirely decomposed into carbon dioxide gas and calcium oxide—
CaCO 3 = CO 2 + CaO. calcium carbonate, carbon dioxide. Calcium oxide.
When it is heated in a confined space, the decomposition only goes on until the liberated gas has attained a certain tension, and as long as the temperature does not vary, the tension remains the same, and is independent of the pro-. portion of the compound decomposed, that is to say, of the amount of calcium oxide present; but the tension increases if the temperature is raised, and diminishes if it is lowered, owing to the recombination of a portion of the carbon dioxide with the calcium oxide; for example, at 800° C. the tension is equivalent to a column of mercury 85 millimetres high, but at1040° C. it is equivalent to a column of 520 mm. Deville applies the term dissociation to changes which occur in this manner; the term only applies to those cases of decomposition in which products are obtained which under the conditions of the experiment are capable of reuniting to form the original substance. When such substances are heated in a confined space, not only does decomposition take place, but a certain proportion of the decomposition products recombine, and for each degree of temperature a definite relation exists between the number of molecules broken up and the number of molecules reconstituted; an equilibrium is thus maintained, and the tension remains constant. But as the temperature rises, a larger number of molecules are decomposed and a proportionally smaller number reformed than at the lower temperature, so that the tension rises with the temperature. When, however, a substance like calcium carbonate is heated in an open vessel, the carbon dioxide escapes as it is formed, and therefore it is impossible for recomposition to occur, and ultimately the whole of the carbonate is decomposed. Apparently whenever solid substances undergo dissociation and furnish a solid and a gaseous product, the tension of dissociation is always independent of the amount of the solid decomposition product present.
The Elements and their Compounds.
The elements are usually divided into two great classes, the metallic and the non-metallic elements, the following being classed as non-metals, and the remainder as metals;— [table]
The metals are mostly bodies of high specific gravity; they exhibit, when polished, a peculiar brilliancy or metallic lustre, and they are good conductors of heat and electricity; whereas the non-metals are mostly bodies of low specific gravity, and bad conductors of heat and electricity, and do not exhibit metallic lustre. The non-metallic elements are also sometimes termed metalloids, but this appellation, which signifies metal-like substances, strictly belongs to certain elements which do not possess the properties of the true metals although they more closely resemble them than the non-metals in many respects; thus, selenium and tellurium, which are closely allied to sulphur in their chemical properties, although bad conductors of heat and electricity, exhibit metallic lustre and have relatively high specific gravities.
But when the properties of the elements are carefully contrasted together it is obvious that no strict line of demarcation can be drawn dividing them into two classes; and if they are arranged in a series, those which are most closely allied in properties being placed next to each other, it is evident that there is a more or less regular alteration in properties from term to term in the series.
When binary compounds, or compounds of two elements, are decomposed by an electric current,the two elements make their appearance at opposite poles. Those elements which are disengaged at the negative pole are termed electropositive, or positive, or basylous elements, whilst those disengaged at the positive pole are termed electro-negative, or negative, or chlorous elements. But the difference between these two classes of elements is one of degree only, and they gradually merge into each other; moreover the electric relations of elements are not absolute, but vary according to the state of combination in which they exist, so that it is just as impossible to divide the elements into two classes according to this property as it is to separate them into two distinct classes of metals and non-metals. The following, however, are negative towards the remaining elements which are more or less positive :— [table]
Elements which readily enter into reaction with each other, and which develop a large amount of heat on combination, are said to have a powerful affinity for each other. The tendency of positive elements to unite with positive elements, or of negative elements to unite with negative elements, is much less than that of positive elements tα∙ unite with negative elements, and the greater the difference in properties between two elements the more powerful is their affinity for each other. Thus, the affinity of hydrogen and oxygen for each other is extremely powerful, much heat being developed by the combination of these two elements; when binary compounds of oxygen are decomposed by the electric current, the oxygen invariably appears at the positive pole, being negative to all other elements, but the hydrogen of hydrogen compounds is. always disengaged at the negative pole. Hydrogen and oxygen are, therefore, of very opposite natures, and this is well illustrated by the circumstance that oxygen combines,, with very few exceptions, with all the remaining elements, whilst compounds of only a very limited number with hydrogen have been obtained.
Hydrides. — The only elements which are known to form, binary compounds with hydrogen are fluorine, chlorine^ bromine, iodine, oxygen, sulphur, selenium, tellurium, carbon, silicon, nitrogen, phosphorus, arsenic, antimony, and copper. Palladium has the property of absorbing a large volume of hydrogen gas, and it is supposed forms a definite compound with it; sodium and potassium, when heated to 350° to 400° C., also absorb hydrogen and appear to form definite compounds. With the exception of carbon, which furnishes an enormous number of compounds with hydrogen, the above-named elements unite with hydrogen only in one proportion, or at most in two or three different proportions. Excepting oxygen, fluorine, chlorine, and bromine, they do not readily enter into reaction with hydrogen, so that in most cases, in order to obtain their compounds with hydrogen, it is necessary to resort to indirect methods of preparation. The compounds of the elements with hydrogen may all be included under the general title of hydrides, but usually they possess specific names; thus, the hydride of oxygen is known as water, and the hydride of nitrogen as ammonia. Some of the hydrides of carbon are gases, but the majority are either liquids or solids; the two known hydrides of oxygen, water and the so-called peroxide of hydrogen, are liquids; a gaseous and a liquid hydride of sulphur are known; three hydrides of phosphorus are said to exist, one of which is liquid, one solid, and one gaseous; hydride of copper is a solid; and the hydrides of the remaining elements are gases. Excepting the hydride of copper, all are colourless substances. The following is a list of the known hydrides, including, however, only the first or simplest hydride of carbon :— [table]
Oxides. — The element oxygen is known to form compounds with all the elements excepting fluorine, but the affinity of the various elements for oxygen, and consequently the stability of the compounds resulting from their union with oxygen, varies within very wide limits; the oxides of many elements which have but a weak affinity for oxygen can only be obtained by indirect means.
Although many of the elements unite with oxygen in several proportions, none are known to furnish more than •five distinct oxides, and these belong to a certain limited number of forms, since they contain one, two, three, four, or five atoms of oxygen in the molecule. But as the same number of atoms of oxygen may unite with a given element in more than one proportion, each of these five classes includes sub-classes. The composition of the oxides may be best illustrated and their behaviour explained by regarding them as formed on the type of the two oxides of hydrogen, hydrogen monoxide or water, OH 2 , and hydrogen dioxide, O 2 H 2 .
The monoxides are formed on the type of a single molecule of water, that is to say, they may be regarded as derived from a single molecule of water by the displacement of the two atoms of hydrogen by the equivalent amount of other elements. If two atoms of a monad element, or single atom of a dyad element, or of a polyad element which functions as a dyad, displace the two atoms of hydrogen, normal monoxides result, examples of which are potassium monoxide, K' 2 O, and copper oxide, Cu"O. But monoxides are known in which two atoms of a dyad or it may be polyad element displace the two atoms of hydrogen in water; it is supposed that in this case the two atoms of polyad metal are united together and therefore function as a dyad compound radicle. Monoxides of this kind are usually termed suboxides; the suboxide of copper, (Cu 2 )"O, and the suboxide of palladium, (Pd 2 )"O, may be cited as examples. The supposed constitution of the three classes of monoxides is represented by the following graphic formulae :—
0
K—O—K Cu=O Cu—Cu.
Potassium oxide, copper oxide. Copper suboxide.
Two classes of dioxides may be distinguished :—the one formed from two molecules of water by the displacement of the four atoms of hydrogen by a single tetrad atom, such as carbon dioxide, C ,τ O 2 , and tin dioxide, Su lτ O 2 j the other formed from a single molecule of hydrogen dioxide by the displacement of the two atoms of hydrogen either by two monad atoms, or by a single dyad atom, such as potassium dioxide, K' 2 O 2 , barium dioxide, Ba"O 2 , and manganese dioxide, Mn"O 2 . These two kinds of dioxide differ enormously in chemical properties; their supposed constitution is represented by the following graphic formtdae :—
K—0 0
I Ba< I O=C=O K—0 0
Carbon dioxide. Potassium dioxide. Barium dioxide.
The dioxides derived from hydrogen dioxide are usually termed peroxides.
The trioxides are divisible into three classes, but all may be regarded as derived from three molecules of water. In the one class the six atoms of hydrogen are displaced by a single hexad atom, as in sulphur trioxide, S τ, O 3 , and tungsten trioxide, W τι O 3 ; in the second class the six atoms of hydrogen are displaced by two polyad (tetrad 1) elements, which together function as a hexad compound radicle, as in ferric oxide, (Fe 2 ) τ, O 3 , and chromic oxide, (Cr 2 )' ’ O 3 . A third class may be regarded as formed by the displacement of the six atoms of hydrogen by two triad atoms, and includes boron trioxide, B"' 2 O 3 , and bismuth trioxide, Bi"' 2 O 3 . There are considerable differences in properties between these three classes of trioxides; their graphic formulae are as follows :—
ø Fe=O Bi=O
II O< I θ<
O=S=O Fe=O Bi=O
Sulphur trioxide. Ferric oxide. Bismuth trioxide.
As examples of tetroxides, osmium tetroxide, OsO 4 , antimony tetroxide, Sb 2 O 4 , the so-called magnetic oxide of iron, Fe 3 O 4 , and lead tetroxide, Pb 3 O 4 , may be quoted. These oxides undoubtedly belong to different classes, but too little is known of them for their relations to each other to be defined.
Lastly, the few pentoxides which are known may, with scarcely an exception, be regarded as derived from five molecules of water, the ten atoms of hydrogen being displaced by two pentad atoms. This is illustrated by the following graphic formulae of phosphorus and arsenic pentoxides;—
0 0 0 0
II II II il
P-O—P As—0—As
II 11 II II
oo oo
Phosphorus pentoxide. Arsenic pentoxide.
Several of the oxides of chlorine and of nitrogen, and one of the oxides of sulphur, are gases; the oxides of hydrogen, and the oxides of chlorine and of nitrogen which are not gaseous, are liquid; and the remaining oxides are solid bodies.
Acids — Bases — Salts. — Many oxides have an extremely powerful affinity for water, and readily combine with it; but the compounds formed in this manner by different oxides differ remarkably in properties. For example, when sulphur trioxide, SO 3 , is added to water, a solution is obtained which has the property of changing the colour of blue vegetable colouring matters, such as litmus, to red, and possesses a sharp sour taste; but when the oxide of a highly positive metal, such as barium oxide, is added to water, a solution capable of restoring the blue colour to reddened litmus is obtained. In both cases the water combines with the oxide, the sulphur trioxide being converted into sulphuric acid: —
SO 3 + OH 2 -SO 4 H 2 , and the barium oxide into barium hydroxide :— BaO + OH 2 = BaO 2 H 2 .
Sulphuric acid and barium hydroxide are representatives of two most important classes of compounds, the acids and bases ; the oxides which furnish acids when combined with water are conveniently termed acid oxides, whilst those which furnish bases may be termed basic oxides.
The acid oxides, however, unite with the basic oxides to form a third class of compounds called salts, which are usually neutral bodies, that is to say, they have no action either on red or on blue litmus; thus, sulphur trioxide and barium oxide unite to form the salt barium sulphate:—
SO 3 + BaO = SO 4 Ba.
The term acid is applied by some chemists to what are here called acid oxides, and what we have called basic oxides are frequently spoken of as bases, the acids being regarded as salts of hydrogen; this view of the constitution of acids is adopted because they are formed by the union of oxides of negative elements with the oxide of the positive element hydrogen, just as what are ordinarily termed salts are formed by the union of oxides of negative with oxides of positive elements. In this article, however, the term acid is restricted to compounds containing hydrogen.
The acids formed from the oxides of the various elements in the manner above explained all enter into reaction with bases, and with especial readiness with those derived from the highly positive elements, in such a manner that salts are produced, the hydrogen of the acid being displaced by the metal of the base, and water formed; for example,
H 2 SO 4 + Ba"O 2 H 2 = BaSO 4 + 2OH 2 . Sulphuric acid. Barium hydroxide. Barium sulphate.
We may therefore define an acid to be a compound containing one or more atoms of hydrogen which may be displaced by a metal by the action of a base; but, as will be evident later on, although applicable to all acids, this definition is of rather too wide a character.
The hydrides of fluorine, chlorine, bromine, and iodine also readily enter into reaction with bases, exchanging their hydrogen for other elements; and as they also possess an acid taste and redden blue litmus, they are universally regarded as acids. The fact that these compounds are powerful acids is of considerable importance, as showing that it is the association of hydrogen with a negative radicle which renders it subject to displacement by positive elements—that is to say, a compound in which hydrogen is associated with a negative radicle will have the properties of an acid, and the more negative the radicle the more readily can the hydrogen be displaced. The composition of all known acids is entirely in accordance with this view; thus, in nitric acid a single atom of hydrogen is associated with the negative compound radicle NO 3 , in sulphuric acid two atoms of hydrogen are combined with the negative compound radicle SO 4 , and in phosphoric acid three atoms of hydrogen are associated with the negative compound radicle PO 4 .
The oxides of the most negative and most positive elements most readily combine with water to form acids or bases, whilst the oxides of the feebly negative or feebly positive elements either do not unite with water or form extremely unstable combinations. Similarly, the oxides of the most negative and most positive elements unite together to form stable salts, much heat being developed by their combination, whereas the oxides of feebly negative and feebly positive elements either do not unite or form very unstable salts. The several oxides of an element often differ remarkably in properties, however, one oxide exhibiting acid characters, whilst another has basic properties.
The oxides of most of the so-called metallic elements are acted upon by acids, such as sulphuric, nitric, and hydrochloric acids, a salt of the metal being formed in all cases in which corresponding salts exist, and the hydro gen of the acid being eliminated in combination with the oxygen of the oxide as water; for example :—
MnO + H 2 SO 4 = MnSO 4 + OH 2 .
Manganese oxide. Sulphuric acid. Manganese sulphate. Water.
If corresponding salts do not exist, however, and action take place, then other products are obtained. Thus, no stable sulphate corresponding with manganese dioxide exists, and when this body is decomposed with sulphuric acid oxygen is evolved :—
2MnO 2 + 2H 2 SO 4 = 2MnSO 4 + 2OH 2 + O 2 .
Manganese Sulphuric Manganese πτ . l ,. n dioxide. acid. sulphate. water ∙ Oxygen.
Actions of this kind invariably occur with the oxides which may be regarded as formed on the type of hydrogen dioxide.
Sulphides. — Sulphur enters into union with most of the elements forming compounds analogous to the oxides in composition and general properties; in fact, we may distinguish acid and basic sulphides, corresponding to the acid and basic oxides, and salts formed by the union of these two classes of sulphides. With very few exceptions the sulphides are solid bodies. Selenium and tellurium, which are closely allied to sulphur, also form compounds with many elements more or less closely resembling the sulphides, but they are of little importance.
Chlorides — Bromides — Iodides — Fluorides. — W ith very few exceptions compounds of all the elements with chlorine have been obtained, and, especially from a theoretical point of view, the chlorides are a class of bodies of the highest importance. The affinity of bromine and iodine for other elements being much lower than that of chlorine, and their compounds much less stable than the corresponding chlorides, comparatively few bromides and iodides are known. Fluorides of many of the elements have also been obtained.
Chlorine unites with many of the elements in two or more proportions, but at most six atoms of chlorine unite with a single atom of another element. The majority of the chlorides are either liquids or solid bodies which may be more or less readily volatilized.
Nitrides — Phosphides. — Nitrogen has but a slight affinity for other elements; its most important compounds are ammonia, NH 3 , cyanogen, C 2 N 2 , and the oxides of nitrogen. The only elements which combine with it readily are tantalum, titanium, tungsten, and vanadium, and most of its compounds can be prepared only by indirect methods.
Phosphorus readily combines with chlorine, bromine, iodine, oxygen, and sulphur, and with most of the metals; its compounds with the non-metallic elements are of considerable importance to the chemist; but its compounds with the metals have been little studied.
Scarcely any of the remaining elements form compounds of importance with elements other than those which have already been considered.
We now proceed to the description of the elements and their more important compounds, commencing with hydrogen, and then passing to the so-called non-metallic elements, which will be considered in the following order :—Hydrogen, oxygen, chlorine, bromine, iodine, fluorine, sulphur, selenium, tellurium, nitrogen, phosphorus, boron, carbon, and silicon; after which the remaining elements will be briefly described as much as possible in the order of their relationship to each other. Hydrogen is placed at the head of the list, because it is the unit or standard of comparison both for the atomic weights and valencies of the remaining elements, and it is now usual also to refer the densities in the state of gas of all compounds to hydrogen. On account of the number and variety of their compounds with other elements the non-metallic elements are by far the most important, and therefore are naturally considered before the metals.
Hydrogen.
Symbol, H; Atomic wt., 1; Molecular wt, 2; Valency, '.
This element was discovered by Cavendish in 1766, and was called by him inflammable air; the name of Hydrogen is derived from vδωp, water, and γewaω, to generate, on account of its forming water when burnt. It occurs in the free state in the gases of volcanoes, and by the aid of the spectroscope has been detected in the sun, stars, and nebulae; it chiefly exists in combination with oxygen as water, and is an important constituent of all vegetable and animal substances.
Hydrogen is obtained by the decomposition of water in various ways. Thus, when two platinum plates, connected with the poles of a voltaic battery, are plunged into water acidulated with a few drops of sulphuric acid, hydrogen is evolved from the plate in connection with the negative pole, oxygen being disengaged from the positive pole. Water is also resolved into its elements when its vapour is heated by passing through an intensely ignited platinum tube, or by discharging electric sparks in an atmosphere of steam; but under these circumstances, owing to the recombination of much of the hydrogen and oxygen, only a small quantity of the mixed gases is obtained.
Hydrogen may also be obtained from water by the action of the highly positive metals, caesium, rubidium, potassium, sodium, and lithium, at ordinary temperatures; in each case a solution of the metallic hydroxide is obtained, and hydrogen evolved, thus:—
2Na + 2OH 2 = H 2 + 2N τ aOH. Sodium Water. Hydrogen. Sodium hydroxide.
Barium, strontium, and calcium also decompose water in the cold, and hydrogen is evolved when magnesium is heated with water at temperatures below the boiling point of the latter. Many other metals, such as iron, manganese, zinc, cadmium, cobalt, nickel, tin, and antimony evolve hydrogen from water when its vapour is passed over the metal heated to redness, the oxide of the metal being formed, thus:—
3Fe + 4OH 2 = 4H 2 + Fe 3 O 4 .
Iron. Water. Hydrogen. Fe ^ e fcrric
Aluminium also decomposes water at a red heat, but owing to the formation of an impermeable coating of aluminium oxide on the surface of the metal the action soon stops. Metals like copper, mercury, silver, gold, and platinum are without action even at a bright red heat. Many metals, however, which do not decompose water unless heated with it, if placed in contact with a more negative element cause the evolution of hydrogen at ordinary atmospheric temperatures; for example, if a plate of zinc, coated with spongy copper by immersion in a solution of copper sulphate, be placed in water, hydrogen is gradually evolved, the reaction which occurs being as follows :—
Zn + 2OH 2 = H 2 + Zn(OH) 2 Zinc. Water. [Hydrogen. Zinc hydroxide.
Hydrogen is usually prepared by the action of zinc or iron on a solution of hydrochloric or sulphuric acid. The change which occurs is represented in the following equations :—
Zn l· 2HC1 = H 2 + ZnCl 2
Zinc. Hydrochloric acid. Hydrogen. Zinc chloride
Zn + H,>⅛O 4 = H 2 + ZnSO 4 . Zinc. Sulphuric acid. Hydrogen. Zinc sulphate.
All metals which readily decompose water when heated readily furnish hydrogen on treatment with hydrochloric and sulphuric acid, and many other metals enter more or less readily (although none so readily) into reaction with these acids; also many other acids may be used in place of hydrochloric or sulphuric acid, but none act more readily. In all cases the action consists in the displacement of the hydrogen of the acid by the metal employed, and if the acid is not one which can enter into reaction with the displaced hydrogen, the latter is evolved as gas. If pure hydrogen is required, it is necessary to employ pure zinc or iron; the impurities in the ordinary metal communicate an extremely disagreeable odour to the gas.
On the large scale nearly pure hydrogen may be prepared by passing steam over charcoal or coke heated to dull redness. If the temperature be kept sufficiently low, hydrogen and carbon dioxide are the sole products :—
C + 2H 2 O-=2H 2 + CO 2 and the latter may be removed by causing the gas tn traverse a vessel filled with slaked lime; but if the temperature be allowed to rise too high carbon monoxide is also produced, and cannot be removed from the mixture.
Pure hydrogen is a colourless, transparent, odourless and tasteless gas. It has never been liquefied, and is very slightly soluble in water, 100 volumes of water dissolving 1'93 volumes, of the gas at all temperatures between 0 β and 24° C. It is the lightest of all known bodies, its specific gravity being ’ 0693, that of air being unity; one litre of hydrogen at 0° C, and under the pressure of 7 60 mm. of mercury, weighs O896 gramme, and it is important to remember this number, since the weight of a litre of any other gas may be at once found by multiplying 0896 by half the molecular weight of the gas, the specific gravity of a gas referred to hydrogen being always half its molecular weight (p. 471).
Pure hydrogen is not poisonous, though it cannot support life; and if mixed with a certain proportion of oxygen, it may be breathed for some time without inconvenience. It is extremely inflammable, and burns in the air with a colourless non-luminous flame forming water; a burning taper is extinguished when plunged into hydrogen, and all bodies which burn in the air are incapable of burning in hydrogen.
Hydrogen does not spontaneously enter into reaction with any of the elements, although it has a powerful affinity for several of them. Thus, when hydrogen and oxygen are mixed nothing occurs, but if a burning taper or a heated wire is applied, a violent explosion ensues, water being produced. Similarly, chlorine and hydrogen are without action upon each other in the dark, but if the mixture is exposed to a bright light or is heated by the passage of an electric spark, the gases at once combine with explosive violence, forming hydrochloric acid. It has already been pointed out that we must suppose that the molecules which constitute free hydrogen, chlorine, and oxygen gases are diatomic, and that hydrogen and chlorine, for example, do not directly combine to form hydrochloric acid, but that an action occurs such as is represented by the equation—
H 2 + C1 2 = HC1 + HC1, and that, therefore, the combination of the atoms of chlorine and hydrogen is preceded by the separation of the twσ atoms of chlorine in the chlorine molecules, and of the two atoms of hydrogen in the hydrogen molecules, from each other, which necessarily involves an expenditure of energy. The application of heat, we may assume, serves to effect this decomposition of some of the molecules of hydrogen, oxygen, and chlorine into their constituent atoms, and since much heat is evolved by the subsequent combination of the atoms of hydrogen and chlorine, or of hydrogen and oxygen, it is only necessary to apply heat to start the reaction, because the remaining molecules are decomposed by the heat developed in the formation of the first portions of water or hydrochloric acid.
Compounds of hydrogen with bromine, iodine, sulphur, and nitrogen may also be obtained directly from these elements and hydrogen, although only with difficulty and in small quantity.
Oxygen.
Symbol, 0; Atomic wt., 15 - 96; Molecular wt., 3Γ92; Valency, u .
Oxygen was first isolated by Priestley in 1774; its name is derived from oξυς, sour, and γεννάω, to generate, in allusion to the circumstance that many of the bodies formed by combining it with other elements dissolve in water, producing sour or acid solutions. It was long believed, in fact, that oxygen was a constituent of all acids.
Oxygen is the most abundant and the most important of all the elements. About one-fifth of the atmosphere consists of free oxygen; it is the chief constituent by weight of water; it is an important constituent of all animal and vegetable substances, and is contained in greater or less proportion in most mineral substances.
Oxygen may be obtained from water in the manner already mentioned under hydrogen, by decomposing it by an electric current, and also by transmitting chlorine gas and steam through a porcelain tube heated to bright redness—
2Cl 2 + 2OH 2 = O 2 + 4HC1;
Chlorine. Water. Oxygen. Hydrochloric acid.
the oxygen is freed from hydrochloric acid and any excess of chlorine by passing it through a solution of sodium hydroxide.
The most interesting method of obtaining oxygen, although it is not a usual method of preparing it, is that by which it was first isolated by Priestley. When mercury is heated in contact with air to a temperature just below its boiling point, it gradually becomes covered with a red scale of mercuric oxide, HgO, and when this red scale is exposed to a considerably higher temperature it is broken up into oxygen and metallic mercury.
The oxides of gold, platinum, and other metals which have but a slight affinity for oxygen, are very readily decomposed when heated, oxygen being evolved and the metal remaining. Many other metallic oxides, and especially those which may be regarded as formed on the type of hydrogen dioxide, which readily breaks up into water and oxygen, when more or less strongly heated, are resolved into oxygen and a lower oxide :—lead dioxide, PbO 2 , barium dioxide, BaO 2 , and manganese dioxide, MnO 2 , for example. In the case of manganese dioxide, 3MnO 2 give Mn 3 O 4 , the decomposition taking place at a red heat. Barium dioxide when strongly heated gives up one-half of its oxygen to produce barium monoxide, 2BaO 2 = O 2 + 2BaO; and by passing moist air over less heated barium monoxide it may be reconverted into the dioxide, which may be decomposed by a stronger heat. By repetitions of these processes it is possible to procure large quantities of oxygen with the aid of a small quantity of barium dioxide, and it has been proposed to employ this method for the preparation of oxygen on the large scale, but in practice there have been found difficulties attending its use.
Ordinarily when pure oxygen is required it is prepared by heating potassium chlorate, which ultimately furnishes potassium chloride and oxygen: 2KC1O 3 = 3O 2 + 2KC1. This decomposition requires a high temperature, and can only be effected in vessels of hard glass, but when the chlorate is mixed with about one-eighth of its weight of a metallic oxide, such as copper oxide, ferric oxide, or manganese dioxide, the oxygen is given off at a considerably lower temperature, and with great rapidity; in this case, however, it is impure, being always contaminated with small quantities of chlorine.
When potassium chlorate is heated alone the salt fuses, but after a considerable evolution of oxygen has taken place the fused mass becomes pasty, and on examination it is found to consist of a mixture of potassium chloride and potassium perchlorate, KC1O 4 , so that the first stage in the decomposition may be represented by the equation—
2KC1O 3 = O 2 + KC1O 4 + KC1;
Potassium chlorate. Oxygen. Potassium perchlorate. Potassium chloride, the potassium perchlorate is resolved on further heating into oxygen and potassium chloride. When a metallic oxide is mixed with the chloride, however, and heat applied, the latter does not fuse, and the formation of potassium perchlorate cannot be detected at any stage of the decomposition. It is difficult to explain the manner in which the metallic oxide acts in promoting the decomposition of the chlorate, since it is found to be unchanged at the conclusion of the reaction; but it is a well-known fact that many bodies which, under ordinary circumstances, do not yield oxygen, readily part with this element when another substance having a tendency to combine or enter into reaction with oxygen is introduced into the sphere of action. Thus, when potassium dichromate is heated with concentrated sulphuric acid, oxygen is evolved, the yellow solution becoming green owing to the formation of chromium sulphate—
2K 2 Cr 2 O 7 + 10H 2 SO 4 =
Potassium dichromate. Sulphuric acid.
3O 2 4- 2Cr 2 (SO 4 ) 3 + 4KHSO 4 + 8OH 2 .
Oxygen. Chromium sulphate. H ’ r ¾ n h ^ a8eiunl Water.
But if the dichromate is dissolved in a considerable quantity of water no change of this kind takes place even after prolonged heating; when, however, a body like sulphurous acid, H 2 SO 3 , which has a tendency to unite with oxygen to form sulphuric acid, is added to the solution, it at once becomes green,—the dichromate being decomposed as represented in the above equation, and the oxygen being fixed by the sulphurous acid. Hence it is not improbable that the metallic oxide acts in virtue of a tendency to form a higher oxide, which higher oxide is no sooner produced, however, than it is broken up into oxygen and the lower oxide. The lower oxide may again undergo conversion into the higher oxide and the latter be broken up into oxygen and lower oxide, so that a relatively small quantity of the metallic oxide may suffice to promote the decomposition of a relatively large quantity of chlorate.
When a small quantity of cobalt sesquioxide, or a few drops of a solution of a cobalt salt, is added to a clear concentrated aqueous solution of bleaching powder, which is then gently heated, oxygen is evolved with great regularity, and the bleaching powder is completely resolved into oxygen and calcium chloride—
2CaCl 2 O = O 2 + 2CaCl 2 . Bleaching powder. Oxygen, calcium chloride.
The bleaching powder solution is not decomposed in this manner when heated alone, and there is no doubt that the cobalt oxide induces the decomposition by acting as a carrier of oxygen; that is to say, it first takes oxygen away and passes to a higher state of oxidation, and then gives it up again, the higher oxide being an extremely unstable body.
When silver oxide is added to a solution of hydrogen dioxide in water, metallic silver, water, and oxygen are produced—
Ag 2 O + H 2 O 2 = O 2 + 2Ag + OH 2 .
oxide H dioxîde n 0χ y≡ en ∙ suτβr ∙ water ∙
This reaction is especially interesting, as it affords experimental evidence in favour of the assumption that the molecule of oxygen consists of two atoms; moreover, a very considerable quantity of heat is developed in the reaction, and since there is every reason for believing that the separation of an atom of oxygen from silver oxide, and of an atom of oxygen from hydrogen dioxide, are both operations which involve an expenditure of energy, there can be little doubt that the heat developed is due to the combination of the atoms of oxygen to form molecules. In other words, this reaction affords evidence that oxygen atoms have a powerful affinity for each other.
The green leaves of plants with the aid of sunlight are enabled to decompose water and carbon dioxide, and evolve oxygen from them.
No method has yet been devised of directly separating pure oxygen frem air, but Graham has shown that it is possible to obtain an “air” containing about 41 - 6 per cent, by volume of oxygen instead of about 20 ’ 8 per cent., which is the amount present in ordinary air, by dialyzing air through india-rubber. For this purpose a bag composed of the thinnest india-rubber supported on cloth is connected with a Sprengel air-pump; if the pump is kept in action, after the air is exhausted from the bag, it is found that it continues slowly to deliver “air” which, however, is richer in oxygen than ordinary air in about the proportions above-mentioned. The gases do not pass through actual pores, but apparently they are dissolved by the rubber, which is thus wetted through by the liquefied gases, and evaporate into the vacuum on the other side of the membrane, the increased amount of oxygen being doubtless due to the greater solubility of oxygen in rubber.
Many other methods of obtaining oxygen are known, but the above given are sufficient to illustrate the nature of the changes by which it is produced.
Oxygen is a colourless, odourless, and tasteless gas, which has hitherto resisted all attempts to liquefy it; it is only slightly soluble in water, 160 volumes of which at ordinary atmospheric temperatures dissolve about three volumes of oxygen. All bodies which burn in the air burn with greatly increased brilliancy and rapidity in oxygen; thus, a glowing splinter of wood bursts into flame when plunged into oxygen, and burns with great brilliancy; and even iron, if heated to redness before it is introduced into the gas, readily burns in it. In all cases in which bodies burn in oxygen the substances burnt combine with the oxygen to form new substances, and the heat and light developed are a consequence of their union. The weight of the products corresponds exactly with the weight of the body burned plus the weight of the oxygen consumed. The combination of two or more bodies when thus accompanied by the development of heat and light is termed combustion, the body burnt being the combustible, and the body in which the burning takes place the supporter of combustion. Thus, hydrogen is a combustible since it burns in oxygen, but oxygen is a supporter of combustion. These terms are merely relative, however, since oxygen may just as readily be burnt in hydrogen as hydrogen in oxygen; and, similarly, air may be burnt in coal gas.
The combination of oxygen with other elements is termed oxidation. In all cases of oxidation heat is developed, but it depends entirely upon the rapidity with which the oxidation is effected whether light is also produced, that is to say, whether what is ordinarily termed combustion takes place. Thus, when iron is burnt in oxygen, the combination of the two elements is effected with great rapidity, and a large amount of heat is developed within a very short space of time, and hence the product of combustion is intensely heated and becomes incandescent; when, however, iron slowly oxidizes or rusts, no light is produced, although actually more heat is developed than when the same weight of iron is burnt in oxygen, the oxide Fe 3 O 4 being formed in the latter and the oxide Fe 2 O 3 in the former case.
The various elements enter into reaction with oxygen with very various degrees of readiness, but as in the case of hydrogen and oxygen it is mostly necessary at least to start the reaction by the application of heat; and if the combination of the two elements can give rise to the development of only a moderate amount of heat, it is usually necessary to continue the application of heat until the oxidation is complete. Phosphorus slowly absorbs oxygen, but the remaining non-metallic elements are not affected by it at ordinary temperatures. The highly positive metals readily absorb it; the majority of the metals, however, when in the massive state are unacted upon in dry oxygen, but undergo oxidation more or less readily in moist oxygen or air. The coating of oxide first formed frequently protects the metal from more than a superficial oxidation. Some of the metals when in a state of very fine division, for instance, lead as obtained by the ignition of its tartrate, undergo oxidation so readily, however, that spontaneous combustion results from their mere exposure to air or oxygen. The spontaneous combustion of substances such as woollen refuse, greasy tow, and hay is a precisely similar phenomenon,—the heat developed by their slow oxidation being to a great extent retained, as they are bad conductors of heat, until finally it becomes sufficient to inflame them. Most animal and vegetable substances when exposed to the air decay, and it is generally supposed that they simply undergo slow oxidation. Pasteur has shown, however, that the oxidizing power of atmospheric oxygen is much exaggerated, and that the decay of animal and vegetable substances exposed to air is not simply the result of the action of the atmospheric oxygen, but of the action of oxygen assisted by microscopic organisms, the decay taking place at an extremely slow and almost imperceptible rate when these organisms are entirely excluded.
When the colourless gas nitric oxide is mixed with oxygen, red fumes of higher oxides of nitrogen are formed, and by this reaction oxygen may with facility be detected and distinguished from other gases. A solution of potassium hydroxide and pyrogallol, or pyrogallic acid, or an ammoniacal solution of cuprous chloride, absorbs oxygen with avidity and may be employed to remove it from a mixture of gases Ozone,
Symbol, O 3 ; Molecular wt., 47'88.
When exposed to the action of electricity, and especially under the influence of what is termed the silent discharge, oxygen undergoes a contraction of volume and acquires remarkably different properties, its chemical activity being greatly enhanced. It has been shown that the change consists in the conversion of the oxygen into an allotropic modification which has received the name of ozone (ο£ω) in allusion to its peculiar odour; and from Sir Benjamin Brodie’s experiments (R. Soc. Trans., 1872, p. 435) there can be no doubt that ozone differs from ordinary oxygen merely in that its molecule contains three atoms; this difference is expressed in the following graphic formulae :—
O=O ∕ 0 ∖
0—0
Oxygen. Ozone.
Various forms of apparatus are employed in ozonizing oxygen, but all are constructed on the principle of the Siemens’s induction tube. This consists of two wide tubes of nearly equal diameter placed one within the other and coated on their exterior surfaces with tinfoil; the coatings of tinfoil are connected with the terminals of a powerful induction coil, and a current of oxygen is passed through the narrow space between the two tubes, and is thus submitted to the action of the electric discharge. In the apparatus employed by Brodie the coatings of tinfoil are dispensed with, but the inner tube is filled with water in which is placed one of the terminal wires of the induction coil, and the outer tube is immersed in a vessel of water in connection with the other terminal wire of the coil.
It is essential that the oxygen submitted to the electric action be pure and in a very dry condition; and it is especially desirable to prevent the elevation of temperature consequent on the electric action, which may be done by placing fragments of ice in the interior of the tube and also in the water contained in the external cylinder in which the induction tube is immersed.
But it is not possible in this manner to convert more than about 15 per cent, of the oxygen into ozone, and Brodie’s experiments prove that there is a fixed limit prescribed by the conditions of experiment, beyond which the formation of ozone cannot pass. The explanation of this appears to be that the formation of ozone belongs to the class of actions termed dissociation phenomena; that is to say, not only is oxygen converted into ozone by the electrical action, but ozone is also reconverted into oxygen, the amount of ozone actually obtained under given conditions of experiment being dependent upon the extent to which these two opposite kinds of change take place. We may therefore expect that the amount of oxygen finally obtained in the form of ozone will be greater the less the ozone is exposed to the electric action; and that this is actually the case is proved by the fact that no advantage is gained by submitting the oxygen more than once to the electric action or by passing it slowly through the induction tube, and also by the behaviour of carbon dioxide. When this gas is submitted to the action of the electric discharge, it is partially resolved into carbon monoxide and oxygen, and the latter is in part transformed into ozone. Brodie has shown that it is not difficult, by passing a rapid current of carbon dioxide through the induction tube, to convert 75 per cent, of the oxygen eliminated from the carbon dioxide into ozone, and has even succeeded in converting as much as 85 per cent.
Pure ozone has never yet been obtained, however. Ozone is also formed in small quantity when water is decomposed by the electric current, employing plates of platinum or gold as electrodes; the amount produced is greater the smaller the electrodes. The slow oxidation of phosphorus in moist air is said to be attended with the production of small quantities of ozone, and it is probable that ozone is formed in other slow oxidations; it appears that in all such cases the formation of ozone is accompanied by that of hydrogen dioxide, a fact which is also true of electrolytic ozone.
Traces of ozone are usually present in the atmosphere, especially in the open country; and it no doubt plays an important part in the removal of organic impurities from the atmosphere.
Ozone is a colourless gas of peculiar unpleasant odour; 100 volumes of water at 0° C. only absorb about ‘5 volume, so that it is considerably less soluble than oxygen. Air charged with ozone exerts an irritating action upon the respiratory organs. Ozone is chiefly remarkable, however, on account of its oxidizing power. Thus, dry mercury, which is not in the least affected by ordinary oxygen, unless heated to near its boiling point, is at once oxidized by ozone, and dry iodine absorbs ozone and is oxidized by it. It bleaches a solution of indigo, which is oxidized by it. It also rapidly corrodes organic substances, such as cork and caoutchouc, which therefore cannot be employed in experiments with ozone. Paraffin, however, is not attacked by it, and an excellent air-tight joint between two glass tubes may be made by means of it. A piece of glass tube, into which they exactly fit, is slipped over the two tubes, which are placed close together, and a fragment of pure paraffin is placed at the external junction of the tubes; on gently heating the paraffin it melts, and the liquid being extremely limpid runs into and fills up the narrow space between the tubes.
In the conversion of oxygen into ozone the volume contracts by one-third, three molecules of oxygen furnishing two molecules of ozone.
3O 2 = 2O 3 ; Oxygen. Ozone.
but when ozone is reconverted into oxygen the volume increases to the same extent that it diminishes when oxygen is converted into ozone, since
2O 3 = 3O 2 . Ozone. Oxygen.
The conversion of ozone into oxygen may be effected by heat. Thus, when ozonized oxygen is passed through a glass tube heated to 110° C. slight decomposition of the ozone takes place; at 200° C. the decomposition is very rapid, about 97 per cent, being converted into oxygen; and below 300° C. decomposition is complete.
Similarly, ozone is decomposed to an unlimited extent by contact with metallic silver, and by manganese dioxide, lead dioxide, and copper oxide, without the substances undergoing more than an excessively minute increase of weight in the reaction. These decompositions afford instances of recurrent action, the active substance being alternately oxidized and reduced by the ozone; in the case of silver, supposing silver monoxide is formed, although the oxide produced is probably a higher oxide, the reaction may be represented in the following manner :—
2Ag + O 3 = O 2 + Ag 2 O;
Silver. Ozone. Oxygen. Silver oxide.
Ag 2 O + O 3 = 2O 2 + 2Ag. Silver oxide. Ozone. Oxygen. Silver.
Many reactions are known which prove that ozone has this power of effecting deoxidation, although itself a most powerful oxidizing agent; thus, by its action on a solution of sodium dioxide, sodium monoxide and oxygen are obtained :—
Na 2 O 2 4- O 3 = 2O. 2 + Na. 2 O. Sodium dioxide. Ozone. Oxygen. Sodium oxide.
The decomposition of ozone in this manner by silver, <fcc., is accompanied by the development of a considerable amount of heat; thus, according to Berthelot, no less than 29,600 units of heat are evolved in the reaction
2O 3 = 3O 2 .
A corresponding amount of energy must therefore be expended in the formation of ozone from ordinary oxygen. But the conversion of ozone into ordinary oxygen is to be regarded as occurring in two stages,—in the first the molecule is resolved into O 2 + 0, and in the second two atoms of oxygen from two molecules of ozone unite forming a molecule of ordinary oxygen. The splitting up of the ozone molecule into O 2 + O probably requires an expenditure of energy, so that the energy developed in the reaction in the form of heat is due to the combination of the atoms of oxygen to form molecules, but is less than that actually developed by their combination by the amount expended in the decomposition of the ozone molecules. As already pointed out, the fact that so large an amount of heat is developed by the combination of oxygen with oxygen indicates that the affinity of the oxygen atoms for each other is very considerable. This being the case, we are enabled also to understand how it is that ozone has such superior power as an oxidizing agent as compared with ordinary oxygen, for it is evident that, in reactions into which oxygen enters, for every 32 grammes employed an amount of energy corresponding to at least 29,600 heat units must be expended in order to separate-the two atoms forming the oxygen molecules from each other; whereas, mostly, when ozone is employed, the amount of energy to be expended is only that required to effect the decomposition of the ozone molecules into O 2 + O, which is probably very small.
The amount of oxidation effected by the same amount of ozone varies, however, according to the nature of the body oxidized, and Brodie’s experiments show that four classes of oxidations may be effected by ozone :—
In the first class of reactions one-third of the molecule is active in effecting oxidation, the remaining two-thirds being obtained as oxygen, so that the volume of oxygen obtained is the same as the volume of ozone operated upon, If the number of atoms of oxygen active in effecting oxidation are placed within brackets, this class of oxidations is expressed by the equation—
O 3 = O 2 + (O) I·
The aption of ozone on a solution of sodium dioxide, quoted above, and its action on a solution of potassium iodide—
2KI + OH 2 + O 3 = O 2 + I 2 + 2KHO, p ≡r ’-«· »“««· °™»· “>««· h p χ≡, on ferrous chloride and sulphate, on an acid solution of potassium ferrocyanide, and on sodium arsenite, are special cases of this class of oxidations.
In the second class of reactions, four atoms of oxygen from two molecules of ozone are active, the remaining two atoms being obtained as oxygen; or
2O 3 = O 2 + 4(O) . . . II.
The oxidation of strongly alkaline solutions of sodium hyposulphite, the oxidation of hydriodic acid, and the oxidation of barium pentasulphide are included in this class.
In a third class of reactions, ten atoms of oxygen from four molecules of ozone are active in promoting oxidation, a single molecule of oxygen being obtained; or
4O 3 ≈O 2 +1O(O) .... III.
The oxidation of sodium sulphydrate and of hydriodic acid at 0° C. are to be referred to this class.
Lastly, in a fourth class of reactions, the entire molecule is active in effecting oxidation; or
O 3 = 3(O) ΓV.
Examples of this class are afforded by the oxidation of tin dichloride, oil of turpentine, and neutral and slightly alkaline solutions of sodium hyposulphite.
That ozone can thus enter into reaction in so many different ways affords the most conclusive evidence of its compound nature, aud by these reactions it may be distinguished from all known gases, and estimated.
The reaction with potassium iodide is usually employed for the detection of ozone, the liberation of iodine being rendered evident by the addition of starch paste, which with iodine furnishes a beautiful blue colour. The liberation of iodine from potassium iodide is not alone sufficient proof, however, of the presence of ozone. By comparing at different times the depth of colour produced on exposing slips of paper moistened with starch paste and potassium iodide solution to the air, for the same length of time, an estimate of the relative amounts of ozone present may be obtained.
Compounds of Hydrogen with, Oxygen. Water.
Symbol, OH 2 ; Molecular wt., 17'96.
Hydrogen and oxygen when mixed do not enter into reaction at the ordinary temperature; but the contact of flame or of any other red-hot body, the passage of the electric spark, and the contact of platinum cause reaction to take place with explosion. The flame and the electric spark act by their intense heat, and the nature of the influence which, perhaps, heat exercises in inducing the combination of hydrogen and oxygen has already been pointed out; but the action of platinum is more obscure. Spongy platinum and the fine powder of that metal called platinum black, although cold, cause the mixed gases to explode as readily as flame does. Even polished slips of platinum, if perfectly clean, will cause the reaction to take place, although more slowly. It appears that at first the absorption of the gases on the surface of the cold metal, and the consequent approximation of their particles, is sufficient to induce the combination of a part of the gases; this being accompanied by the development of heat, the metal becomes warmed and acts more vigorously, and the combination of further quantities of the gases takes place, by which more heat is developed, so that by degrees the metal becomes red-hot, and if any of the mixed gases be still uncombined it causes them to explode. In the case of the powder or the sponge, especially the former, all this takes place so rapidly from the enormous surface of the metal, that it becomes red-hot as soon as it is introduced, and fires the mixture as rapidly as a flame. Now, it is well known that metallic platinum, and especially spongy platinum, or platinum black, readily condenses gases upon its surface; but the condensation is attended by the development of heat, and it seems not improbable, therefore, that its action in causing the combination of hydrogen with oxygen is of the same character as the action of flame,—that the platinum, in fact, does not itself induce the combination, but that the heat resulting from the condensation of the gases by the platinum is the direct cause of the reaction. Graham’s researches have shown also that heated platinum absorbs hydrogen, and the absorption is doubtless accompanied by the development of heat; this is an additional reason for the action of platinum becoming more vigorous as it becomes warmer.
A series of experiments by Von Meyer (Journal für praktische Chemie, 1876, p. 121) may here be referred to, as illustrating in a remarkable manner the influence of finely-divided platinum in causing the combination of gases. It was first shown by Bunsen that when a mixturo of hydrogen and carbon monoxide is exploded with insufficient oxygen to convert these two gases respectively into water and carbon dioxide, whatever the proportions employed, the quantities oxidized are always in a simple molecular ratio. From Von Meyer’s experiments it appears that, when a similar mixture is placed in contact with finely-divided platinum, although the oxidation is effected very slowly, the same law obtains. Whereas, however, when the mixture of the three gases is exploded, always relatively more hydrogen than carbon monoxide is oxidized, the proportion in which the two gases are burnt being on the average about as 3 to 1, the reverse is the case when their oxidation is slowly effected with the aid of platinum; in the latter case the number of molecules of carbon monoxide oxidized to the dioxide is from 7 to 8 times as great as the number of hydrogen atoms oxidized to water.
When oxygen and hydrogen, from whatever cause, enter into reaction, it is always in the proportions to form water, —that is, invariably in the proportion of 2 volumes of hydrogen gas to 1 volume of oxygen gas. Any excess of either is left uncombined. If the water which is formed be measured as steam or gas, and its volume compared with the volume of the mixed gases at the same temperature before combination, it is found that condensation has taken place, 2 volumes of water gas being obtained from 2 volumes of hydrogen and 1 volume of oxygen gas hence—∙
2H 2 + O 2 =2H 2 O. Hydrogen. Oxygen. Water.
In the liquid state, however, the volume of the water formed is so small as compared with the volume occupied by the mixed gases that it may be neglected, so that if the contraction in volume which occurs on explosion is measured at the ordinary temperature, two-thirds of the contraction represents the hydrogen, and one-third the oxygen, which has disappeared. This enables us to use hydrogen to determine the amount of free oxygen in air, or in any gaseous mixture, one-third of the contraction which occurs when a measured quantity of the gas is exploded with a measured excess of hydrogen representing the amount of oxygen present.
The affinity of oxygen for hydrogen, as measured by the heat developed by their combination, is very great, 68,376 units of heat, according to Thomsen, being evolved in the combination of 16 grammes of oxygen with 2O05 grammes of hydrogen, the product being liquid water at 18’ C. By burning the two gases together from a jet the most intense artificial heat that is known, except that of the discharge of a powerful galvanic battery, is produced.
It is well here to call attention to the circumstance that the equations ordinarily employed to express the formation of water from its elements, and vice versa the resolution of water into its elements, viz.—
2H 2 + O 2 =2H 2 O, and
2H 2 O = O 2 + 2H 2 , do not take into account that in the one case a large amount of heat is developed, and that in the other case a corresponding amount of energy is expended. These equations merely represent, in fact, the distribution of weight in the changes, and therefore are but imperfect expressions of what really occurs, since the development of heat in the formation of water, and the absorption of heat, or the expenditure of a corresponding amount of some other form of energy in its decomposition, are integral parts of the changes, the amount of heat developed or absorbed, under given conditions, being as definite and constant as the weights of the substances which enter into reaction and are produced. The same may be said of all equations employed to represent chemical change.
The composition of water may be determined by burning a known weight of hydrogen in an excess of oxygen, and weighing the water produced; then the difference between the weight of the hydrogen burnt and of the water produced is the amount of oxygen combined with the hydrogen. Or hydrogen is passed over a weighed quantity of copper oxide in a tube heated to redness; the hydrogen thou reduces the oxide or removes the oxygen from it, forming water, which is carefully collected and weighed, and the loss of weight which the tube of copper oxide suffers is carefully determined. The loss of weight of the copper oxide gives the amount of oxygen, and the difference between this and the amount of water produced is the amount of hydrogen combined with this amount of oxygen.
The amount of oxygen combined with 2 parts by weight, or 2 atoms, of hydrogen is usually stated to be 16 parts, and 16 is generally regarded as the atomic weight of oxygen. But from the examination of the determinations which have been made by various elements of the composition of water, Staas, to whom we are indebted for the most exact determinations of atomic weights yet made, arrives at the conclusion that this number is too high, and that the atomic weight of oxygen is certainly not higher than 15 ’ 96. Thomsen has recently determined the amount of water produced by burning 2 litres of hydrogen in an excess of oxygen, and taking Regnault’s numbers for the specific gravities of oxygen and hydrogen, he obtains a number for the atomic weight of oxygen which is in complete accordance with that given by Staas.
At the ordinary temperature of the air water is a clear, transparent, tasteless, and odourless liquid; it appears colourless when seen in small quantity, but that it really has a pale blue colour is apparent when a shining white object is viewed through a column several feet in thickness.
Water is solid at temperatures below 0° C., 0° C. being the temperature at which frozen water or ice melts; the melting point is diminished by increase of pressure to the slight extent of 0'00757° C. for each additional atmosphere. Water expands in freezing, its density compared with water at 0° C. being , 92. The conversion of liquid water at 0° C. into solid water or ice is accompanied by the liberation of heat, and heat is rendered latent or absorbed to the same extent in the melting of ice,—the quantity of heat absorbed or liberated in the melting of ice or freezing of water being sufficient to raise the temperature of 7 9 times its weight of water from 0° to l o C.
Water evaporates at all temperatures when in contact with atmospheric air or other gases, and the vapour given off has a density and tension determined by the temperature; the tension of the vapour rapidly rises with the temperature, until at 100° C. it is equal to the ordinary atmospheric pressure (760 mm.), and the water boils. The boiling point, however, rapidly rises with increase of pressure, and sinks when the pressure is diminished; thus under the pressure of two atmospheres water boils at 121° C., and under the pressure of twelve atmospheres at 190° C. When water boils under the ordinary atmospheric temperature it is converted into 1600 times its volume of vapour. The conversion into vapour is attended with the absorption of a large amount of heat,—the quantity of heat absorbed or rendered latent in the conversion of water at 100° C. into steam of the same temperature being sufficient to raise the temperature of 536 times the weight of water converted into steam from 0° to l 0 C.
Water, chemically speaking, is a remarkably neutral substance, and hence its great value to the chemist as a solvent. There are very few substances which are not to some extent dissolved by it, but the solubility of different substances is very unequal. Heat generally increases its solvent power, whilst cold diminishes it; there are many exceptions to this rule, however. The dissolution of substances which may again be separated from the solution undecomposed is accompanied, in the majority of cases, with an absorption of heat, as will be evident on inspection of the table on p. 485. In the first column the name of the substance is given, and in the second its formula; the third exhibits the number of molecules of water (in grammes) in which one molecule (in grammes) of the substance is dissolved at about 18° C.; the fourth column shows the number of units of heat developed or absorbed,the — sign indicating that heat is absorbed, and the + sign that it is developed.
It is extremely difficult to interpret the meaning of numbers such as are contained in the table, especially as we are almost entirely ignorant of the condition of substances in solution in water. But there is no doubt that the heat developed or absorbed on dissolving a solid substance is the mean result of several distinct operations, which partly, perhaps, involve an absorption, and partly a development of heat. Thus, in the first place, there is a change of state from the solid to the liquid, which in most if not all cases involves an expenditure of energy; then, many substances on dissolving in water combine with it, the combination probably being always attended with development of heat. Contraction also generally takes place in the dissolution of salts in water, and is accompanied by a considerable development of heat, arising from the great resistance which water offers to compression. Chemists, moreover, are inclined to the belief that very many if not all substances, even those which are ordinarily regarded as stable in presence of water, enter to a greater or less extent into reaction with water when dissolved in it; a solu tion of sodium chloride, for example, may be regarded as containing a certain amount of sodium hydroxide and hydrochloric acid formed by the reaction—
NaC1 + OH 2 = HC1 + NaOH .
Sodium chloride. Water. Hydrochloric acid. Sodium hydroxide.
Table of Phenomena attending Solution of Salts in Water. [table]
The occurrence of reactions of this kind would in many cases involve an absorption, but in others a development of heat. The only two substances mentioned in the above table which develop heat to any extent when dissolved in water, it will be seen, are magnesium and aluminium chlorides. Both of these, however, are known to form compounds with water, and both probably enter to a very considerable extent into reaction with it in the manner above pointed out. It is even probable that the latter cannot exist as such when dissolved in water; the development of so large an amount of heat is therefore readily understood. But at present we are unable satisfactorily to account for the difference observed between salts such as potassium and sodium iodides, which so closely resemble each other in most respects, one of which, it will be noticed, absorbs 5110 units, whilst the other develops 1220 units of heat when dissolved in water.
From this it will be obvious that the study of the condition of salts in solution is beset with difficulties; the thermochemical method of investigation appears in most cases to be the only one which is applicable, since the introduction of new substances at once introduces a new set of conditions, but on account of the compIexity of the phenomena attending dissolution, even the results obtained by this method possess only a limited value, and at present only general conclusions can be drawn from them.
Water, as we have already stated, enters into combination with oxides of many of the elements, forming two classes of compounds,—the acids, and the metallic hydroxides or hydrates. The general properties and relations of these two classes of compounds may with advantage now be discussed.
The monoxides of the highly positive monad elements caesium, rubidium, potassium, sodium, and lithium form with water easily soluble hydroxides, which cannot be decomposed by heat; their solutions are soapy to the touch, and restore the blue colour to vegetable infusions which have been reddened by an acid. These hydroxides are usually termed alkalies,—a solution which has the power of restoring the blue colour to reddened litmus being said to exhibit an alkaline reaction. The term alkali is of Arabic origin, and was at first given to the crude sodium carbonate obtained from the ashes of sea-weed, a solution of which is soapy to the touch and restores the blue colour to reddened litmus, and like the above-mentioned hydroxides also has powerful cleansing properties.
The hydroxides derived from the monoxides of barium, strontium, and calcium, which are less positive elements, also exhibit an alkaline reaction, but they are not nearly so soluble in water; and, with the exception of barium hydroxide, they are decomposed on ignition, yielding the oxide and water. None of the remaining positive elements, except thallium, furnish soluble hydroxides which exhibit an alkaline reaction.
Similar differences may be observed between the oxides of negative elements, which furnish acids when combined with water. Thus, the monoxide of the highly negative element chlorine readily dissolves in water, but the acid produced is exceedingly unstable; similarly, the acids derived from the oxides of nitrogen are soluble, but of low stability. The oxides of the less negative elements, sulphur and phosphorus, however, furnish very soluble acids, which exhibit considerable stability, being with difficulty resolved into the oxide and water.
The lower oxides of the most positive elements enter into combination with water in such proportions as to produce compounds containing an equal number of atoms of hydrogen and oxygen; for example, sodium monoxide and water furnish sodium hydroxide—
Na 2 O + OH 2 = 2NaOH;
Sodium monoxide. Water. Sodium hydroxide.
and barium oxide and water yield barium hydroxide—
BaO 4- OH 2 = BaO 2 H 2 or Ba(OH) 2 .
Barium oxide. Water. Barium hydroxide.
Hence these hydroxides may be looked upon as combinations of the respective elements with the monad compound radicle (OH)' or hydroxyl; and the various compounds obtained from the oxides of the remaining elements by the action of water may either be regarded as similarly constituted, or may be viewed as combinations of one or more OH groups with compound radicles formed by the union of the elements with oxygen. Thus, we may regard nitric acid, HNO 3 , as a compound of the two monad radicles, (NO 2 )'and (OH)'; sulphuric acid, H 2 SO 4 , as a combination of the dyad radicle (SO 2 ) with two OH groups; and phosphoric acid, H 3 PO 4 , as a combination of the triad radicle (PO) with three monad hydroxyl groups.
The number of OH groups which may be associated with a single atom of a given simple radicle or element, or with a compound radicle, entirely depends upon the nature of the radicle, but does not appear ever to exceed four; the stability of the compound also varies with the radicle, the tendency to form stable compounds with OH being the greater the more positive the radicle. In the case of compounds of feebly positive radicles with several OH groups, there is always a tendency for the elements of one or more molecules of water to separate from the compound, thus producing a body which is to be regarded as a combination of one or more OH groups with an oxygenated radicle. For instance, ferric hydroxide, Fe 2 (OH) 6 , very readily loses water, forming compounds such as Fe 2 O(OH) 4 and Fe 2 O 2 (OH) 2 .
Nearly all the bodies produced by the action of water on the oxides of the non-metals are to be regarded as compounds of oxygenated radicles with hydroxyl, and are more or less acid in character. As a rule, the more negative the radicle with which OH is associated the more powerful will be the acid, and since the addition of oxygen to the radicle renders it more negative, the acid furnished by the higher oxide of an element is usually more stable, and also a more powerful acid, than that furnished by the lower oxide; thus, sulphurous acid, H 2 SO 3 , is extremely unstable as compared with sulphuric acid, H 2 SO 4 , and a much less powerful acid, but the former may be regarded as a combination of the dyad radicle (SO), and the latter as a combination of the dyad radicle (SO 2 ), with twice the monad radicle (OH). The behaviour of the positive elements is exactly complementary of this, since the greater the amount of the negative element oxygen associated with them the less is the tendency to furnish corresponding hydroxides when combined with water, and the less the stability of the resulting hydroxides; and, as a rule, also the basic properties both of the oxides and of the hydroxides corresponding to them become much less pronounced as the relative proportion of oxygen increases.
In all cases in which oxides unite with water to produce stable combinations, much heat is developed; the stability of the compounds formed by the combination of water with oxides, in fact, appears to be directly in proportion to the amount of heat developed in their formation, and it has been shown that the formation of many of the most unstable hydroxides from their elements would be accompanied by an absorption of heat; hence their instability is readily understood.
When the hydroxides are added in sufficient quantity to solutions of the acids, mostly neutral solutions are produced,—that is to say, solutions which do not affect either blue or red litmus. The acid is then said to be neutralized by the hydroxide, or vice versa. The basic oxides which furnish hydroxides, and indeed the basic oxides generally, behave in a similar manner with acids. In these cases a salt is produced; thus, by the addition of sodium hydroxide to sulphuric acid the salt sodium sulphate is obtained—∙
2NaOH + H 2 SO 4 = Na 2 SO 4 + 2H 2 O. Sodium hydroxide. Sulphuric acid. Sodium sulphate. Water.
Obviously, we may regard the salts produced in this manner as formed by the displacement of the hydrogen of the acid by more or less positive elements or simple radicles, and they may actually in many cases be prepared by the action of the metals on the acids; for example—
Zn + H 2 SO 4 = H 2 + ZnSO 4 Zinc. Sulphuric acid. Hydrogen. Zinc sulphate.
They are also obtained, as we have seen, by the union of a more or less basic with a more or less acid oxide. But a large number of salts are known derived from the acids by the displacement of hydrogen by compound radicles; thus, many of the salts of the element vanadium are formed by the introduction of the group (V 2 O 2 ) lv , which functions as a tetrad element, in place of the hydrogen of acids, vanadyl sulphate being (V 2 O 2 ) ιv (SO 4 ) 2 . The ammonium salts are formed in a similar manner by the displacement of hydrogen in acids by the monad compound radicle (N τ H 4 )'. But the hydrogen of many hydroxides may also be displaced by positive elements or radicles; for example, zinc hydroxide, although insoluble in water, dissolves in solutions of the alkalies, and on heating metallic zinc with a concentrated solution of potassium hydroxide, it dissolves with evolution of hydrogen—
Zn + 2HOK = Zn''(OK) 2 + H 2 .
Zinc hydroxide and similar compounds, therefore, display both basic and acid functions; and it is difficult to deny the compounds formed from them by introducing positive elements in place of hydrogen the title of salts.
Another class of oxy-salts are formed by the union of two different oxides of the same element; the compounds of fluorine, chlorine, bromine, and iodine with positive elements, and of acid with basic sulphides, are also termed salts, the compounds of sulphur being distinguished as sulpho-salts or thio-salts, whilst those of the four other elements are distinguished as haloid salts.
From this it will be obvious that the term salt is of very wide application, and that it is almost impossible to define it in a few words. It is also extremely diffi cult strictly to define an acid, for, although the very greatest difference is observable between the compounds of the oxides of highly negative and of highly positive elements with water, the differences become less and less marked as we pass from one end of the series to the other. The only definition which really separates bodies which are usually regarded as true acids from hydroxides possessing acid properties is afforded by the fact that, whilst the hydrogen in all hydroxides which exhibit basic properties may be displaced by negative and also in some instances by positive radicles, the hydrogen in acids can only be displaced by positive radicles. For example, the hydrogen in zinc hydroxide may be displaced by the positive radicle potassium and by the negative radicle NO 2 , but it is only possible to displace the hydrogen in sulphuric acid by positive radicles. This definition is not applicable, however, to organic acids.
The extent to which hydrogen may be displaced in an acid by positive radicles is termed its basicity, an acid which contains a single atom of displaceable hydrogen being termed monobasic, and acids containing two or three atoms of displaceable hydrogen, dibasic or tribasic. On the other hand, the hydroxides and basic oxides are frequently spoken of as monacid, diacid, or triacid, <fcc., according as they are capable of neutralizing a single molecule, two, or three molecules of a monobasic acid; thus, sodium hydroxide is monacid since a single molecule neutralizes a single molecule of monobasic nitric acid, forming the salt sodium nitrate—
NaOH + HNO 3 ≈ NaNO 3 + OH 2 , Sodium hydroxide. Nitric acid. Sodium nitrate. Water.
and barium hydroxide is a diacid base since it neutralizes two molecules of nitric acid—
BaO 2 H 2 + 2HNO 3 = Ba(NO 3 ) 2 + 2OH 2 . Barium hydroxide. Nitric acid. Barium nitrate. Water.
But we may also regard the salts formed by the action of acids on hydroxides as derived from the latter by the displacement of hydrogen by negative compound radicles,—sodium nitrate, for example, as sodium hydroxide in which the atom of hydrogen has been displaced by the monad compound radicle NO 2 ; hence we may define the acidity of hydroxides to be the extent to which hydrogen may be displaced in them by negative radicles.
In polybasic acids, that is to say, in acids containing two or more atoms of displaceable hydrogen, it is possible to displace the hydrogen step by step; thus, by the addition of a single molecule of sodium hydroxide to a single molecule of sulphuric acid, only one-half the hydrogen in the acid is displaced—
NaOH + H 2 SO 4 = NaHSO 4 + OH 2 Sodium hydroxide. Sulphuric acid. Sodium hydrogen sulphate. Water, but by the addition of a second molecule of the hydroxide the second atom of hydrogen is also displaced—
NaOH + NaHSO 4 = Na 2 SO 4 + OH 2 .
Sodium Sodium hydrogen Sodium water hydroxide. sulphate. sulphate.
Salts derived from acids by the partial displacement of their hydrogen are termed acid salts, those in which the hydrogen displaceable by metals is entirely displaced being termed normal salts. Similarly, when the hydrogen in polyacid hydroxides is entirely displaced by negative radicles, normal salts are obtained, but when it is only partially displaced the so-called basic salts are produced, which bear precisely the same relation to the hydroxides that the acid salts bear to the acids. Thus, from the triacid hydroxide, bismuth hydroxide, Bi"'(OH) 3 , and the monobasic acid, nitric acid, we obtain normal bismuth nitrate and two basic bismuth nitrates :—
í NO 3 ( NO 3 ( NO 3
Bi { NO 3 Bi √ NO 3 Bi 7 OH .
( NO 3 (OH (OH
Normal salt. Basic salts.
These sales are strictly comparable with the normal and two acid salts produced from the tribasic acid, phosphoric acid, and the monacid hydroxide, sodium hydroxide :—
(ONa (ONa (ONa
PO √ ONa PO J ONa PO 7 OH .
( ONa (OH (OH
Normal salts. Acid salts.
Two explanations may be given of the constitution of the salts formed by displacing the hydrogen in polybasic acids by polyad elements; the formula of barium sulphate, for example, may be represented by either of the following formulae :—
SO 2 {°>Ba, SO s {g according as it is supposed that two atoms of hydrogen in a single molecule of the acid are displaced, or that two molecules of the acid are concerned in the formation of the salt, one atom of hydrogen in each being displaced by each of the barium atoms. There is not at present sufficient evidence to enable us to decide definitively in favour of one or the other of these views, but the great tendency which the polybasic acids exhibit to form double salts appears to render the latter the more probable.
The formation of salts by the action of acids on metallic oxides and hydroxides is in all cases attended with a development of heat. This subject has been very carefully studied by several chemists, especially by Thomsen, whose results for a number of acids are given in the following tables. The amount of heat developed in the reaction— __ zatt . λ .
αNaOHAq, Q Aq,—
that is to say, on adding a solution of α molecules (in grammes) of sodium hydroxide to a solution of one molecule (iu grammes) of the acid, represented by Q,—is given for a number of acids in the first of the following tables. The second table exhibits the amount of heat developed in the reaction—
NaOHAq , αQAq; t.e., on adding a solution of one molecule of sodium hydroxide to a solution of α molecules of the acid Q. Each molecule (in grammes) of sodium hydroxide was dissolved in 200 molecules (in grammes) of water, and the solutions of the acids were of equivalent strength; that is to say, the number of grammes of acid required to form a normal salt with 23 + 16 + 1 or 40 grammes of sodium hydroxide were dissolved in 18 × 200 or 3600 grammes of water. The temperature at which the two solutions were mixed in all the experiments was about 18° C. The results are expressed in hundreds of heat-units, and according to Thomsen they may all be relied on within 1 per cent.
(αNaOHAq, QAq). [table]
(NaOHAq,αQAq). [table]
To facilitate comparison between the various acids they are arranged in the following table almost in accordance with the amounts of heat developed on the addition of sodium hydroxide to equivalent quantities. [table]
From these tables it will be evident—(1), that when a molecule of sodium hydroxide in aqueous solution enters into reaction with an acid, the heat developed is very nearly proportional to the quantity of acid present until this amounts to 1, ½, ½, or ∣ molecule, according as the acid is mono-, di-, tri-, or tetrabasic; but that when the amount of acid added exceeds that requisite to form the normal salt, the different acids behave differently, heat being in some cases developed, and in others absorbed, according to the constitution of the acid; and (2), that mostly when a molecule of an acid in aqueous solution enters into reaction with sodium hydroxide, the amount of heat developed increases almost in proportion to the amount of the latter, and until 1, 2, 3, or 4 molecules are added, according as the acid is mono-, di-, tri-, or tetrabasic; the further addition of sodium hydroxide is not then attended with any considerable development of heat.
Very different amounts of heat, it will be observed, are developed on neutralizing the different acids, but there is mostly a remarkable similarity in the results obtained in the case of acids which from chemical evidence are known to be closely allied. Thus, hydrochloric, hydrobromic, and hydriodic acids exhibit the same heat of neutralization; and the numbers for sulphuric and selenic acids, and for phosphoric and arsenic acids, are very similar. Hydrofluoric acid, it will be seen, differs considerably from the allied monobasic acids since the addition of the first half molecule of sodium hydroxide produces less heat than the second, which is not improbably owing to the formation of the acid salt HF 2 Na; it differs also by its high heat of neutralization.
The tables show also that the dibasic acids may be divided into several groups, according to the amounts of heat developed on the addition of the first and second molecules of sodium hydroxide. The first group includes hydrofluosilicic and hydrochloroplatinic acids, the amount of heat developed by the second molecule of hydroxide being equal to that developed on the addition of the first molecule. The second group includes sulphuric, selenic, oxalic, and tartaric acids; with these acids less heat is developed by the first than by the second molecule; thus :— [table]
In the third group, which includes sulphurous, selenious, carbonic, and boric acids, and probably also chromic, phosphorous, and succinic acids, the contrary is the case [table]
The tribasic acids exhibit similar differences; thus :— [table]
These differences which acids of the same basicity exhibit when submitted to thermochemical investigation correspond to differences in their chemical behaviour. For example, when a solution of citric acid is neutralized with sodium carbonate and evaporated to crystallization, the trisodium salt Na 3 C 6 H 5 O 7 is readily obtained, but when a solution of phosphoric or arsenic acid is similarly treated, the disodium salt Na 2 HPO 4 or Na 2 HAsO 4 is formed; the trisodium salts of these acids can only be procured by adding sodium hydroxide. Apparently the trisodium salts of phosphoric and arsenic acids are partially decomposed by water, as their solutions are strongly alkaline; hence the third molecule of hydroxide does not effect the complete conversion of the di-into the tri-sodium salt. The behaviour of sulphuric acid will be discussed later on.
All soluble hydroxides when in solution appear to have nearly the same heat of neutralization, as will be seen from the following table, which exhibits the number of units of heat developed on neutralizing solutions of equivalent quantities of various hydroxides with a solution of one molecule (in grammes) of sulphuric acid, or of the equivalent quantities (2 molecules) of hydrochloric or nitric acid at 18° C. :— [table]
But very different amounts of heat are developed on dissolving the hydroxides which are insoluble in water in acids. The following numbers represent the heat of neutralization of various hydroxides insoluble in water by a molecule of sulphuric acid ∙— [table]
Thomsen, however, considers that the differences between the amounts of heat developed when these hydroxides are dissolved in acids, and when soluble hydroxides are neutralized, are due to the heat absorbed in rendering them soluble; or in other words, that solutions of these hydroxides would have the same heat of neutralization as solutions of soluble hydroxides.
One of the most interesting results of the thermochemical investigation of the behaviour of acids with metallic hydroxides is the proof that when sodium hydroxide, for example, is added to a mixture of two acids in insufficient quantity to neutralize both, the sodium salts of the two acids are seldom formed in the proportions in which the two acids are mixed. Thus, when a solution of 2 molecules of sodium hydroxide is mixed with a solution containing two molecules of monobasic nitric acid and one molecule of dibasic sulphuric acid, two-thirds of the sodium hydroxide enter into reaction with the nitric acid and one-third with the sulphuric acid. The avidity, as it is termed by Thomsen, of nitric acid to enter into reaction with sodium hydroxide is thus twice as great as that of sulphuric acid. The avidity of hydrochloric acid for sodium hydroxide is equal to that of nitric acid, but the avidities of hydrobromic, hydriodic, and hydrofluoric acids are less than that of hydrochloric acid in the proportion HC1: HBr :HI :HF= 100 :89 :79 :5.
In concluding this brief account of the thermochemical behaviour of some of the acids and hydroxides, we may observe that the interpretation of most of the remarkable results to which we have alluded has not been given. The study of this branch of chemistry is yet in its infancy, and is beset with difficulties, but its importance can scarcely be exaggerated. At present we scarcely know more than that in the formation of a given substance a certain amount of heat is finally developed or absorbed, but in most cases we are ignorant of the value of one or more of the several distinct operations which we believe are involved in the reaction by which it is produced, so that the results given by different substances cannot as a rule be satisfactorily compared.
Many substances belonging to all classes of chemical compounds unite with water, forming combinations from which it may be again expelled by the application of a greater or less degree of heat, and from the circumstance that many of these combinations are crystalline, the water is said to be present in the form of water of crystallization. Thus, potassium hydroxide crystallizes with 2 molecules of water, as KHO + 2H 2 O; copper sulphate with 5 molecules, as CuSO 4 + 5H 2 O; and sodium sulphate with 7 and with 10 molecules. The water may usually be expelled from these compounds by the application of a gentle heat, and many salts part with their water of crystallization on exposure to the air; but others are not deprived of it, or more than a part, unless very strongly heated. Copper sulphate, for example, readily loses 4 of the 5 molecules of water with which it combines, but the remaining molecule is only expelled by heating to about 200° C.; 5 of the 7 molecules of water in crystallized magnesium sulphate are readily removed, and the sixth is expelled at about 150° C., but the seventh is retained at 200° C.
The compounds containing water of crystallization are usually denied the title of atomic compounds ordinarily applied to combinations of two or more elements in which it is supposed the constituent atoms are all associated in a single molecule under the influence of the force which has received the name of chemical affinity, and in contradistinction are termed molecular compounds, being regarded as combinations of two or more separate molecules. This, although perhaps true of many of the compounds containing water of crystallization, is certainly not true of all, and notably of the sulphates which are only deprived of their water by heating to high temperatures.
Hydrogen Dioxide, H 2 O 2 = 33∙92.
This compound cannot be obtained directly from its elements, but there appears to be little doubt that it may be formed by the combination of oxygen with water, since it is produced in many cases of slow oxidation in presence of water, and particularly in processes where ozone is formed. It is produced in small quantity when water is decomposed by an electric current, especially when small electrodes are employed. Its formation in this manner may be regarded as an oxidation of water, but it has been suggested that it is the direct product of electrolysis, and that the oxygen evolved is a secondary product derived from the decomposition of the hydrogen dioxide; thus—
2H 2 O = H 2 + H 2 O 2 j 2H 2 O 2 = O 2 +2H 2 O .
It is usually prepared from barium dioxide by double decomposition with hydrochloric or carbonic acid :—
BaO 2 + 2HC1 = H 2 O 2 + BaCl 2 .
Barium dioxide. Hydrogen chloride. Hydrogen dioxide. Barium chloride.
Hydrogen dioxide or peroxide is an exceedingly unstable substance, and readily decomposes even in aqueous solution into water and oxygen, especially on heating. The solution is more stable if slightly acid. A dilute solution may be concentrated by evaporation in vacuo over sulphuric acid, and hydrogen dioxide was obtained in this manner by its discoverer Thénard as a colourless transparent liquid, of specific gravity 1 ’ 452, which did not freeze at - 30° C.; it is doubtful, however, whether it has ever been prepared quite free from water. It did not redden litmus, but had a harsh bitter taste; when placed upon the hand it instantly turned the cuticle white.
Hydrogen dioxide exhibits the closest resemblance to ozone. Thus, it is decomposed by mere contact with finely-divided metals, such as silver, gold, and platinum, with evolution of oxygen. Like ozone it is a powerful reducing agent, entering into reaction with silver oxide, for example, to form water, oxygen, and metallic silver :—
H 2 O 2 + Ag 2 O = H 2 O + 2Ag + O 2 .
According to Fairley, about 37,000 units of heat are developed in the production of 32 grammes of oxygen by this reaction, which thus affords further proof of the strength of the affinity of oxygen for oxygen, as probably the heat developed is chiefly, if not entirely, due to the combination of the oxygen atoms. Many other oxides are reduced by it, either to the metallic state or to lower oxides, oxygen being evolved and water produced.
But hydrogen dioxide is also a powerful oxidizing agent. Thus, it decolorizes a solution of indigo; it converts sulphurous acid, H 2 SO 3 , into sulphuric acid, H 2 SO 4 ; lead sulphide, PbS, into lead sulphate, PbSO 4 ; and many oxides into higher oxides. It decomposes a solution of hydriodic acid with separation of iodine :—
H 2 O 2 + 2HI = 2H 2 O +1 2 .
With the aid of this reaction an important series of experiments to ascertain the connection between the con ditions of a chemical change and its amount have been made by Harcourt and Esson. The experiments consisted in adding successive equal portions of sodium hyposulphite to a solution containing hydrogen dioxide, hydriodic acid, and a little starch. By this reagent the iodine which is continually being liberated by the action of the dioxide on the hydriodic acid is instantly reconverted into iodide, so that the liquid, though it contain starch, and though iodine is being formed in it, remains quite colourless as long as any hyposulphite is present. But when the last trace of hyposulphite has been removed by the action of the iodine, the portion of iodine next formed remains free, and the liquid becomes suddenly blue. The addition of another small portion of hyposulphite again removes the colour, and until all the hyposulphite is decomposed the solution remains colourless, and then again becomes suddenly blue. The intervals at which the blue colour appeared were carefully noted, and the amount of hydrogen dioxide decomposed being known from the amount of hyposulphite employed, the quantity of dioxide decomposed in a given time was thus determined. The observed results are given in the following table :— [table]
The general conclusion deducible from these experiments is, that the amount of change at any moment varies directly with the amount of dioxide present in the solution; iu accordance with this law, the quantities of dioxide at the end of a series of times taken in arithmetical progression are themselves in geometrical progression. This law of chemical action has been corroborated by the investigation of other reactions, and it is probably of very general application.
When hydrogen dioxide solution is mixed with a concentrated solution of barium hydroxide, crystalline hydrated barium dioxide, BaO 2 + 6H 2 O, separates—
H 2 O 2 + Ba(OH) 2 = BaO 2 + 2H 2 O.
Hydrogen dioxide. Barium hydroxide. Barium dioxide. Water.
In a similar manner, peroxides of many metals are precipitated on the addition of their salts to a solution of hydrogen dioxide.
Hydrogen dioxide, it will be evident, differs remarkably from hydrogen monoxide or water. Its instability, and its tendency to enter into reaction with other bodies with separation of oxygen, appear to be explained by the fact that its decomposition into water and oxygen is attended with the development of a very considerable amount of heat. The behaviour of ozone and hydrogen dioxide, in fact, strikingly illustrates one of the most important laws of chemical action, viz., that those decompositions and reactions which are attended with the development of heat always take place more readily than those which require an absorption of heat, and they take place the more readily the greater the amount of heat which is liberated. The heat developed by the decomposition of ozone and hydrogen dioxide, we have seen, is to be traced to the same cause, being due, it can scarcely be doubted, to the combination of the oxygen atoms.
Fluorine — Chlorine — Bromine — Iodine. [table]
These four elements form with metals compounds analogous to sea salt, and from this circumstance the name halogens, or salt-producers (from αλs, sea-salt), has been applied to them, their compounds with other radicles being frequently termed haloid compounds. They are always classed together on account of their close analogy in properties, but there are numerous and very important distinctions between them.
The element fluorine is not known in the free state, all attempts to isolate it having failed on account of the impossibility of obtaining vessels which can withstand its action. Chlorine is a gas of a greenish yellow colour, whilst bromine, at ordinary atmospheric temperatures, is a mobile red liquid, so deep in colour as to be almost opaque, and iodine is a black, crystalline, and very brittle solid, which exhibits metallic lustre.
None of these elements are ever met with in the free state, but their compounds are very widely distributed, and they are to be detected in most rocks and soils, in spring and sea-water, and in the ashes of plants and animals. Fluorine occurs most abundantly in combination with calcium as fluor-spar, and chlorine in combination with sodium as ordinary salt, large deposits of which exist in various parts of the globe; considerable deposits of bromine in combination with potassium have within recent years been discovered in Stassfurt, but no abundant source of iodine has hitherto been discovered.
Chlorine was discovered by Scheele in 1774, and was so named on account of its colour (from χλωρός, green), but its elementary nature was first established by Davy in 1810. Bromine was first described in 1826 by Balard, who obtained it from bittern, the mother liquor of seawater, after the less soluble salts have been extracted by evaporation and crystallization; it was named on account of its irritating odour (from βρώμας, a stench). Iodine was discovered by Courtois in 1811, in the waste liquors from the manufacture of sodium carbonate from the ashes of sea-weed; it received its name from the beautiful violet colour of its vapour (ioeιδι)v, violet-coloured).
Chlorine is usually prepared, both in the laboratory and on the large scale, by gently heating a concentrated solution of hydrochloric acid with manganese dioxide; the reaction appears to take place in two stages,—the first consisting in the formation of the manganese chloride corresponding to manganese dioxide—
MnO 2 + 4HC1 = MnCl 4 + 2H 2 O;
Manganese Hydrochloric Manganese tetra- ∙ w ∙ ntp - dioxide. acid. chloride. water ∙
but this compound is so unstable that it breaks up into chlorine and a lower chloride of manganese—
MnCl 4 = Cl 2 + MnCl 2 .
Manganese tetra-Manganese chloride. Chlorine. dichl0ridβ .
It may be procured directly from salt by acting on a mixture of salt and manganese dioxide with sulphuric acid, but in this case hydrochloric acid is first produced by the action of the sulphuric acid on the salt—
NaCl + H 2 SO 4 = HC1 + NaHSO 4
The hydrochloric acid then enters into reaction with the manganese dioxide, as explained above, but the manganese dichloride is converted into manganese sulphate and hydrochloric acid by the sulphuric acid. The following equation therefore expresses the final result, but does not take into account that there are several stages in the reaction—
2NaCl + 3H 2 SO 4 + MnO 2 Sodium chloride. Sulphuric acid. Manganese dioxide.
= Cl 2 + 2OH 2 + MnSO 4 + 2NaHSO 4 .
Γhlnrin<* water Manganese Sodium hydrogen chlorine. Water. sulphate. sulphate.
Bromine and iodine are obtained in a precisely similar manner by treating the liquid containing the bromide or iodide with manganese dioxide and sulphuric acid.
Another process for the manufacture of chlorine has been recently introduced by Deacon, which consists in passing a mixture of hydrochloric acid gas and air through a heated brickwork chamber filled with pieces of firebrick or marble which have been soaked in a solution of copper sulphate and dried. The hydrogen of the hydrochloric acid forms water with the oxygen of the air, and chlorine is liberated, but the precise nature of the changes by which this is effected is not known. The reaction is at its maximum at a temperature of about 500° C., which is much below that at which oxygen alone acts upon hydrochloric acid to form water and chlorine; the copper sulphate is but little affected, although a certain amount is always transformed into chloride. There can be no doubt, however, that the decomposition is the result of recurrent action, and another somewhat similar method of obtaining chlorine may be quoted in illustration. Thus, when a mixture of silica and sodium chloride is heated to redness in a current of oxygen, chlorine is evolved, and the chloride is converted into silicate—
2SiO 2 + 4NaCl .+ O 2 = 2N τ a 2 SiO 3 + 2C1 2 ;
Silica. Sodium chloride. Oxygen. Sodium silicate, chlorine.
but on passing hydrochloric acid gas over the silicate it is re-converted into the chloride and silica—
Na 2 SiO 3 + 2HC1 = SiO 2 + 2NaCl + OH 2 Sodium silicate. Hydrochloric acid. Silica. Sodium chloride. Water.
Hence, by passing hydrochloric acid together with air over the mixture of silica and chloride, a continuous evolution of chlorine is obtained, reaction in the one direction having no sooner taken place than reaction in the opposite direction sets in. There can be little doubt that the production of chlorine by Deacon’s process is, in a similar manner, the result of a series of changes in opposite directions, the nature of which, however, remains to be discovered.
Chlorine is very readily condensed by a pressure of about 4 atmospheres at 15° C., or by cold alone at about — 50° C., into a yellow mobile liquid of specific gravity about 1 ∙33, but it has not been solidified. Bromine has at 0° C. the specific gravity 3T 88; it boils at 59 o, 5 C., yielding a dense red vapour, and solidifies at - 24° ∙5 C. Iodine melts at 107° C., and boils at about 175° C., yielding a magnificent purple vapour. Chlorine gas and the vapour of bromine and iodine have a pungent, irritating, disagreeable odour, and are irrespirable. Water at 10° C. dissolves about 2‘5 times its bulk of chlorine, and about 3 per cent, of bromine, but iodine is very sparingly soluble in water. "When chlorine is passed into a mixture of crushed ice and water a crystalline hydrate, Cl 2 + 10H 2 O, is produced, and a similar compound of bromine, Br 2 +10H 2 O, may be obtained, but iodine does not furnish a hydrate.
The halogens furnish with hydrogen compounds in which a single atom of hydrogen is united with a single atom of the halogen, viz. :— [table]
The conditions under which hydrogen and chlorine enter into reaction have already been described, and we have seen that no change of volume occurs in the formation of hydrogen chloride from its elements. A mixture of hydrogen and bromine vapour does not explode on the application of flame, but hydrobromic acid is slowly formed when the mixture is heated; hydriodic acid is also produced when iodine is heated in hydrogen. Gore has shown that hydrofluoric acid may be prepared by heating silver fluoride in an atmosphere of hydrogen—
2AgF + H 2 = 2HF + 2Ag; Silver fluoride. Hydrogen. Hydrofluoric acid. Silver, and he finds that one volume of hydrogen in uniting with fluorine produces two volumes of hydrofluoric acid gas, thus proving that the constitution of hydrogen fluoride is similar to that of hydrogen chloride, bromide, and iodide. Hydrogen chloride or hydrochloric acid is usually prepared by gently heating a mixture of sodium chloride or common salt and concentrated sulphuric acid, diluted with a small quantity of water—
NaCl + H 2 SO 4 = HC1 + NaHSO 4 .
Sodium chloride. Sulphuric acid. Hydrogen chloride. Sodium hydrogen sulphate Hydrogen bromide and iodide cannot be prepared in this manner from metallic bromides and iodides, as they are partially decomposed by concentrated sulphuric acid, bromine or iodine being liberated; this decomposition takes place much more readily with hydriodic acid than with hydrobromic acid—
2HI + H 2 SO 4 = I 2 + SO 2 + 2OH 2 .
Hydrogen iodide. Sulphuric acid. Iodine. Sulphur dioxide. Water.
Hydrobromic and hydriodic acid are therefore usually prepared by the action of water on the compounds of phosphorus with bromine and iodine; thus—
PBr 3 + 3H 2 O = 3HBr + PH 3 O 3 Phosphorus tribromide. Water. Hydrogen bromide. Phosphorous acid.
PI 3 + 3H 2 O = 3HI + PH 3 O 3 .
Phosphorus triiodide. Water. Hydrogen iodide. Phosphorous acid.
Hydrogen fluoride is procured by gently heating finely powdered calcium fluoride or fluor-spar with concentrated sulphuric acid in leaden or platinum vessels, or by strongly heating potassium hydrogen fluoride in a platinum re tort—
CaF 2 + H 2 SO 4 = 2HF + CaSO 4 Calcium fluoride. Sulphuric acid. Hydrogen fluoride, calcium sulphate.
HKF 2 = HF + KF.
Potassium hydrogen fluoride. Hydrogen fluoride. Potassium fluoride.
The compounds of hydrogen with the halogens, or, as they are frequently termed, the haloid acids, are not only similar in composition, but they resemble each other very closely in properties. Thus, hydrogen chloride, bromide, and iodide at ordinary atmospheric temperatures are colourless transparent gases, which fume strongly in moist air, whilst hydrogen fluoride is a highly volatile colourless liquid, of specific gravity ’ 988 at 12 0, 7 C., which boils at 19° ∙4 C.; its vapour fumes strongly in moist air. At a temperature of 10° C., under a pressure of 40 atmospheres, hydrogen chloride is condensed to a colourless liquid of specific gravity 1 *27, but it has never been solidified; hydro-, gen bromide and iodide are still more readily liquefied, and both maybe caused to solidify, the former at - 73° C., and the latter at - 51° C.
The circumstances which attend the formation of hydrogen chloride, bromide, and iodide from their elements, and.
their general behaviour, show that the affinity of chlorine for hydrogen is much greater than that of bromine, whilst that of bromine is greater than that of iodine; and as already pointed out (p. 475), a considerable amount of heat is developed in the formation of hydrogen chloride, but a less amount in the formation of hydrogen bromide, and the production of hydrogen iodide from its elements involves the absorption of heat. We have seen, however, that the heat developed or absorbed in the formation of these compounds cannot be regarded as an absolute measure of the affinity of chlorine, bromine, and iodine for hydrogen, since there is an expenditure of energy due to the change of state, the three compounds being gases, but one being formed from two gases, the second from a gas and a liquid, and the third from a gas and a solid. Moreover, as we believe that the molecules which enter into reaction consist each of two atoms, the combination of the dissimilar atoms to form molecules of hydrogen chloride, bromide, and iodide must be preceded by the separation of the similar atoms which constitute the molecules of hydrogen, and of chlorine, bromine, and iodine, and there must be an expenditure of energy to effect this separation; it appears probable that different amounts of energy are expended in separating the atoms of chlorine, of bromine, and of iodine from each other, and we may further assume as probable that the amount of energy expended in the separation of the atoms is greater in the case of chlorine than of bromine, and greater in the case of bromine than of iodine, although at present there is scarcely any experimental evidence which enables us to draw conclusions of any value in proof of this.
Hydrogen fluoride is miscible with water in all proportions, and much heat is developed on mixing the two substances; hydrogen chloride, bromide, and iodide are also extremely soluble in water, forming colourless strongly acid solutions. Water at 0° C. absorbs ’ 825 of its own weight, or about 492 times its bulk, of hydrogen chloride gas, increasing in volume about one-third, and acquiring a density of about 1’23. This saturated solution consists of hydrogen chloride and water in about the proportions indicated by the formula HC1 + 3H 2 O; it fumes strongly, and when heated it gives off hydrogen chloride until at 112° C. a more dilute solution distils over, corresponding approximately in composition with the formula HC1 + 8H 2 O. A weaker solution when heated parts with water until it acquires this composition, and then distils unchanged. In like manner, a solution containing about 48 per cent, of hydrogen bromide distils unchanged at 125° C., and a solution containing about 57 per cent, of hydrogen iodide boils constantly at 127° C.; these solutions correspond approximately in composition with the formulae HBr + 5H 2 O and 2HI + 11H 2 O. A solution of hydrogen fluoride approximately of the composition IIF + 2H 2 O distils unchanged at 120° C. Roscoe’s experiments have shown, however, that the composition of these solutions varies with the pressure, and that there exists for each pressure a solution of corresponding strength which undergoes no change in composition when distilled under this pressure, and therefore has a constant boiling point. Hence, the apparent constancy of composition cannot be regarded as proof of the existence of a definite hydrate corresponding in composition to that of the solutions obtained on distilling under a given pressure. There is no evidence, however, to invalidate, but much to support the conclusion that the haloid acids do not dissolve in water as such, but that they form definite hydrates; and if it be admitted that their solutions contain definite hydrates, their behaviour is precisely similar to that of other compounds which decompose when heated, but furnish products which can reunite to form the parent compound.
According to Thomsen’s determinations, the amounts of heat developed when gaseous hydrochloric, hydrobromic, and hydriodic acids are dissolved in water, in the proportion of 1 molecule of the acid to 400 molecules of water, are as follows—
In the dissolution of hydrogen chloride, 17,314 heat-units. [table]
The development of so large an amount of heat, although, doubtless, in a great measure arising from the change from the gaseous to the liquid state, may be regarded as evidence that the acids really enter into combination with water, and are not merely dissolved; it may be noticed also that the same amounts of heat are developed in dissolving hydrogen bromide and iodide, and that their solutions, which under the ordinary atmospheric pressure distil unchanged, have nearly the same boiling point, and correspond approximately in composition.
Solutions of hydrogen chloride and bromide may be preserved unchanged, but the solution of hydrogen iodide gradually becomes brown, especially on exposure to light, owing to the separation of iodine, the hydrogen being oxidized by the oxygen of the air—
4HI + O 2 = 2I 2 + 2OH 2 .
Hydrogen iodide. Oxygen. Iodine. Water.
On the other hand, solutions of chlorine or bromine in water may be preserved unchanged in the dark, but when they are exposed to light oxygen is evolved, and hydrogen chloride or bromide produced—
2C1 2 + 2H 2 O = O 2 + 4HC1.
chlorine. water. Oxygen. Hydrogen chloride.
A solution of iodine does not undergo alteration. In the presence of substances which have a tendency to enter into reaction with oxygen, water is often readily decomposed in this manner by chlorine and bromine without the aid of light, but iodine seldom effects this change; thus, on passing chlorine into a solution of sulphurous acid, the latter is oxidized to sulphuric acid—
H 2 SO 3 + H 2 O + Cl 2 = H 2 SO 4 + 2HC1. Sulphurous acid. Water. Chlorine. Sulphuric acid. Hydrogen chloride. The bleaching action which is powerfully exerted by chlorine in presence of water, and less powerfully by bromine, is in most cases the result of a similar change, the substance bleached being usually one which has a tendency to enter into reaction with oxygen.
Chlorine readily displaces bromine and iodine from their compounds with hydrogen and other elements, and, similarly, bromine displaces iodine; indeed the affinity of chlorine for other elements, with few exceptions, is superior to that of bromine, the affinity of bromine being in like manner superior to that of iodine. The decomposition of many chlorine compounds by the corresponding bromine and iodine compounds, and of bromine compounds by the corresponding iodine compounds, would appear at first sight to be in contradiction to this statement, but on examination of all such reactions it is found that they involve a development of heat, and therefore are perfectly normal in character. For example, silver chloride is converted into silver iodide by digestion with an aqueous solution of hydriodic acid—
AgCl + HI = AgI + HCl.
It will be obvious that, in a molecule of silver iodide and a molecule of hydrogen chloride, the same quantity of the same elements is present as in a molecule of silver chloride and of a molecule of hydrogen iodide. But in the formation of the latter system from its elements, less heat is developed than in the formation of the former; hence reaction in the direction indicated by the above equation is attended with the development of heat. Thus, in the formation of a molecule of dissolved, hydriodic acid from its elements 13,170 units of heat are evolved, and about 34,800 units of heat are developed in the production of a molecule of silver chloride; whereas 39,320 units of heat are developed in the formation of a molecule of dissolved hydrochloric acid from its elements, and about 18,650 units in the production of a molecule of silver iodide. But —
(34,800 +13,170) - (18,650 + 39,320) = - 10,000 that is to say, the action of hydriodic acid on silver chloride is attended with the development of no less than 10,000 units of heat.
All metals which decompose water at a red heat or at lower temperatures decompose a solution of hydrogen chloride with evolution of hydrogen, and form the corresponding metallic chloride, but with very different degrees of readiness; thus—
Zn + 2HC1 = H 2 + ZnCl 2 .
Zinc. Hydrogen chloride. Hydrogen. Zinc chloride.
According to Thomsen, the amounts of heat developed in the production of 2 grammes of hydrogen by the action of various metals on a dilute solution of hydrogen chloride are as follows :— [table]
Exact comparative observations of the degrees of readiness with which the various metals evolve hydrogen from a solution of hydrogen chloride have not as yet been made, but from ordinary observations it appears that those act most readily which develop the greatest amount of heat when dissolved. It is very difficult to compare the behaviour of different metals, however, since minute quantities of impurity exert a most remarkable influence; thus, ordinary zinc and iron dissolve with the greatest readiness in a dilute solution of hydrogen chloride, but the pure metals are only slowly dissolved.
Hydrogen bromide and iodide closely resemble hydrogen chloride in their behaviour with metals; mercury, however, which is not affected by hydrogen chloride or bromide, slowly dissolves in hydrogen iodide. Hydrogen fluoride not only acts with great readiness upon all metals which are dissolved by the other haloid acids, but likewise on many which these acids do not attack,—copper and silver, for example; and it also dissolves many elements which are insoluble in all other acids, such as silicon, boron, titanium, tantalum, and zirconium, with evolution of hydrogen. Gore has shown that anhydrous liquid hydrogen fluoride and chloride are, in most cases, much less energetic in their chemical reactions than their aqueous solutions; this is probably because the products which are formed in the first instance, being insoluble or difficultly soluble in the acid, form a coating on the surface of the substance submitted to the action of the acid, and thus preserve it from further action, or cause the action to take place more slowly; when water is present the product is dissolved, and thus a fresh surface is continually exposed.
The haloid acids readily enter into reaction with most metallic oxides, forming water and the corresponding compound of the metal with the halogen; thus—
ZnO + 2HC1 = ZnCl 2 + OH 2
Zinc oxide. Hydrogen chloride. Zinc chloride. Water. Fe 2 O 3 + 6HC1 = Fe 2 Cl 6 + 3OH 2 .
Ferric oxide. Hydrogen chloride. Ferric chloride. Water.
Hydrogen fluoride is capable of dissolving some oxides, such as silica, SiO 2 , for example, which are not affected by the other haloid acids. It is in consequence easily distinguished from the other haloid acids by the powerful corrosive action which it exerts on glass.
The haloid acids also readily enter into reaction with, or neutralize, the hydroxides, forming water and the corresponding fluoride, chloride, bromide, or iodide—
NaOH + HC1 = NaCl + H 2 O;
Sodium hydroxide. Hydrogen chloride. Sodium chloride. Water
Ba(OH) 2 + 2HBr = BaBr 2 + 2H 2 O.
Barium hydroxide. Hydrogen bromide. Barium bromide. Water. The thermochemical behaviour of hydrochloric acid with hydroxides has already been discussed.
Chlorine enters into reaction with many of the elements at ordinary atmospheric temperatures, and with all of them, with few exceptions, at more or less elevated temperatures; its union with phosphorus, and with finely divided arsenic, antimony, tin, and copper, is attended with combustion. Bromine and iodine closely resemble chlorine in their behaviour with other elements; the union of bromine with other elements, however, appears always to take place less readily, and to be accompanied by the development of less heat than is the case with chlorine, and the affinity of iodine for most other elements is still weaker than that of bromine. The iodides also are usually less stable than the corresponding bromides, which are usually less stable than the corresponding chlorides. The compounds of the halogens with metals and with positive elements generally are mostly stable in presence of water, but their compounds with negative elements are mostly decomposed by water, and the halogen separated in the form of haloid acid.
The same element does not always enter into combination with chlorine, bromine, and iodine in the same proportions; thus, phosphorus forms a stable trichloride, PC1 3 , and an unstable pentachloride, PC1 5 ; a stable tribromide, PBr 3 , and a pentabromide, PBr 5 , less stable than the pentachloride; but with iodine it forms the compounds PI 3 and P 2 I 4 . Many elements furnish compounds with fluorine containing a higher proportion of fluorine than corresponds to the amount of chlorine in their highest chlorides, or they form stable fluorides but unstable chlorides, bromides, and iodides; for example, phosphorus pentafluoride, PF 5 , is not decomposed by heat, whereas the pentachloride, PC1 5 , is readily resolved into the trichloride and chlorine by heating.
The halogens do not only combine with hydrogen and other elements, however, but they are capable of forming compounds with each other. Thus, by the action of iodine on silver fluoride, Gore has obtained iodine pentafluoride, IF 5 , as a colourless highly volatile liquid; iodine at once enters into reaction with chlorine, and is first converted into liquid iodine chloride, ICI, but ultimately yields iodine trichloride, IC1 3 , which is a crystalline solid; with bromine iodine appears to furnish only a liquid monobromide, IBr. In like manner, bromine forms with chlorine a liquid bromine chloride, BrCl.
On account of its great affinity for hydrogen, chlorine readily enters into reaction with a large number of hydrogenized carbon compounds, and displaces the hydrogen more or less completely. Thus, methane, CH 4 , when submitted to the action of chlorine, is successively converted into monochloromethane, CH 3 C1, dichloromethane, CH 2 C1 2 , trichloromethane, CHC1 3 , and tetrachloromethane, CC1 4 ,—the displaced hydrogen being evolved as hydrogen chloride. Bromine behaves similarly, but enters into such reactions much less readily; iodine seldom, if ever, behaves in this way.
Compounds of the Halogens with Oxygen.
Oxides of fluorine are unknown, and the oxides of the remaining halogens are very imperfectly investigated. Three oxides of chlorine have been prepared, viz. :—
C l 2 O, C1 2 O 3 , and C1O. 2 Hypochlorous chlorous chlorine anhydride. anhydride. peroxide.
At ordinary temperatures these compounds are greenish-yellow gases, remarkable on account of their instability, being decomposed with explosive violence by very slight elevations of temperature. Their instability is accounted for by the circumstance that their decomposition is attended with the development of a very considerable amount of heat; thus, according to Thomsen, the formation of gaseous hypochlorous anhydride from its elements would involve the absorption of 18,040 units of heat per molecule, and consequently this amount is developed when the anhydride is resolved into its elements. The explanation of the absorption of heat in the formation of the oxides of chlorine would appear to be that more energy must be expended in separating the atoms of oxygen from each other, and in separating the atoms of chlorine from each other, than is produced by the combination of the chlorine and oxygen atoms.
Oxides of bromine are also unknown, and since thermochemical investigation has shown that oxygen has less affinity to bromine than to chlorine, their non-formation in reactions corresponding to those by which the oxides of chlorine are produced is no longer surprising.
Iodine differs entirely from chlorine and bromine with regard to its affinity for oxygen. Only one oxide of iodine is known, viz., iodic anhydride, I 2 O 5 , but this is a white crystalline substance of considerable stability, and, according to Thomsen’s determination, the formation of this compound from its elements would involve the development of no less than 44,860 units of heat per molecule. It has not been directly produced from its elements, although its formation in this manner is not improbable.
Chlorous anhydride is prepared by passing dry chlorine gas over dry mercuric oxide—
2HgO + 2C1 2 = C1 2 O + Hg 2 OCl 2 .
It may be condensed by surrounding the receiver with a mixture of ice and salt, aud thus obtained as a deep red liquid which emits a vapour of a deeper colour than that of chlorine, having a peculiar suffocating chlorous smell. It is with the greatest readiness decomposed into chlorine and oxygen, the warmth of the hand being sufficient to cause it to explode. Water dissolves about 200 times its bulk of the gas, forming a pale yellow solution of hypochlorous acid, which has an acid but not sour taste—
C 1 2 O + H 2 O = 2HOC1.
By neutralizing this solution with metallic hydroxides or carbonates corresponding metallic hypochlorites are formed, but these salts are usually prepared by passing chlorine into water in which the metallic hydroxide or carbonate is suspended or dissolved, the liquid being carefully cooled—
2KOH + Cl 2 = KOC1 + KCΙ + OH 2 .
Potassium r ,∙, Potassium Potassium hydroxide. Chlorine. hypochlorlte . ch ] O ride. Water
A concentrated solution of hypochlorous acid decomposes rapidly, even when kept in ice; a dilute solution is more stable, but is decomposed by boiling into chloric acid, water, chlorine, and oxygen. These substances are probably the products of several distinct changes, such as are represented by the following equations—
2HOC1 = O 2 + 2HC1
3HOC1 = HO 3 C1 + 2HC1.
Hypochlorous acid. chloric acid. Hydrogen chloride.
Hypochlorous acid, however, readily decomposes hydrochloric acid, and the chlorine produced is probably the result uf reaction between these bodies—
HC1 + HOC1 = C1 2 + OH 2 .
Hypochlorous acid is a very powerful oxidizing agent, and rapidly converts many of the elements into their highest oxides, at the same time liberating chlorine; indeed, in many respects it is not unlike ozone and hydrogen dioxide in its chemical behaviour, and it enters into reaction with the latter in the manner represented by the equation—
HOC1 + H 2 O 2 = O 2 + HC1 + H 2 O.
Like chlorine, it possesses powerful bleaching properties. The hypochlorites also act as bleaching agents, probably by yielding oxygen to the substance submitted to their action, but less powerfully than the free acid.
The bleaching powder or chloride of lime of commerce is prepared by exposing slaked lime or calcium hydroxide to the action of chlorine gas. There has been much discussion among chemists as to the constitution of the compound so produced, and it is only quite recently that evidence has been obtained which enables us to pronounce a decision. The composition of bleaching powder is represented by the formula CaCl 2 O, whereas calcium hypochlorite would be represented by the formula CaCl 2 O 2 j it was therefore supposed that the bleaching powder was a mixture of calcium chloride and calcium hypochlorite, since CaCl 2 + CaCl 2 O 2 = 2CaCl 2 O. Apparently, however, it is a distinct compound intermediate between calcium chloride and calcium hypochlorite, thus—
Caí C1 CaJ C1 Ca!(° ci >∙
ca iα ’ 1 (Oci) , υa {(OCi) ’ calcium chloride cakittm chlorohypochlorite. calcium calcium cnιonde. (Bleaching powder). hypochlorite.
and its formation from calcium hydroxide may be represented by the following equation—
Ca{θ+c⅝=Ca{α cv + °H 2 .
On dissolving in water, according to Kingzett, it is resolved into calcium chloride and calcium hypochlorite, which may be obtained in crystals by carefully concentrating the solution in vacuo, so that a solution of bleaching powder is correctly regarded as a solution of calcium hypochlorite—
20a∣θ 1 c 1 = CaCl 2 + Cajθ^}.
Bleaching powder, calcium chloride. Calcium hypochlorite.
Solutions of hypochlorites are very unstable, and when boiled they furnish a mixture of chloride and chlorate; thus—
3KOC1 ≈ KO 3 C1 + 2KC1.
Potassium hypochlorite. Potassium chlorate. Potassium chloride.
It is on this account that it is necessary in preparing hypochlorites to act on a cold solution of the hydroxide, since if the solution be heated chlorate is formed by the decomposition of the hypochlorite.
The chlorates are very stable compounds as compared with the hypochlorites. They may be obtained well crystallized, and are all soluble in water. They are destitute of bleaching properties. The most important chlorate is potassium chlorate, KC1O 3 , which is chiefly interesting as a source of oxygen, and is largely employed in the manufacture of fireworks; the resolution of this salt into oxygen and potassium chloride is attended with a development of heat amounting to 9700 units per molecule decomposed.
By exactly decomposing a solution of barium chlorate with sulphuric acid, a solution of chloric acid, HC1O 3 , is obtained, which may be separated from the insoluble barium sulphate by filtration, and concentrated by evaporation in vacuo until it forms a syrupy liquid :—
Ba(ClO 3 ) 2 + H 2 SO 4 = 2HC1O 3 + BaSO 4 .
Chloric acid has a faint chlorous odour, and a powerfully acid reaction; it is instantly decomposed by contact with organic matter, with charring, and frequently even with ignition. In diffused daylight it gradually undergoes spontaneous decomposition, and when heated to a little above 100° C. it is rapidly converted into perchloric acid, water, chlorine, and oxygen—
8HC1O 3 = 4HC1O 4 + 2H 2 O + 3O 2 + 2C1 2 .
Chloric acid. Perchloric acid. Water. Oxygen, chlorine.
It is a powerful oxidizing and bleaching agent.
Chloric acid serves for the preparation of the remaining oxides and oxy-acids of chlorine. To prepare chlorous anhydride from it, 10 parts of pure benzene are dissolved in 100 parts of concentrated sulphuric acid, and when the mixture has cooled 12 parts of pulverized potassium chlorate are added. The chloric acid liberated from the chlorate under these circumstances is deprived of a portion of its oxygen by the benzene, which is oxidized, and at the same time, by the dehydrating influence exercised by the sulphuric acid, the elements of a molecule of water are removed and chlorous anhydride is liberated; thus— 2HC1O, - O., - OIL = Cl.,O.,.
The mixture is heated to about 50° C. in a flask having a delivery tube fitted to its neck by grinding, and blown out in several places into small bulbs; the gas, after passing through water contained in these bulbs, is conducted into a glass tube surrounded by a mixture of ice and salt. The condensed liquid deposits crystals of a hydrate of chlorous acid, and the supernatant liquid decanted immediately into a small cooled glass cylinder forms the nearly anhydrous oxide C1 2 O 3 . Liquid chlorous oxide or anhydride thus obtained is very mobile, is of a deep red colour, and has a specific gravity of 1*330 to 1 ’ 387 at 0° C.; it boils a little above 0° C., and it explodes somewhat readily at a few degrees above 0° C. The specific gravity of the gaseous oxide accords with that required by the formula C1 2 O 3 . Chlorous anhydride is not very soluble in water, 100 grammes of water at 23° C. dissolving 5 ’ 65 grammes of the oxide (Brandau). The solution of chlorous acid, HC1O 2 , thus obtained, has powerful bleaching and oxidizing properties; when it is heated, chloric and hydrochloric acids are formed and chlorine evolved. By neutralizing the solution with hydroxides of the alkali or alkaline earth metals corresponding chlorites are obtained, which are soluble; and from these the insoluble chlorites of lead and silver are obtained by double decomposition. Chlorites are decomposed by the feeblest acids, even by carbonic acid.
When fused potassium chlorate is carefully treated with concentrated sulphuric acid, chlorine dioxide or peroxide is evolved; the reaction by which it is produced, it is stated, is represented by the following equation—
3KC1O 3 + 2H 2 SO 4 = 2C1O 2 + H 2 O + KC1O 4 + 2KHSO 4 .
Potassium Sulphuric chlorine w ater Potassium Potassium chlorate. acid. peroxide. * perchlorate, hydrogen sulphate.
The so-called cuchlorine gas, which is obtained when a mixture of a chlorate and hydrochloric acid is gently heated, appears to be a mixture of chlorine peroxide with chlorine. Chlorine peroxide is gaseous at ordinary temperatures, but by means of a mixture of ice and salt it may be condensed to a red liquid. It has a deep greenish-yellow colour, and peculiar sweet chlorous odour; a slight degree of heat and mere contact with organic matters at once determine its explosion. Like all the oxides of chlorine, it acts rapidly upon mercury and most metals. From Pebal’s recent experiments there appears to be no doubt that chlorine peroxide is correctly represented by the formula C1O 2 , and it is, therefore, to be included with nitric oxide and the few other compounds in which we must suppose an odd number of affinities are free (see p. 473). It is absorbed by alkaline solutions, forming a mixture of a chlorite and a chlorate—
2C1O 2 + 2KHO = KC1O 2 + KC1O 3 + OH 2 .
chlorine Potassium Potassium Potassium yr a t er peroxide. hydroxide. chloride. chiorate.
In this respect especially it differs from the other oxides of chlorine.
The only remaining chlorine compound to be considered is perchloric acid, HC1O 4 . It may be formed from chloric acid by oxidation, but, according to Roscoe, the best method of preparing it consists in boiling down a solution of chloric acid; lower oxides of chlorine escape, and a solution of perchloric acid is left. It may also be obtained from potassium perchlorate, farmed by heating potassium chlorate, by distilling it with sulphuric acid.
Aqueous perchloric acid may be concentrated by boiling till it attains a temperature of 203° C., after which it passes over in the form of an oily liquid which approximately has the composition HC1O 4 + 2H 2 O. By distilling this hydrate with twice its volume of sulphuric acid, nearly pure perchloric acid is obtained.
The pure acid is a colourless very soluble liquid, of specific gravity 1 , 782 at 15° '5 C. In this state it is one of the most powerful oxidizing agents known; a single drop brought in contact with charcoal, wood, or almost any organic substance, immediately causes an explosive combustion, which in violence almost equals the sudden decomposition of the so-called chloride of nitrogen. It produces frightful burns if allowed to fall upon the skin. It undergoes spontaneous decomposition at ordinary temperatures, and cannot be distilled unchanged. It unites very energetically with water, and when mixed with it in suitable proportions forms the hydrate HC1O 4 + H 2 O; this is a white crystalline substance, which melts at 50° C∙, and undergoes decomposition when heated to 110° C., splitting up into the pure acid and the oily hydrate above mentioned.
Aqueous perchloric acid has a sour taste, and reddens litmus strongly, but does not bleach; it dissolves iron and zinc with evolution of hydrogen, and when dilute it is unaffected by hydrogen sulphide and sulphurous acid, which reduce all other oxy-acids of chlorine. Perchloric acid, in fact, is distinguished from all other oxidized chlorine compounds by its superior stability, oxygen being withdrawn from it with comparative difficulty.
The results of Thomsen’s thermochemical examination of various reactions which are involved in the formation of some of the oxidized chlorine compounds are exhibited in the following table. The symbols in the first column indicate the nature of the reaction, and the numbers in the second column the units of heat developed or absorbed,—the reaction being supposed to take place between the substances of which the symbols are separated by commas. Thus, the ^ ne Cl 2 , O - 18,040 Gaseous oxide is to be read as meaning that, in the formation of a molecule (in grammes) of gaseous chlorine oxide from chlorine and oxygen, 18,040 units of heat would be absorbed Similarly, the reaction C1 2 O, Aq, that is to say, the absorption of gaseous chlorine oxide by water is attended with the development of 9440 units of heat; hence the reaction Cl 2 , O, Aq, or the formation of a solution of hypochlorous acid (two molecules) from chlorine, oxygen, and water would involve the absorption of 8600 units of heat. For these, as in all Thomsen’s experiments, the determinations were made at about 18° C.
[table]
In this table, and in all similar tables, the reactions are represented empirically, and the figures after the symbols merely indicate the number of atoms which enter into the composition of the compound. Thus, the reaction expressed as Cl s , 0, merely refers to the formation of a molecule (in grammes) of hypochlorous anhydride, C1 2 O, from chlorine and oxygen, and the arrangement of the symbols in this manner is not intended to indicate that the compound is formed from a molecule qì chlorine and an atom of oxygen. The symbol Aq is intended to indicate that an excess of water is employed, ι.e., such an amount that no appreciable development of heat would be caused by the addition of a further quantity.
When bromine is added to an aqueous solution of sodium hydroxide, sodium hypobromite, NaOBr, is formed, but this is much less stable than the corresponding hypochlorite, and is readily converted into bromate and bromide. The bromates and bromic acid resemble the chlorates and chloric acid, but are much less stable; thus, a solution of bromic acid is decomposed at 100° C., giving off bromine and oxygen, and all reducing agents decompose it with facility. The inferior stability of the oxidized compounds of bromine as compared with those of chlorine appears to indicate that the affinity of bromine to oxygen is less than that of chlorine; this is entirely confirmed by thermochemical investigation, Thomsen having obtained the following values for bromic acid :— [table]
When iodine is added to solutions of alkaline hydroxides or carbonates, alkaline hypoiodites corresponding to the hypochlorites are perhaps formed, but as yet have not been isolated. By dissolving iodine in a warm solution of potassium hydroxide a mixture of potassium iodide and iodate is produced— [table]
Iodine is converted into iodic acid by heating with the strongest nitric acid, and by the action of many other oxidizing agents; but iodic acid is usually prepared from barium iodate, which is decomposed with the necessary quantity of sulphuric acid. Potassium iodate is most readily prepared by passing chlorine into water in which iodine is suspended until the latter is dissolved, then adding a corresponding quantity of potassium chlorate and warming; the iodine is converted into monochloride, which enters into reaction with the chlorate, and chlorine is evolved—
KC1O 3 + IC1 = KIO 3 + C1 2 .
By adding a solution of barium hydroxide, insoluble barium iodate is precipitated, from which, as pointed out, iodic acid may be prepared.
Iodic acid separates from its solution on spontaneous evaporation in crystals of the composition HIO 3 ; when heated to 120° C. these crystals lose water, and apparently form the compound HI 3 O 8 or HIO 3 + I 2 O 5 ; at about 170° C. this compound is resolved into water and iodic anhydride, I 2 O 5 . Iodic anhydride is a crystalline substance, which readily dissolves in water, forming iodic acid; when heated to about 370° C. it is resolved into oxygen and iodine. Iodic acid is a powerful oxidizing agent, being easily decomposed by deoxidizing agents; hydriodic acid reduces it to iodine and water—
HIO 3 + 5HI = 3I 2 + 3H 2 O.
The chlorates and bromates are mostly easily soluble, and the corresponding salts are isomorphous; but the iodates are mostly difficultly soluble, and with the exception of barium iodate, are not isomorphous with the corresponding chlorates and bromates. Moreover, chloric and bromic acids furnish but one class of salts, of which potassium chlorate, KC1O 3 , is typical; but iodic acid, in addition to the normal iodates, such as potassium iodate, KIO 3 , forms acid salts. These differences appear to indicate that iodic acid differs in constitution from the other acids, and Thomsen has pointed out that it is probably a dibasic acid of the formula H 2 I 2 O c , since it is readily converted into water and the anhydride I 2 O 5 , a property which is characteristic of dibasic but not of monobasic acids; and since it furnishes acid salts, which is also characteristic of polybasic but not of monobasic acids, and is isomorphous with succinic acid, which is dibasic. Normal potassium iodate would therefore have the formula K 2 I 2 O 6 , and the so-called diiodate is probably the acid salt KHI 2 O 6 .
This view of the constitution of iodic acid is rendered the more probable by the behaviour of periodic acid, the disodium salt of which separates as a crystalline powder on passing chlorine into a solution of sodium iodate and sodium hydroxide; the reaction is empirically represented by the equation—
NaIO 3 + 3NaOH + Cl 2 = Na 2 TH 3 O 6 + 2NaCl.
Sodium iodate. Disodium periodate.
To prepare the acid from this salt, it is dissolved in nitric acid, and silver nitrate added; the precipitate of disilver periodate is dissolved in hot dilute nitric acid, and the solution concentrated at a moderate heat until a salt of the composition AgIO 4 crystallizes out. By treating this salt with cold water it is decomposed, disilver periodate a∏(Z periodic acid being produced :—
2 AgIO 4 + 4OH 2 = Ag 2 IH 3 O 6 + H 5 IO 6 ; to obtain the latter the solution is evaporated.
Periodic acid separates from its aqueous solution in colourless crystals of the composition H 5 IO 6 , which are not altered at 100° C., but melt at 130° C., and when heated to a higher temperature give off water and oxygen, leaving iodic anhydride.
Periodic acid furnishes a complex but extremely interesting series of salts. Thus, it forms two potassium salts, one of which has the composition KIO 4 , and is obtained on passing chlorine into a solution of potassium iodate and potassium hydroxide; the second is represented by the formula K 4 I 2 O 7 , and is obtained on dissolving the former in potassium hydroxide solution. The latter crystallizes with 9 molecules of water, which, however, it loses when exposed over sulphuric acid. Corresponding sodium salts exist. That which has the composition NaIO 4 crystallizes either in the anhydrous state, or with 3 molecules of water, which it readily loses in dry air; the second salt, the formation of which was above described, separates in crystals of the composition Na 4 I 2 O 9 + 2H 2 O, and is only rendered anhydrous by heating to 220° C. Three silver salts are known, represented by the formulae AgIO 4 , Ag 4 I 2 O 9 + 3H 2 O, and Ag 5 IO 6 ; the second of these loses 2 molecules of water at 100° C., but the third only at 200° C. The barium salt Ba 2 I 2 O 9 separates in the anhydrous state from strongly acid solutions, but when prepared by precipitating a solution of the acid with barium hydroxide, or of an alkali salt with a barium salt, it contains 6 or 7 molecules of water, which are only completely removed by heating to 300° C. It also furnishes a lead salt of the composition Pb 3 I 2 O 10 + 2H 2 O, and amongst other magnesium salts one of the composition Mg 4 I 2 O 11 , which crystallizes with 6 and with 9 molecules of water.
The question now arises—What is the nature of the relation between these various 'salts? By the chemical method alone it is extremely difficult, if not impossible, to decide, but from Thomsen’s thermochemical investigation of the acid there can be little doubt as to the answer we should make to this question. When successive molecules of potassium hydroxide are added to a solution of 1 molecule of the acid H 5 IO 6 the amounts of heat evolved are as follows :— [table]
The first and second molecules, therefore, together cause the development of 26,590 units of heat, or of 13,295 units per molecule. But we have seen (p. 486) that on neutralizing a large number of ι acids, between 13,750 and 13,150 units of heat are developed per molecule of hydroxide added; and as the addition of further quantities of the hydroxide causes a comparatively slight development of heat there is little doubt from these results that the molecule H 5 IO 6 is dibasic. But on account of the existence of salts, such as K 4 I 2 O 9 , it appears desirable to double this formula, and to represent the molecule of periodic acid by the formula H 4 .1 2 O 12 H 6 . Several of the salts above alluded to may be regarded as derived from this molecule by the partial or total displacement of the hydrogen by metals; and we may term those which are formed by displacing 4 of the 10 atoms of hydrogen normal salts, while those in which 2, 4, or 6 of the remaining atoms of hydrogen are displaced may be called basic salts. Thus we have—
Normal wits.
Na 4 .1 2 O 12 H 6 ; Ag 4 .1 2 O 12 H 6
Basic salts.
Pb 2 ". I 2 O 12 H 4 Pb"i Mg 2 ". I 2 O 12 H 2 Mg 2 " 5 Ag 4 '. I 2 O 12 Ag 6 '. That the tendency to form basic salts is slight in the case of the highly positive elements is also shown by the small amount of heat developed on the addition of the third, fourth, and fifth molecules of potassium hydroxide.
The anhydrous salts such as K 4 I 2 O 9 are to be regarded as derived from a distinct acid formed from the molecule H 4 .1 2 O 12 H 6 by the withdrawal of the elements of 3 molecules of water. The salts, such as KIO 4 , which on account of its isomorphism with potassium permanganate, K 2 Mn 2 O 8 , is more probably represented by the formula K 2 I 2 O 8 , are, it may be supposed, also derived from a distinct acid, formed in a similar manner by the withdrawal of the elements of a fourth molecule of water; these salts have the same empirical composition as the perchlorates, and as potassium perchlorate is isomorphous with the compound K 2 I 2 O 8 , it is probable that it corresponds with it in composition, and that perchloric acid therefore is represented by the formula H 2 C1 2 O 8 . If this conclusion be correct, and chloric acid be correctly represented by the formula HC1O 3 , we have an explanation of the great difference which is observed in the properties of these two acids. Perchloric acid, we have seen, has a great tendency to combine with water, and its hydrates may be regarded as distinct acids, the liquid hydrate bearing the same relation to the acid H 2 C1 2 O 8 that crystallized periodic acid bears to the hypothetical acid H 2 I 2 O 8 , from which salts such as K 2 I 2 O 8 are derived. Basic perchlorates corresponding to the basic periodates are not known, but a Cuprammonium salt and a lead salt have been obtained which apparently are derived from the crystalline hydrate of perchloric acid, the latter having the composition Pb"Cl 2 O 8 + PbH 2 O 2 or Pb"Cl 2 O 10 H 2 Pb".
It will have been noticed that the amount of heat developed on the addition of the first molecule of potassium hydroxide to the solution of periodic acid regarded as H 5 IO 6 is much less than is usually observed with other acids, and the amount developed on the addition of the second molecule much greater; the two molecules together, however, produce an effect comparable with that observed in the case of other acids. But the salt produced on passing chlorine into a solution of potassium hydroxide has the composition K 2 I 2 O 8 , so that the first action of potassium hydroxide on a solution of the acid H 4 .1 2 O 12 H 6 apparently does not merely consist in the direct displacement of hydrogen in this acid by potassium; the elements of four molecules of water are also withdrawn, and since this latter operation involves the absorption of heat, the amount of heat finally developed in the reaction is but small. On the addition of a further quantity of hydroxide, however, change in the reverse order is effected, and hence the normal amount of heat is developed by the combined action of the two molecules of hydroxide. It is interesting also to observe that more heat is developed on adding the third than by the fourth and fifth molecules of hydroxide together, and to contrast this with the circumstance that the silver salt Ag 4 I 2 O j2 H 6 loses the elements of two molecules of water at 100ð C., but the elements of a third only at 200’ C. Apparently there is a tendency on the part of potassium hydroxide to enter into reaction with the acid after the production of the normal salt K 4 I 2 O 12 H 6 ; but salts containing a relatively larger proportion of potassium evidently cannot exist except in solution, and even then only to a limited extent.
We are now in a position also to explain the formation of sodium periodate by the action of chlorine on a solution of sodium iodate and sodium hydroxide. We have learnt that the action of chlorine on the latter is to produce sodium hypochlorite, and that this is a powerful oxidizing agent—
NaOH + Cl 3 = NaOCl + HC1, and there can be little doubt, therefore, that the iodate is at first oxidized by it, and converted into the salt Na s IO 8 j thus—
Na 2 I 2 O 6 + 2NaOCl = Na 2 I 2 O 8 + 2NaCl. A A U A A O
In presence of sodium hydrate and water, however, this salt is at once converted into the normal periodate—
Na 2 I 2 O 8 + 2NaOH + 2H 2 O = Na 4 I 2 O 10 H 6 .
The following are the results of Thomsen’s thermochemical examination of iodic and periodic acids :—
[table]
It will be noticed that the amounts of heat developed in the formation of these two acids in aqueous solution from their elements are very simply related; thus—
(I 2 ,O 6 ,H 2 , Àq) = 111,420 heat-units or 3 × 37,140
(I,O li ,H δ , Aq) =184,400 „ „ 5 × 36,880
Again, both acids are easily decomposed by hydriodic acid, yielding water and iodine, and a similar simple relation obtains between the amounts of heat developed in the two reactions—
(I 2 O β H 2 Aq, 10HIAq) = 167,040 heat-units, or 5 × 33,408 (IO e H 5 Aq, 7HIAq) = 133,770 „ „ 4 × 33,442
The analogous decomposition of chloric acid by hydrochloric acid would involve the absorption of 30,920 units of heat, whilst that of bromic acid by hydrobromic acid would be attended with the development of 101,520 units of heat; the difference between these numbers, 132,440, is almost equal to the amount evolved in the decomposition of periodic acid, and on comparing the several reactions it will be evident that the differences between the amounts of heat evolved are approximately simple multiples of a common constant; thus— [table]
The differences are thus to each other as 4:2:1. The meaning of this remarkable relation is at present unknown.
Without further description, it will be obvious that whilst chlorine, bromine, and iodine and their compounds exhibit very considerable general resemblance, yet there are many most important differences between them; and also that chlorine is much more closely related to bromine than is the latter to iodine. [table]
These three elements are usually classed together, since most of their compounds with other elements are of the same type. In some respects they resemble each other closely, but in others they differ very widely,—selenium being much more closely related to sulphur on the one hand than to tellurium on the other.
Sulphur occurs native in most volcanic districts; it is also found abundantly in combination with metals, such as copper, iron, lead, and zinc, and it is widely distributed in the form of sulphates of barium, calcium, magnesium, and strontium; it is a necessary constituent of animal and vegetable tissues.
Selenium is of rare occurrence, and is always met with in combination with other elements; it is frequently associated in minute proportions with sulphur in its ores. Tellurium is also a rare element; it is occasionally found native, but generally in combination with various metals, such as bismuth, copper, lead, gold, or silver, and it is usually accompanied by selenium.
Sulphur is a very brittle solid of a lemon-yellow colour, insoluble in water, and therefore tasteless. It is a bad conductor of heat, and a non-conductor of electricity. Its specific gravity in the native crystalline state is 2O5. Sulphur melts at 115° C., and at 120° C. is converted into a perfectly limpid pale yellow liquid; but as it is gradually heated it becomes darker and more viscid until, at 200°-240° C., it is so thick that the vessel in which it is contained may be inverted without any outflow taking place. At 250°-300° C. it again liquefies, but does not become so fluid as when first melted, and at 440° C. it boils, yielding an orange-coloured vapour.
The density of sulphur vapour at 1000° C. is 32 times as great as that of hydrogen at the same temperature; so that as the atomic weight of sulphur is 32, the molecules of sulphur at this temperature are diatomic; but at about 500° C. its density is three times as great, so that at this temperature its molecules must be regarded as hexatomic.
Selenium is a deep brown-coloured brittle solid, of the specific gravity 4 , 5 in the crystalline state; like sulphur, it is insoluble in water. It boils at a temperature below a red heat, and gives off a deep yellow vapour, the density of which at about 1400° C. is such as to show that it consists of diatomic molecules, but at lower temperatures, like sulphur, its molecule is more complex.
Tellurium is much like a metal in appearance, resembling bismuth in colour; it is very brittle; its specific gravity is 6 ’ 1 to 6 ’ 33, and it conducts heat and electricity, though not very readily. It melts at about 500° C., and at a high temperature is converted into a yellow vapour, the density of which at 1440° C. is such as to show that the molecule of tellurium at this temperature is diatomic. No tendency has hitherto been observed on the part of tellurium to form a more complex molecule in the manner characteristic of selenium and sulphur; and it is noteworthy also that, whereas the density of sulphur vapour at temperatures not far removed from its boiling point is three times as great as at much higher temperatures, the density of selenium vapour under similar circumstances is only about one and a half times as great as at about 1400° C.
Sulphur is chiefly remarkable for the great number of allotropie forms in which it can exist. Thus, native sulphur is found crystallized in rhombic octahedra of the specific gravity 2 , 05. This variety melts at 114° *5 C.; it is readily soluble in carbon disulphide, the sulphide of chlorine, S 2 C1 2 , benzene, turpentine, <fcc., and sulphur is usually deposited in this form on spontaneous evaporation of its solution in carbon disulphide; it is, in fact, the stable form into which all other allotropes tend to change. When, however, melted sulphur is allowed to cool slowly it crystallizes in transparent, yellowish-brown, monoclinic prisms, of lower specific gravity, viz., 198, but of higher melting point (120° C.); in the course of a few days, although the crystals maintain their prismatic form, they become opaque, and, on examination, are found to be converted into aggregations of minute octahedrons. This change of prismatic into octahedral sulphur is attended with the development of a considerable amount of heat, which is especially noticeable when the conversion is caused to take place rapidly by scratching the crystals. The transformation of the octahedral variety into the prismatic takes place at about the melting point of the former, a transparent crystal of octahedral sulphur being converted into an opaque mass of prismatic crystals when heated for some time to a temperature of 105°-115° C.; also, when a saturated solution of sulphur in hot turpentine is allowed to cool, crystals of prismatic sulphur are at first deposited, but after a time, when the liquid has become comparatively cool, the crystals which separate are octahedral.
Prismatic sulphur is soluble in carbon disulphide and other menstrua which dissolve the octahedral modification; another modification, which is also soluble, but destitute of crystalline form, is obtained on addition of acids to solutions of alkaline polysulphides—
K 2 S 5 + 2HC1 = 4S + H 2 S + 2KC1.
Potassium Hydrogen < 5nl . nr Hydrogen Potassium pentasulphide. chloride. 0u1puur. su ]phide, chloride.
It has a greenish-white colour, and constitutes ordinary milk of sulphur; by keeping, it becomes gradually converted into octahedral sulphur. Sublimed sulphur, or flowers of sulphur, is probably closely allied to this form, but always contains a small proportion of insoluble sulphur.
When sulphur is heated to 260°-300° C. and then poured in a thin stream into cold water, it is converted into the so-called plastic modification, and is obtained as a soft yellowish-brown semitransparent mass, capable of being drawn out into fine elastic threads. In this state sulphur has the specific gravity 1 95, and is insoluble in carbon disulphide; in the course of a few hours, however, it again becomes brittle and almost entirely reconverted into the octahedral modification, the change being accompanied by the development of heat.
When the chloride of sulphur, S 2 CJ 2 , is decomposed by water, it furnishes sulphur, hydrochloric acid, and thiosulphuric acid; the sulphur thus obtained is amorphous and insoluble in carbon disulphide, but is converted into octahedral sulphur by fusion, or by exposure for some time to a temperature of 100° C. Thiosulphuric acid also gradually decomposes into sulphur and sulphurous acid, and if a solution of a thiosulphate is decomposed by hydrochloric acid—
Na 2 S 2 O 3 + 2HC1 = S + H 2 SO 3 + 2NaCl Sodium Hydrogen ς 1 h Sulphurous . un4 □ n thiosulphate. chloride. Sulphur. ac j<j. Sodium chloride.
the sulphur which separates is soluble in carbon disulphide, but is obtained on evaporation of the solution in an amorphous condition, and cannot then be again dissolved in carbon disulphide; it is stated, however, that if the hyposulphite is decomposed by dilute sulphuric acid, the sulphur which separates is insoluble in carbon disulphide. When a solution of ordinary sulphur in carbon disulphide is exposed to sunlight, an amorphous insoluble modification separates; the reconversion of this modification into octahedral sulphur is attended with the absorption of heat.
Several minor modifications of sulphur have been described, but it is not known whether the differences which they exhibit are inherent, or whether, as is not improbable, they are due to the presence of impurity.
There appears to be a relation between the modification formed in p reaction and the compound or compounds from which it is obtained, since, generally, the sulphur separated from compounds in which it is associated with positive elements is soluble, whilst that separated from compounds in which it is associated with negative elements is insoluble in carbon disulphide; but the conditions under which the sulphur is separated doubtless exercise an important influence.
At present we have no knowledge as to the manner in which the various allotropes are related to each other; it has been suggested, however, that the differences between them are at least in some cases due to differences in molecular composition.
Several allotropic modifications of selenium are also known, but a tendency to form allotropes has not been observed in the case of tellurium.
Three principal modifications of selenium, corresponding to the three principal modifications of sulphur, may be distinguished, viz., (1), black crystalline selenium, of specific gravity about 4 - 80, which may be placed by the side of the octahedral variety of sulphur, being the form in which selenium separates when solutions of metallic selenides are exposed to the air, and since all other modifications are converted into it when heated to about 150° C.; (2), red crystalline selenium, of specific gravity 4'46 to 4’51, which probably corresponds to monoclinic sulphur, as it has the same crystalline form, and may be obtained by rapidly cooling melted selenium; and (3), red amorphous selenium, of specific gravity about 4 ’ 3. This last corresponds to the amorphous allotrope of sulphur, and, as in the case of sulphur, it occurs in two modifications, the one soluble and the other insoluble in carbon disulphide. Black crystal line selenium, unlike octahedral sulphur, however, is insoluble in carbon disulphide, but the red crystalline allotrope is soluble, although to a much less extent than sulphur. The conversion of these modifications of selenium into each other is attended with development or absorption of heat, just as in the case of sulphur, and the melting points of the several modifications are different. Thus, if amorphous selenium, which melts a few degrees above 100° C., be heated to about 96° C., it quickly becomes crystalline, the change being attended with considerable rise of temperature, and the melting point rises to 217° C. The electrical conductivity of selenium is found to vary in a remarkable manner with the temperature, and is also influenced by light. Amorphous selenium is a non-conductor up to 80° C., but from this temperature up to 210* Ç. its conductivity gradually increases, after which it again diminishes. Selenium which has been kept for several hours at 210° C. and then gradually cooled is especially sensitive to the influence of light, its conductivity increasing with the intensity of the light.
Sulphur, selenium, and tellurium form gaseous compounds with hydrogen, analogous in composition to water :— [table]
Hydrogen sulphide may be produced directly from its elements by passing hydrogen gas into boiling sulphur, but it is always prepared by the action of a solution of hydrochloric or sulphuric acid on a metallic sulphide, that of iron being commonly employed—
FeS + 2HC1 = H 2 S + FeCl 2 . Ferrous sulphide. Hydrogen chloride. Hydrogen ∙ulphlde. Ferrous chloride. The compounds of selenium and tellurium with hydrogen are obtained in a similar manner from metallic selenides and tellurides. Hydrogen and selenium also directly enter into reaction; the quantity of hydrogen selenide formed is apparently a function of the temperature; it increases from 250° to 260° C., and decreases regularly from the latter point to 700° C.
Hydrogen sulphide, hydrosulphuric acid, or sulphuretted hydrogen gas, is transparent and colourless; it possesses a most offensive odour, and cannot be breathed with impunity, frequently giving rise to nausea and vertigo even when much diluted. It may be condensed by powerful pressure to an extremely mobile liquid, which solidifies when cooled in a bath of solid carbon dioxide and ether to a white transparent mass, which melts at — 85° C. Hydrogen sulphide is readily inflammable, burning in air with a blue flame, and forming water and sulphur dioxide—
2SH 2 + 3O 2 = 2SO 2 + 2H 2 O. Hydrogen sulphide. Oxygen. Sulphur dioxide. Water.
Most metals when heated in the gas decompose it, a metallic sulphide being produced, and hydrogen liberated. Hydrogen sulphide dissolves in water, a saturated solution containing about three times its volume of the gas; the solution has the odour and taste of the gas, and a slight acid reaction. It is gradually decomposed on exposure to the air, sulphur being deposited—
2SH 2 + O 2 = 2S + 2OH,.
It is decomposed in a similar manner by nearly all oxidizing agents, and by the action of chlorine, bromine, and iodine—
SH 2 + C1 2 = 2HC1 + S.
Iodine, however, cannot decompose the gas at ordinary temperatures except in presence of water. This is because the reaction requires an absorption of heat. Thus, in the production of a molecule of hydrogen sulphide from hydrogen and sulphur in the state in which it separates when hydrogen sulphide is decomposed by iodine, 4500 units of heat are developed; but in the production of a molecule of hydriodic acid from its elements, 6000 units of heat are absorbed, so that the reaction indicated by the equation H 2 S + I 2 = 2HI + S would involve the absorption of 4500 + 2 × 6000 = 16,500 units ofheat. That it takes place in presence of water is due to the fact that the dissolution of the hydriodic acid produced is attended with the development of 2 × 19200 = 38,400 units of heat; hence, when the reaction is effected in a dilute solution, heat is developed to the extent of 38,400-16,500 = 21,900 units of heat. But it is found that the reaction takes place the less readily as the concentration of the solution increases, and that it ceases when the solution has attained a specific gravity of 1 - 56 at ordinary temperatures; in more concentrated solutions sulphur even dissolves with production of hydrogen sulphide and liberation of iodine. A simple explanation of this apparently anomalous result, however, is afforded by the observation that the heat developed by the absorption of equal quantities of hydriodic acid is less as the quantity of acid already dissolved in the water is greater. The amount of heat developed, therefore, diminishing as tke quantity of hydriodic formed by the reaction in the liquid increases, at a certain point becomes equal to that absorbed in the decomposition of the hydrogen sulphide by the iodine, and the reaction ceases since it can no longer be attended with a development of heat.
This behaviour of iodine with hydrogen sulphide alone, or in presence of water, is one of the most striking illustrations of the fact that reactions involving the expenditure of energy cannot take place directly, and are only possible when the conditions are such that one or more of the products of the reactions enter into secondary reactions, so as to cause the development of more heat than is absorbed in the primary reaction.
The hydrogen in hydrogen sulphide may be displaced by metals,—the compounds formed by displacing one-half the hydrogen being termed sulphydrates or hydrosulphides, whilst those in which the whole of the hydrogen is displaced are termed sulphides. These two classes of compounds correspond to the metallic hydroxides and the metallic oxides respectively, and in many respects closely resemble them; the sulphur compounds, however, are, with few exceptions, far less stable than the corresponding oxygen compounds.
Hydrogen sulphide enters directly into reaction with the metallic hydroxides, exchanging its hydrogen for the metal; for example—
NaOH + H 2 S = NaSH + HOH.
Sodium Hydrogen Sodium hydroxide. sulphide. hydrosulphide. nater.
It therefore exhibits the behaviour of an acid. From Thomsen’s experiments it appears that 7740 units of heat are developed on the addition of a solution of one molecule of sodium hydroxide to a solution of one molecule of hydrogen sulphide, and that the further addition of the hydroxide is without effect. Hydrogen sulphide is thus proved to be a monobasic acid, and this result also shows that when soluble sulphides, such as sodium sulphide, are dissolved in water, double decomposition occurs, thus: Na 2 S + OH 2 = NaSH + NaOH, just as when sodium oxide, for example, is added to water: Na 2 O + OH 2 = 2NaOH. It is uncertain, however, whether the decomposition of the sulphides by water in this manner is complete, or whether it is only partial, and the more complete the greater the quantity of water present.
The highly positive metals lithium, sodium, potassium, calcium, strontium, barium, and magnesium form soluble sulphides and hydrosulphides, but most of the sulphides of other metals are insoluble. The nature of many of the compounds precipitated from metallic solutions by hydrogen sulphide or an alkaline hydrosulphide is not well established; but in many cases apparently they are intermediate in composition between the hydroxides and hydrosulphides; the precipitate formed on the addition of an alkaline hydrosulphide to a solution of a zinc salt, for example, is probably a compound of this kind, and may be represented by the formula HO.Zn.SH,—zinc hydroxide being HO.Zn.OH, and zinc hydrosulphide HS.Zn.SH. The solutions of salts of heavy metals, such as mercury and lead, furnish precipitates of the corresponding sulphides with hydrogen sulphide or alkaline hydrosulphides.
The hydrosulphides of certain elements, such as aluminium and chromium, cannot exist in presence of water, but enter into reaction with it with evolution of hydrogen sulphide; hence, on the addition of an alkaline hydrosulphide to solution of their salts, the corresponding hydroxides are precipitated :— [table]
Sulphur unites with all the metals and with most of the non-metallic elements; the sulphides are therefore usually prepared directly from their elements. Two classes of sulphides corresponding to the basic and acid oxides may be distinguished, but the distinction between them is much less marked than that between the two classes of oxides. The sulphides of the non-metallic elements and the sulphides of arsenic, antimony, tin, molybdenum, tungsten, vanadium, gold, and platinum, which are soluble in solutions of alkaline hydrosulphides, belong to the class of acid sulphides, and the remaining sulphides are basic. These two classes of sulphides are capable of uniting together to form sulphur salts, just as the basic and acid oxides combine formiûg oxy-salts. As a rule, the sulphides and oxides of the same element have similar formulae and correspond in their general behaviour. Occasionally there are oxides to which there are no corresponding sulphides, but more fre quently sulphides to which there are no corresponding oxides. All sulphides are decomposed more or less readily by hydrochloric acid gas, hydrogen sulphide and a chloride being produced, and in some cases sulphur is also liberated. Many of the sulphides which are decomposed when heated in an atmosphere of hydrochloric acid gas are little, if at all, affected by a boiling solution of the acid.
Hydrogen selenide and telluride closely resemble hydrogen sulphide in properties, but are far less stable compounds. The former has a most offensive acrid odour, impairing or even destroying the sense of smell for several hours, and producing inflammation of the eyes. Their aqueous solutions are decomposed on exposure to the air with separation of selenium and tellurium respectively; and with solutions of salts of most metals they produce precipitates of tire corresponding selenhydrates or selenides and tellurhydrates or tellurides. The selenides and the tellurides, like the sulphides, may be formed by the direct combination of their elements, and are usually so prepared; in their général behaviour they resemble the corresponding sulphides.
A higher sulphide of hydrogen is known of which the composition has not yet been satisfactorily determined, owing to the difficulty of obtaining it in a pure state; but as it is obtained on adding a solution of potassium pentasulphide, K 2 S 5 , to a dilute solution of hydrochloric acid, and is formed without evolution of hydrogen sulphide or separation of sulphur, it appears most probable that it is the pentasulphide H 2 S 5 . It separates as an oily liquid, heavier than water, possessing a peculiar sulphurous disagreeable odour; it is soluble in water. It is a very unstable substance, undergoing decomposition into hydrogen sulphide and sulphur with great facility; this decomposition is instantaneous under the influence of substances such as finely divided platinum, gold, and charcoal. Its stability is increased by the presence of moderately strong acids, but diminished by alkalies. With an alcoholic solution of potassium hydrosulphide it furnishes potassium pentasulphide, hydrogen sulphide being evolved—
H 2 S 5 + 2KHS = K 2 S 5 + 2H 2 S.
This reaction is analogous to that between sulphuric acid and potassium hydroxide :—
H 2 SO 4 + 2KHO = K 2 SO 4 + 2H 2 O; hydrogen pentasulphide may, in fact, be regarded as the analogue of sulphuric acid.
Sulphur, selenium, and tellurium form compounds with chlorine, bromine, and iodine, similar in composition, but differing greatly in stability.
With chlorine, sulphur forms the three compounds— S 2 Cl 2 , SCl 2 , and SC1 4 , all of which are liquid. The chloride S 2 Cl 2 is obtained by passing chlorine over sulphur, which is gently heated; it is a mobile reddish-yellow liquid, having a peculiar, penetrating, and most disagreeable odour. It boils without suffering decomposition at 137° C. It is slowly decomposed by water, yielding hydrogen chloride, sulphur, and thiosulphuric acid —
2S 2 C1 2 + 3OH 2 = 4HC1 + 2S + H 2 S 2 O 3 .
When saturated with chlorine at about 10° C. it is converted into the dichloride SC1 2 , but if saturated with chlorine at about - 22° C. it furnishes the tetrachloride SC1 4 . Both of these compounds are so unstable, however, that they are resolved into the lowest chloride S 2 Cl 2 and chlorine when very slightly heated.
The chlorides of selenium and tellurium are also obtained by the direct action of chlorine on the elements. Two chlorides of selenium are known, Se 2 Cl 2 and SeCl 4 , the former being liquid and the latter a white crystalline solid.
Tellurium forms the two chlorides TeCl 2 and TeCl 4 , both of which are solid. These chlorides of tellurium, as well as selenium tetrachloride, appear to volatilize without decomposition.
Very little is known of the bromides of sulphur and selenium, but from the observations which have been made there is no doubt that the sulphur bromides are much less stable than the chlorides; sulphur, therefore, like oxygen, appears to have less affinity for bromine than for chlorine. Tellurium furnishes two crystalline bromides, TeBr 2 and TeBr 4 , both of which may be sublimed without decomposition.
The only iodide of sulphur which is known has the composition S 2 I 2 , and is obtained by the direct union of its elements; it is a black crystalline solid, insoluble in water, and readily decomposes when heated. Iodides of selenium are not known with certainty, but two tellurium iodides have been prepared, Tel 2 and Tel 4 ; both are black crystalline bodies, which give off iodine when heated.
The chlorides and bromides of selenium and tellurium, like the chlorides and bromides of sulphur, are decomposed by water, but the iodides of these three elements are comparatively stable and are scarcely affected unless heated with water.
Sulphur, selenium, and tellurium burn in oxygen or air, forming the dioxides
SO 2 , SeO 2 , TeO 2 .
Under certain conditions, sulphur dioxide takes up an additional atom of oxygen, and is converted into the trioxide, SO 3 ; a tellurium trioxide, TeO 3 , may also be obtained, but the corresponding oxide of selenium is not known. A third oxide of sulphur, S 2 O 3 , has recently been described.
The dioxides and trioxides of sulphur, selenium, and tellurium have the property in common of forming corresponding acids when combined with water :—
H 2 SO 3 H 2 SeO 3 H 2 TeO 3 ∙
Sulphurous acid. Seleuious acid. Tellurous acid.
H 2 SO 4 H 2 SeO 4 H 2 TeO 4 .
Sulphuric acid. Selenic acid. Telluric acid.
These compounds differ greatly in stability and in their properties generally—the sulphur and selenium compounds being closely related, whilst the tellurium compounds are widely different from them in most respects. .
Sulphtfr dioxide or sulphurous anhydride is a colourless gas, of pungent suffocating odour; by a pressure of three atmospheres, or a refrigerating mixture of ice and salt, it is readily condensed to a colourless mobile liquid, which boils at about - 10° C. The liquid anhydride freezes at - 76° C., forming a transparent colourless crystalline solid, which melts at about -79° C. Water at 0° C. dissolves about 68 times its bulk of the gas, but only 32 times its bulk at 24°, forming a solution of sulphurous acid, H 2 SO 3 , which is readily decomposed by heat.
Selenium dioxide or selenious anhydride is a white infusible substance, which volatilizes at a temperature below redness; its vapour condenses in snow white deliquescent prisms, which dissolve in water forming selenious acid, H 2 SeO 3 .
Tellurium dioxide, however, is only very slightly soluble in water, and the solution does not exhibit an acid reaction; the attraction of this oxide for water, in fact, is so slight, that when a solution of tellurous hydrate (tellurous acid), H 2 TeO 3 , is heated to about 40° C. the oxide separates. It fuses readily and volatilizes, and the fused oxide, which is a transparent, deep yellow liquid, solidifies on cooling to a white crystalline mass.
The difference in physical properties between sulphur and seleqium dioxides is so great as to suggest that the molecule oí the latter is not represented by the formula SeO 2 , but that it is more complex; and although the vapour density of selenium dioxide is such as to negative this assumption for the gaseous substance, there is no evidence to show that the expression SeO 2 is the correct molecular formula for the solid oxide.
Sulphur trioxide or sulphuric anhydride is a white, highly volatile substance which dissolves in water, forming sulphuric acid, H 2 SO 4 , the combination being attended with the development of so much heat that when the anhydride comes in contact with the water a violent hissing is produced, just as when a red-hot iron is plunged into water. Selenium trioxide has not been isolated, but selenic acid, H 2 SeO 4 , the acid corresponding to sulphuric acid, is known, and very closely resembles the latter compound in many of its properties. Tellurium trioxide, TeO 3 , is an orange-yellow solid, insoluble in water, and therefore altogether different in properties from sulphur trioxide; moreover, the resemblance between the compound H 2 TeO 4 , from which tellurium trioxide is obtained by heating to a temperature below redness, and sulphuric and selenic acids scarcely extends beyond the similarity of their formulae, since telluric acid is nearly insoluble in water, and its solution has a metallic rather than an acid taste, and reddens litmus but slightly.
Sulphur dioxide is usually prepared for laboratory purposes by heating a metal, such as copper or mercury, with concentrated sulphuric acid; the most probable explanation of its formation in this manner appears to be that the metal acting upon the acid produces the corresponding sulphate and hydrogen, and that the latter whilst in the nascent state acts upon another portion of the acid, reducing it to sulphurous acid, which splits up into sulphur dioxide and water; thus— [table]
We may also suppose, however, that the metal merely withdraws an atom of oxygen from the sulphuric acid, and that the oxide produced dissolves in the acid, forming a sulphate— [table]
The kind of action represented by the first of these equations undoubtedly takes place when carbon in the form of charcoal is heated with sulphuric acid, whereby carbon dioxide, sulphur dioxide, and water are produced—
C + 2H 2 SO 4 ≈ CO 2 + 2SO 2 + 2H 2 O. carbon. Sulphuric acid, carbon dioxide. Sulphur dioxide. Water.
When copper is employed not only are sulphur dioxide, water, and cupric sulphate produced, but cuprous and cupric sulphides, Cu 2 S and CuS, it is stated, are also formed; their production, however, is due to secondary action.
Selenium dioxide is obtained either by burning selenium in a stream of oxygen, or by evaporating a solution of selenious acid to dryness.
Sulphurous acid is always prepared by passing sulphur dioxide gas into cold water; the combination is attended with a slight elevation of temperature. Selenious acid is produced in a similar manner by dissolving the oxide in water, or by dissolving selenium in concentrated nitric acid, and evaporating to expel the excess of nitric acid. Teliurous acid may be obtained by dissolving tellurium in nitric acid of specific gravity 1 ’ 25, and pouring the colution, after the lapse of not more than a few minutes, into water. If the precipitation be delayed for a longer time, the oxide TeO 2 is thrown down instead of the hydrate. It is best prepared by decomposing tellurium tetrachloride with water—
TeCl 4 + 3H 2 O = H 2 TeO 3 + 4HCI. Tellurium tetrachloride. Water. Tellurous hydrate. Hydrogen chloride.
It is stated that anhydrous sulphurous acid, H 2 SO 3 , has been obtained in crystals by cooling a saturated aqueous solution to 0° C.; a crystalline hydrate of the composition H 2 SO 3 + 8H 2 O was obtained by Pierre by cooling to - 6° C. a saturated solution through which a current of the gas was being transmitted. The solution of sulphur dioxide in water is strongly acid, and effervesces with carbonates; by passing a current of the gas through water in which metallic hydroxides or carbonates are dissolved or suspended metallic sulphites are produced. Two classes of metallic sulphites may be thus formed :— the acid sulphites or bisulphites in which one-half the hydrogen in sulphurous acid is displaced by a metal; and normal sulphites, in which the whole of the hydrogen is displaced. It is also possible to displace the two atoms of hydrogen in sulphurous acid by two different metals, and thus to obtain so-called double salts. The following table shows the composition of some of the sulphites, disregarding the water of crystallization which several of them contain :— [table]
The acid sulphites of barium, strontium, calcium, and magnesium, and the acid and normal sulphites of the alkali metals, are soluble in water, although only the sulphites of alkali metals are freely soluble; but most other sulphites are insoluble, and may be prepared by precipitation with an alkaline sulphite. The acid sulphites are prepared by saturating a solution of the metallic hydroxide or carbonate with sulphur dioxide gas, and then adding to it as much of the hydroxide or carbonate as it originally contained; thus—
Na 2 CO 3 + OH 2 + 2SO 2 = 2NaHSO 3 + CO 2 Sodium carbonate. Water. Sulphur dioxide, s ° di ^ l lp h ^° geπ <‰ a ‰ 0 f
2NaHSO 3 + Na 2 CO 3 = 2Na 2 SO 3 + H 2 O + CO 2 . “ϊΧ" c =, S<->"∙~W∙ W.« Ä The alkali and alkaline-earth metals are the only ones, however, which readily produce acid sulphites; in fact, the tendency to form acid salts is almost restricted to these metals in the case of all acids, and it is difficult to obtain acid salts containing heavy metals.
A solution of sulphurous acid slowly absorbs oxygen from the air and is converted into sulphuric acid; in like manner, the sulphites, particularly if in solution, become converted into sulphates on exposure to the air. All sulphites are decomposed at a red heat, either into sulphate and sulphide, or into sulphur dioxide and metallic oxide. They are also decomposed by all acids excepting carbonic and boric acids, sulphurous acid being liberated. Sulphurous acid possesses considerable bleaching powers, and is extensively employed in bleaching straw, wool, and many other articles, which would be injured by chlorine. The articles to be bleached are moistened and suspended in closed chambers in which sulphur is burnt; the sulphur dioxide produced is then absorbed by the damp goods, and their colour is discharged. The manner in which it acts is not well understood, but it appears to be by forming colourless compounds with the colouring matters; it does not, like chlorine, decompose colouring matters, for the colour may usually be restored by acids or alkalies. The reproduction of the yellow colour of new flannel, when it is washed with an alkaline soap for the first time, is an illustration of this. Sulphurous acid is also a powerful antiseptic, and is highly valuable as a disinfecting agent.
Selenious acid is deposited from a hot aqueous solution on slow cooling in colourless prismatic crystals like saltpetre, It closely resembles sulphurous acid in properties, and like it furnishes acid selenites, such as potassium hydrogen selenite, KHSeO 3 , and normal selenites, such as potassium selenite, K 2 SeO 3 , but it also forms so-called hyperacid salts with the alkali metals, e.g., HKSeO 3 + H 2 SeO 3 , The selenites of alkali metals are soluble in water, but the other selenites are insoluble.
Tellu∙ous hydrate, prepared by decomposing tellurium tetrachloride with water, is a bulky precipitate, which, when dried over sulphuric acid, forms a light white earthy powder, having a bitter metallic taste. It is only slightly soluble in water, but dissolves in alkalies and alkaline carbonates. With the alkali metals tellurous acid forms three classes of salts corresponding to the three classes of selenites, of which the following are examples— [table]
With the alkaline earth metals tellurous acid forms normal salts, and also certain peculiar salts, such as BaTe 2 O 5 or BaTeO 3 + TeO 9 , and BaTe 4 O 9 or BaTeO<, + 3TeO 2 . With the heavy metals it appears to form only normal salts. The tellurites of alkali metals are soluble in water, but those of other metals are difficultly soluble or insoluble.
In its oxides, therefore, it will be evident tellurium differs widely from selenium and sulphur, but it much resembles metals like bismuth and antimony, since it forms salts with acids. Thus, a tellurium sulphate of the composition Te(SO 4 ) 2 , it is stated, is formed by dissolving tellurium in concentrated sulphuric acid. Tellurous hydrate also dissolves readily in acids; the solution in hydrochloric acid probably contains either tellurium chloride, TeCl 4 , or a compound intermediate ia composition between tellurium chloride and tellurium hydroxide, Te(OH) 4 , formed from the latter by the partial displacement of the OH groups by chlorine; when it is dissolved in oxy-acids, apparently salts are formed in which the group TeO displaces hydrogen, corresponding to the antimony salts in which the radicle SbO displaces hydrogen.
Sulphurous acid, on account of its tendency to form sulphuric acid, has considerable power as a reducing agent. Thus iodine and sulphurous acid, in presence of a large quantity of water, yield hydriodic and sulphuric acid—
I 2 + H 2 O + H 2 SO 3 = 2HI + H 2 SO 4 . Sulphurous acid. Sulphuric acid.
With chlorine and bromine similar reactions occur still more readily. On this account, sulphurous acid is largely •employed as an “antichlor” to remove the excess of chlorine from articles bleached with bleaching powder. When solutions of sulphurous and selenious acid are mixed, the latter is reduced, especially on warming, and red amorphous selenium precipitated; similarly, a black precipitate of tellurium is formed on warming a solution containing tellurous and sulphurous acids. But these acids also readily part with at least a portion of their oxygen. For instance, when hydrogen sulphide gas is passed into a solution of sulphurous acid, sulphur is deposited, and the solution contains pentathionic acid—
5H 2 S + 5H 2 SO 3 =< 5S + H 2 S 5 O 6 + 9OH 2 . Hydrogen sulphide. 8u ¾ r ° ug Sulphur pe " t c a ^ ionic Water. When a solution of selenious acid is similarly treated, a precipitate is thrown down, which apparently consists chiefly of selenium sulphide, SeS 2 , mixed with a lower sulphide and free sulphur. Solutions of tellurous hydrate when thus treated furnish a brown precipitate of tellurium sulphide.
Sulp>hur Trioxide, SO 3 .
Dry sulphur dioxide gas and oxygen readily combine when the mixture is passed over gently heated platinum black or platinized pumice, forming sulphur trioxide or sulphuric anhydride, SO 3 . This compound may be obtained from sulphuric acid by distilling it with phosphoric anhydride :—
H 2 SO 4 + P 2 O 5 = SO 3 + 2HPO 3 .
Sulphuric ael<L 5** Sulphur trloxldu
It is usually prepared from N^ordhausen sulphuric acid,, which gives off sulphur trioxide when gently heated, ordinary sulphuric acid remaining behind.
Sulphur trioxide usually crystallizes in white slender needles, but it exists in two modifications. Thus, when the liquid oxide is cooled, it solidifies at 16° C. in long colourless prisms, which melt at the same temperature; but if kept at temperatures below 25°, it changes into a mass of fine white needles. This second modification gradually liquefies at temperatures above 50°, and again passes into the first modification; it dissolves with extreme slowness in sulphuric acid, whereas the liquid oxide is miscible in all proportions with the acid. Liquid sulphuric anhydride undergoes very great expansion by heat, its mean coefficient of expansion between 25° and 45° C. being , 0027 for l o C.; it is quite colourless when pure. Sulphur trioxide very readily parts with one of its atoms of oxygen, converting phosphorus trichloride, for example, into phosphorus oxy trichloride—
PC1 3 + SO 3 = SO 2 + POC1 3 . Phosphorus Sulphur Sulphur Phosphorus trichloride. trioxide. dioxide. oxytrichloride.
This reaction takes place when the substances are merely mixed together, although a strong red heat is necessary in order to resolve the trioxide into sulphur dioxide and oxygen.
When finely divided sulphur is added in small quantities to liquid sulphuric anhydride, drops of a deep blue colour sink to the bottom and solidify immediately. These consist of the compound S 2 O 3 . The temperature during the operation must be kept at 15° C., for if it is lower the anhydride does not remain liquid, and if higher the substance decomposes. When about a gramme of sulphur has been added the anhydride is poured off, and the solid residue freed from any that remains by a gentle heat. A bluish-green crystalline mass is thus obtained, which decomposes at ordinary temperatures, giving off sulphur dioxide, and leaving sulphur; water immediately decomposes it, forming sulphuric, sulphurous, and probably thiosulphuric acids. The corresponding selenium compound, SeSO 3 , may be prepared in a similar manner; Sulphuric Acid, H 2 SO 4 .
Of all chemical compounds this is probably the most important, on account of its numerous practical applications, and enormous quantities of it are now manufactured. The method employed consists essentially in oxidizing sulphurous acid by atmospheric oxygen, which is accomplished with the aid of nitric oxide gas in the following manner. Sulphur dioxide gas is prepared by burning sulphur, or iron pyrites, FeS 2 , in a properly constructed furnace, sulphur being always employed when the purest qualities of acid are required, and this gas is conveyed into a large leaden chamber, into which steam and air are also continuously admitted. At the back of the furnace in which the sulphur is burnt, a small vessel is placed containing sodium nitrate and sulphuric acid, which continue to generate nitric acid for some time; thus—
NaNO 3 + H 2 SO 4 = HNO 3 + NaHSO 4 .
Sodium nitrate. Sulphuric acid. Nitric acid. Sodium hydrogen sulphate.
The exact changes which occur among the various substances thus introduced into the chambers are but imperfectly understood; the nitric acid, however, is apparently decomposed by the sulphurous acid in such a manner that finally nitric oxide is produced—
2HNO 3 + 3SO 2 + 2H 2 O = 2NO + 3H 2 SO 4 .
But nitric oxide combines directly with oxygen at ordinary temperatures, forming the oxides N 2 O 3 and N 2 O 4 , and it is generally supposed that the nitric oxide thus formed acts as a carrier of oxygen,—that on coming in contact with the oxygen of the air introduced into the chamber, it is converted into a higher oxide, which, however, is no sooner produced than it is reduced again to nitric oxide by the sulphurous acid, which is oxidized to sulphuric acid; supposing the oxide formed to be nitric peroxide, we may represent the change by the following simple equation :—
2SO 2 + 2H 2 O + N 2 O 4 = 2H 2 SO 4 + 2NO.
The nitric oxide again absorbs oxygen to form the higher oxide, which again enters into reaction with sulphur dioxide and water, with liberation of nitric oxide, and so on continuously. Theoretically, therefore, a limited quantity of nitric oxide should suffice to produce an indefinitely large amount of sulphuric acid, but practically this is not the case, chiefly owing to the loss occasioned by the dissolution of a certain quantity in the sulphuric acid which collects on the floor of the chambers; a certain amount also escapes from the chambers with the nitrogen of the air, but this is usually absorbed by passing the escaping vapours through a tower filled with pieces of coke, kept moistened with sulphuric acid, or over which water is constantly pouring; when acid is employed it is afterwards pumped into a similar scrubber, and deprived of its nitric oxide gas by means of the sulphur dioxide gas from the burning sulphur or pyrites, and when water is used it is pumped into, and distributed on the floors of, the leaden chambers.
In practice the amount of sulphuric acid produced usually approximates very closely to the theoretical quantity.
It is beyond question that the reactions which occur in the formation of sulphuric acid are not all of the simple character represented by the equations above given. Thus, the formation of nitric oxide from nitric acid is probably the final result of three distinct reactions. In the first place, from the sulphur dioxide and nitric acid a compound which is known as the lead chamber crystals, on account of its forming crystals, is probably produced; thus—
I. SO 2 + NO 2 .OH = SO 2 {g⅛.
Sulphur dioxide. Nitric acid. Lead chamber crystals.
This compound, it will be observed, is very closely related to sulphuric acid, and may be regarded as formed from it by the displacement of one of the OH groups by the group NO 2 . It is decomposed by water, yielding sulphuric and nitrous acids :—
II SO 2 {g°≈ + HOH = SO, {°≡ + NO 2 H. Lead chamber crystals. Water. Sulphuric acid. Nitrous acid. From the nitrous acid thus formed nitric oxide may be produced in two ways,—either by its spontaneous decomposition in accordance with the equation
3HNO 2 = 2NO + HNO 3 + H 2 Oj Nitrous acid. Nitric oxide. Nitric acid. Water.
or by the action of sulphur dioxide—
2HNO 2 + SO 2 = 2NO + H 2 SO 4 . Nitrous acid. Sulphur dioxide. Nitric oxide. Sulphuric acid.
The latter reaction appears to be the more probable on account of the large amount of sulphur dioxide present, but it is not unlikely that both occur.
With regard to the part which the nitric oxide plays in the conversion of sulphurous into sulphuric acid, it is highly probable that, by the conjoint action of the oxygen and water present, it is converted into nitrous acid, from which nitric oxide is reproduced by the action of the sulphur dioxide in the manner above pointed out—
4NO + O 2 + 2OH 2 = 4HNO 2 .
If this be the case, nitrous acid is the active agent in the conversion of sulphur dioxide into sulphuric acid. It is known to effect the oxidation of sulphur dioxide with the greatest readiness, whereas dilute nitric acid has very little action. On this account, it is probable also that the nitric oxide is not converted into the peroxide, as sometimes is supposed, and that the latter is the active agent in the formation of sulphuric acid, for our general experience appears to warrant the conclusion that, as sulphur dioxide enters so readily into reaction with nitrous acid, this acid would Be produced from the nitric oxide almost entirely to the exclusion of other oxidation products. Moreover, nitric peroxide in presence of water forms nitrous and nitric acids: N 2 O 4 + H 2 O = HNO 2 + HNO 3 j so that, even supposing that the nitric oxide is converted into nitric peroxide, the formation of sulphuric acid would still in a large measure be due to the action of nitrous acid. Although sulphur dioxide acts very slowly on dilute nitric acid, it readily reduces it to nitric oxide in presence of moderately concentrated sulphuric acid; hence any nitric acid formed in the process, when carried down to the floor of the chamber and mixed with the moderately concentrated sulphuric acid there collected, would also undergo reduction by the sulphur dioxide.
The acid from the leaden chambers has generally a specific gravity of about 1 ‘55 ; by concentration in shallow leaden pans this is raised to above 1 ’ 7. The further concentration is effected in glass or platinum retorts. The commercial acid, known as oil of vitriol, has a specific gravity of about 1 ‘84; it is frequently of a more or less brown colour from the presence of organic matter, and always contains lead. Other impurities, such as arsenic and nitrous or nitric acid, are not unfrequently present.
Pure sulphuric acid, H 2 SO 4 , is an oily, colourless, inodorous liquid, of specific gravity 1 , 842; it solidifies at - 35° C. It cannot be distilled unchanged, and when heated gives off the anhydride until a liquid remains containing about 1 - 5 per cent, of water and 98 5 per cent, of sulphuric acid, which boils at 338° C.; acid of this strength may always be obtained by boiling down acid of any other degree of concentration. When the acid is vaporized, dissociation takes place very rapidly, the vapour at temperatures above 400° C. consisting entirely of a mixture of water and sulphuric anhydride. Sulphuric acid chars most organic substances containing oxygen and hydrogen on account of its great affinity for water, causing the separation of these elements in the form of water and the liberation of carbon; when exposed to moist air it will even double its weight in the course of a few days, and on this account it is much employed as a desiccating agent. Much heat is developed when it is mixed with water, as will be evident from the following table, which represents the amounts of heat developed on mixing one molecule of the acid (98 grammes) with n molecules of water at about 18° C., according to Thomsen’s determinations :—
[table]
The last number in the table represents the amount of heat which, according to calculation, would be developed on mixing a molecule of sulphuric acid with an infinitely large quantity of water. It will be seen that the addition of the first two molecules is accompanied with the development of an amount of heat equal to half the entire amount developed on adding an excess of water; this appears to be accounted for by the formation of definite compounds of sulphuric acid with water. The first of these hydrates, H 2 SO 4 + H 2 O, may be obtained crystallized in colourless six-sided prisms, by cooling a mixture of the acid and water in these proportions to about 8° C.; it boils at 205°-210° C., and may be produced by evaporating any more dilute sulphuric acid at 205° C. until it ceases to lose water. The second hydrate, H 2 SO 4 + 2H 2 O, may, in a similar manner, be obtained by evaporating any more dilute acid at 100° C. until it ceases to lose water; it boils at 193° C. The specific gravity of the first hydrate is 1 - 78, and that of the second 1∙62; when water is added to the acid in the proportions to form the latter, the maximum condensation (about 8 per cent.) is observed which occurs on mixing these substances.
Further evidence in support of the view that sulphuric acid forms definite compounds with water which may be regarded as distinct acids is afforded by its behaviour on neutralization, by the persistent manner in which many sulphates retain one or two molecules of water, and by the existence of so-called basic sulphates. Thus, when it is neutralized by sodium hydroxide, considerably less heat is developed by the first than by the second molecule of hydroxide; this behaviour does not appear remarkable if it be supposed that it exists in solution in combination with water, and that, therefore, as in the case of periodic acid, not only is hydrogen displaced by sodium on the addition of the hydroxide, but the elements of one or more molecules of water are also separated, the latter being an operation which necessarily would involve the expenditure of energy, as the combination of the acid with water is attended with the development of heat (see page 488).
By displacing one-half the hydrogen in sulphuric acid by metals acid sulphates are formed, and normal sulphates are produced by displacing the whole of the hydrogen. Many of the normal sulphates crystallize with the same number of molecules of water, and are isomorphous, as, for example— [table]
Five of the seven molecules of water are removed with facility from these salts, but the sixth is less readily removed, especially from magnesium sulphate, and they are not deprived of the seventh unless heated to 200°-300° C. Copper, cadmium, and manganese sulphates behave similarly, and are only with difficulty deprived of the last molecule of water. The last molecule retained with such persistency by these salts was termed, by Graham, water of constitution, to distinguish it from water of crystallization. In the present state of our knowledge it is impossible absolutely to define the meaning of these terms, or even to say that there is an absolute difference between the so-called water of constitution and water of crystallization, and not merely one of degree; but since we have independent evidence tending to prove the existence of distinct acids formed by the combination of sulphuric acid with water, we may conveniently regard the above-mentioned salts as derived from these acids. Magnesium sulphate, for example, we may regard as the normal magnesium salt of the acid H 6 SO 6 , crystallized with five molecules of water. It will be evident that such salts as this bear a relation to the acid from which they are regarded as derived similar to that which the normal periodates bear to crystalline periodic acid, H 5 IO 6 j indeed, a very considerable general resemblance may be traced between the periodates and the sulphates. Thus, the sulphates of potassium, K 2 SO 4 , and of silver, Ag 2 SO 4 , separate from aqueous solution in anhydrous crystals; and sodium sulphate, Na 2 SO 4 , is also deposited in the anhydrous state at temperatures above 34° C.; it is only with metals such as magnesium, zinc, and copper, that salts are obtained which may be regarded as derived from the acids H 4 SO 5 or H 6 SO 6 . Similarly, the periodates KIO 4 , NaIO 4 , and AgIO 4 , derived from the at present hypothetical acid HIO 4 , are readily produced, but with other metals chiefly salts derived from the acid H 5 IO 6 are obtained.
But that these hydrates of sulphuric acid are distinct acids derives its most important confirmation from the existence of the so-called basic sulphates, such as Zn 2 SO 5 or ZnSO 4 + ZnO, Hg 3 SO 6 or HgSO 4 + 2HgO, and Cu 3 SO β + 3H 2 O or CuSO 4 + 2CuO + 3H 2 O; these salts are strictly comparable with the basic periodates.
A large number of double and mixed salts formed by the union of two or more sulphates are known. Thus, magnesium sulphate and the isomorphous sulphates form isomorphous double salts with potassium, sodium and ammonium sulphates, which crystallize with six molecules of water, and which usually are therefore regarded as formed by the displacement of the molecule of so-called water of constitution by the alkaline sulphate— [table]
It has-already been pointed out that sulphuric acid may be regarded as a compound of the dyad radicle SO 2 with the monad radicle OH, or as (SO 2 )"(OH) 2 ', this view of its “constitution” being chiefly founded on its behaviour with the chlorides of phosphorus. Thus, by the action of phosphorus pentachloride on sulphuric acid, an atom of oxygen and an atom of hydrogen, in other words, the group OH, are displaced by a single atom of chlorine—
HO.SO 2 . OH + PC1 5 = HO. S0 2 . Cl + POC1 3 + HC1. Although evidence of the existence of a second OH group in sulphuric acid cannot be obtained in a similar manner by the continued action of the phosphorus pentachloride, it is afforded by the behaviour of the compound SO 2 C1 2 , produced by the direct union of chlorine and sulphur dioxide with water, which successively converts it into the compound HO. SO 9 . Cl and sulphuric acid— [table]
Sulphuric chloride, SO 2 C1 2 , and sulphuric chlorhydrate, SO 2 (OH)Cl, are both colourless liquids; the former boils at about 70° C., and the latter at 158° C. Sulphuric chlorhydrate may be formed by the direct combination of hydrogen chloride with sulphuric anhydride, and when heated in closed vessels to 170° C., it furnishes sulphuric chloride and sulphuric acid :—
2C1.SO 2 .OH = C1.SO 2 .C1 + H0.S0 2 .0H.
Sulphuric acid enters into reaction with a large number of hydrogenized carbon compounds in such a manner that water is formed and hydrogen displaced by the monad group SO 2 . OH; for example—
HO.SO 2 .OH + C 6 H 6 = C 6 H 5 .SO 2 .OH + OH 2 .
Sulphuric, acid. Benzene. Benzenesulphonic acid.
The sulphonic adds, as the compounds thus produced are termed, are powerful monobasic acids. But the action may proceed further, both OH groups in sulphuric acid being displaced by monad compound radicles, in which case neutral bodies called sulphones are formed; thus— HO.SO 2 .OH + 2C 6 H 6 = C 6 H 5 .SO 2 .C 6 H 5 + 2OH 2 .
Sulphuric acid. Benzene. Benzenesulphone.
Selenic Add, H 2 SeO 4 .
This acid may be obtained by oxidizing selenious acid by the action of chlorine, or bromine, and water—
H 2 SeO 3 + Bγ 2 + H 2 O = H 2 SeO 4 + 2HBr.
It is a transparent colourless liquid, which in the most concentrated state boils at 280° C., and has a specific gravity of 2 6. In this state it is not quite pure selenic acid, but contains a little water, which cannot be driven off without decomposing the acid into selenium dioxide, oxygen, and water. Selenic acid resembles sulphuric acid in many of its properties, being very hygroscopic, and when it is added to water a considerable amount of heat is developed. It is reduced to selenious acid when boiled with hydrochloric acid—
H 2 SeO 4 + 2HC1 = H 2 SeO 3 + Cl 2 + H 2 O; but it is not decomposed by the hydrogen produced when zinc or iron is dissolved in its solution. Like sulphuric acid it precipitates barium salts, even in presence of other mineral acids.
The selenates or salts produced by displacing the hydrogen in selenic acid by metals correspond in composition, and also very closely in their properties, to the sulphates, with which they are isomorphous.
Telluric Add, H 2 TeO 4 .
The potassium salt of this acid is obtained by fusing tellurium or tellurous oxide with potassium nitrate. From this salt barium tellurate may be prepared by precipitating its solution with a barium salt, and the acid is produced by decomposing barium tellurate with sulphuric acid.
Telluric acid crystallizes from water in large prisms of the composition H 2 TeO 4 + 2H 2 O. This hydrate has a metallic taste, and reddens litmus slightly. It dissolves slowly in cold water, but freely in boiling water. It loses its water of crystallization at a little above 100° C. The compound H 2 TeO 4 is nearly insoluble in cold water, but dissolves on boiling; when heated to a temperature below redness it furnishes the oxide TeO 3 . It is reduced to tellurous acid by boiling with hydrochloric acid, and it is also decomposed by sulphurous acid, which is without action on selenic acid, with precipitation of tellurium;
In addition to acid and normal tellurates, such as KHTeO 4 and K 2 TeO 4 , and hyperacid tellurates, of which the salt KHTeO 4 + H 2 TeO 4 is an example, telluric acid also furnishes so-called di-and tetra-tellurates, such as (NH 4 ) 2 Te 2 O 7 and (NH 4 ) 2 Te 4 O 13 or (NH 4 ) 2 TeO 4 + 3TeO 3 , and basic tellurates, such as Ag 6 TeO 6 and Ag 6 Te 2 O 9 . The ditellurates may be regarded as salts of the acid H 2 Te 2 O 7 , analogous to anhydrosulphuric acid, and the basic salt Ag 6 TeO 7 as formed from the hydrate H 6 TeO 6 or H 2 TeO 4 + 2H 2 O; whilst salts such as Ag 6 Te 2 O 9 axe apparently derivatives of an acid formed by the withdrawal of the elements of three molecules of water from two molecules of the acid H 6 TeO β .
Other Adds of Sulphur.
In addition to sulphurous and sulphuric acids, a number of acids, more or less closely related to them, may be obtained. The following is a complete list of the known acids of sulphur :— [table]
Hyposulρhurous Add. — This acid is obtained by the action of zinc on a solution of sulphurous acid, which dissolves the metal, forming zinc sulphite; the hydrogen, which is the accessory product of this reaction, is not evolved, however, but reduces a portion of the acid, forming hyposulphurous acid; thus—
H 2 SO 3 + Zn = 2H + Z11SO 3
Sulphurous acid. Zinc sulphite.
H 2 SO 3 + 2ÏÏ = H 2 SO 2 + H 2 O.
Sulphurous acid. Hypostdphurous
The solution thus obtained is very unstable, and rapidly decomposes with separation of sulphur. Sodium hyposulphite, NaHSO 2 , is a more stable substance, and is produced in a similar manner by the action of zinc on a solution of sodium hydrogen sulphite. This salt crystallizes in slender colourless needles; it is soluble in water, but insoluble in alcohol. When exposed to the air in a moist state it becomes very hot, and is converted by oxidation into hydrogen sodium sulphite, but in the dry state it is not affected by oxygen. The remarkable observation has been made, however, that when a solution of sodium hyposulphite is oxidized by free oxygen, that is to say, when water saturated with oxygen is added to a solution of the hyposulphite, only one-half the oxygen is employed in causing the conversion of the hyposulphite into the sulphite, the remaining half becoming affixed to water, forming hydrogen dioxide; the reaction may be expressed by the following equation—
NaHSO 2 + O 2 + H 2 O = NaHSO 3 + II 2 O 2 .
Sodium bypwnlpMio. " 1 ≡°
Hyposulphurous acid has a much greater decolorizing and reducing power than sulphurous acid; it immediately reduces the metals from mercury and silver salts, and it precipitates copper hydride, Cu 2 H 2 , from a solution of copper sulphate.
Thiosulphuric Add, H 2 S 2 O 3 .
This acid is formed from sulphurous acid by combining it with sulphur; thus, when a solution of sodium sulphite is digested with sulphur, sodium thiosulphate is produced—
Na 2 SO 3 + S = Na 2 S 2 O 3 .
Sodium sulphite. Sodium thiosulphate.
It will be obvious that this reaction is precisely analogous to that which occurs when sodium sulphite is converted into sodium sulphate by the action of oxygen.
Thiosulphuric acid cannot be isolated, on account of its instability, and when an acid—hydrochloric acid, for example—is added to a solution of a thiosulphate, the thiosulphuric acid which is produced rapidly decomposes into sulphur and sulphurous acid: H 2 S 2 O 3 = H 2 SO 3 + S. The thiosulphates of alkali and alkaline earth metals are crystalline and soluble in water, and are fairly stable salts; the thiosulphates of the heavy metals, however, which are precipitated on the addition of solutions of metallic salts to a solution of sodium thiosulphate, are very unstable, and rapidly decompose into a metallic sulphide and sulphuric acid, thus—
Ag 2 S 2 O 3 + H 2 O = Ag 2 S + H 2 SO 4 . Silver thiosulphate. Silver sulphide. Sulphuric acid.
Sodium thiosulphate is largely employed in photography for the purpose of dissolving the chloride, bromide, and iodide of silver. These salts are readily soluble in a solution of the thiosulphate until they are exposed to light, when they become insoluble. If, therefore, a sheet of paper coated with one of these silver salts be exposed to light under an object which is in part transparent and in part opaque, such as a piece of lace, for example, the silver salt is rendered insoluble in those parts upon which the light has fallen, and by then immersing the paper in a solution of sodium thiosulphate the unaltered silver salt is removed, and the picture is thus “fixed.” The solvent power of the thiosulphate is due to the formation of a soluble silver sodium thiosulphate—
Na 2 S 2 O 3 + AgCl ≡ NaAgS 2 O 3 + NaCl; many other metallic salts insoluble in water are dissolved by a solution of sodium thiosulphate, probably also in consequence of the formation of similar mixed salts.
Anhydrosulphuric Acid, H 2 S 2 O 7 .
This acid is obtained by dissolving sulphuric anhydride in sulphuric acid in the requisite proportions: H 2 SO 4 + SO 3 = H 2 S 2 O 7 , It crystallizes in large colourless transparent prisms. When gently heated it is decomposed into sulphuric anhydride, which distils over, and sulphuric acid, which remains. Water dissolves it, forming sulphuric acid.
The so-called Nordhausen or fuming sulphuric acid consists chiefly of this acid. It is prepared at Nordhausen, in Saxony, by distilling an impure ferric sulphate, Fe 2 (SO 4 ) 3 , obtained by exposing ferrous sulphate, FeSO 4 , to a moderate heat in contact with the air,—the distillate, consisting chiefly of sulphuric anhydride, being received in sulphuric acid. This acid readily dissolves many hydrogenized carbon compounds which are only with difficulty acted upon by ordinary sulphuric acid, converting them into sulphonic acids; it is employed on this account to dissolve indigo, and in the preparation of anthraquinonedisul phonic acid, from which alizarin is artificially produced.
Salts of this acid, such as Na 2 S 2 O 7 , Ag 2 S 2 O 7 , BaS 2 O 7 , are obtained by combining the corresponding normal sulphates with sulphuric anhydride; water decomposes them, forming corresponding acid sulphates. The acid salt, KHS 2 O 7 , or hydrogen potassium anhydrosulphate, is obtained by dissolving the normal salt in anhydrosulphuric acid; it crystallizes in prisms.
The chloride of anhydrosulphuric acid, S 2 O 5 C1 2 , is produced by the action of sulphuric anhydride on many chlorinated compounds, thus—
2SO 3 + CC1 4 = S 2 O 5 C1 2 + COC1 2 Sulphuric carbon Anhydrosulphuric Carbonic anhydride. tetrachloride. chloride. chloride
It is a colourless oily liquid, of specific gravity 1 ’ 829 at 18° C., and boils at 146° C. In contact with water it decomposes slowly and noiselessly, forming sulphuric and hydrochloric acids, and is thus distinguished from sulphuric chlorhydrate, SO 2 (OH)C1, which is rapidly decomposed with almost explosive violence when thrown into water.
The lead chamber crystals (p. 504), when heated, furnish a magnificently crystalline body, which is also a derivative of anhydrosulphuric acid—
2 s ° 2 {θ⅛ = S 2 ° 5 (NO 2 ) 2 + H 2 Oj it distils at about 360 β C. without decomposing.
Dithionic Acid, H 2 S 2 O 6 .
The manganese salt of this acid is obtained by the action of a solution of sulphurous acid on manganese dioxide—
MnO 2 + 2H 2 SO 3 = MnS 2 O 6 + 2H 2 O.
This salt may be converted into the barium salt by treatment with barium hydroxide, and from a solution of the barium salt the acid is produced by adding exactly sufficient sulphuric acid to precipitate the barium. By concentrating the solution in a vacuum over sulphuric acid it may be obtained of specific gravity 1 - 347, but on further concentration the acid is resolved into sulphurous acid and sulphur dioxide—
H 2 S 2 O 6 = H 2 SO 4 + SO 2 .
A dilute solution decomposes in a similar manner when heated. In contact with the air dithionic acid is slowly oxidized to sulphuric acid.
The dithionates produced by decomposing barium dithionate, BaS 2 O β , with the corresponding sulphates, are all soluble in water and crystallize well. They exhibit considerable stability, but when heated are more or less readily converted into the corresponding normal sulphate and sulphur dioxide.
Trithionic Acid, H 2 S 3 O 6 .
Potassium trithionate is produced, together with potassium thiosulphate, when a saturated solution of hydrogen potassium sulphite is gently heated with sulphur; it is not improbable that the thiosulphate is a secondary product, formed by the action of hydrogen sulphide resulting from the action of the sulphur on the hydrogen potassium sulphite—
2HKSO 3 + 2S = K 2 S 3 O 6 + H 2 S.
Hydrogen potassium Potassium sulphite. trithionate.
Trithionic acid may be obtained from a concentrated solution of the potassium salt by precipitating the potassium by hydrofluosilicic acid. The solution has a sour and somewhat bitter taste, and is inodorous; it is permanent in a dilute state only, and on concentrating it in a vacuum over sulphuric acid it begins to decompose, even at 0° C., sulphur dioxide being evolved, and sulphur deposited, while sulphuric acid remains in solution. The salts of trithionic acid are but little known; they are extremely unstable, and even when boiled with water are decomposed and converted into sulphates, with evolution of sulphur dioxide and deposition of sulphur—
K 2 S 3 O 6 = K 2 SO 4 + SO 2 + S. Potassium trithionate. Potassium sulphate.
Tetrathianic Acid, H 2 S 4 O 6 .
The salts of this acid are produced by the action of iodine on thiosulphates; for example—
2Na 2 S 2 O 3 + I 2 = Na 2 S 4 Oθ + 2NaI. Sodium thiosulphate. Sodium tetrathionate.
The acid may be prepared from the barium salt by double decomposition with sulphuric acid. A dilute solution may be boiled without decomposing; but a concentrated solution is decomposed when heated, yielding sulphur, and sulphurous and sulphuric acids. Its salts are all soluble in water, and are much more stable than the trithionates, but for the most part their solutions cannot be evaporated without decomposition.
Pentathionic Acid, H 2 S 5 O 6 .
A solution of this acid is obtained by alternately passing sulphur dioxide and hydrogen sulphide gases through water—
5H 2 SO 3 + 5H 2 S = H 2 S δ O 6 + 9H 2 O + 5S; the solution may be concentrated by a gentle heat till it attains a specific gravity of 1*25-1'3, and may then be further concentrated in a vacuum to the specific gravity of about 1 ’ 6. The solution is colourless and inodorous, and has a strongly acid taste; it may be preserved unchanged at ordinary temperatures, but on heating a concentrated solution of the acid, hydrogen sulphide and sulphur dioxide are evolved, sulphur is deposited, and sulphuric acid remains.
The salts of pentathionic acid are so unstable that it is difficult to obtain them in the solid state. In their formation from the acid there is a great tendency for the fifth atom of sulphur to be separated, tetrathionates being produced, which have greater stability, and sometimes two atoms of sulphur are given up and trithionates are formed.
A number of the reactions involved in the formation of various sulphur compounds, and also of a few selenium and tellurium compounds, have been submitted to thermochemical investigation by Thomsen. The results deduced from his experiments are given in the following table :— [table]
The number 4510 for the reaction S, H 2 applies strictly only to sulphur in the state in which it separates when hydrogen sulphide is decomposed by iodine, and would require a small correction to make it applicable to rhombic sulphur.
On comparing the numbers representing the amounts of heat developed in the formation of sulphurous and sulphuric acids and of the corresponding selenium and tellurium compounds, thus— [table]
it will be evident that the affinity of selenium to oxygen, as measured by the heat developed, is less than that of sulphur to oxygen, and also less than that of tellurium to oxygen. In this respect, therefore, sulphur, selenium, and tellurium form a series corresponding to that in which chlorine, bromine, and iodine may be arranged. Bromine, we have seen, is intermediate in its properties between chlorine and iodine, and its atomic weight is also almost the mean of the atomic weights of these two elements, and its affinity to oxygen is less than that of either chlorine or iodine; the atomic weight of selenium, which is intermediate in its properties between sulphur and tellurium, is also nearly the mean of the atomic weights of these elements.
Constitution of the Sulphur Compounds.
The constitution of the compounds of sulphur with monad elements may be readily deduced; thus, the only formulae by which the chlorides of sulphur can be represented graphically are as follows :—
Cl
Cl—S—S—Cl Cl—S—Cl Cl—8—Cl
I
Cl
But the constitution of each of the oxides of sulphur may be expressed in two ways, thus—
S O—S—0
gCo xχ d x
0 II O=S=O O=S=O
Sulphur dioxide. Sulphur trioxide.
In like manner, two corresponding formulae may be assigned to each of the acids formed by combining these oxides with the elements of a molecule of water. Sulphuric acid, for instance, may be represented by either of the following formulae—
0
II
H—0—0—S—0—0—Il H—0—S—0—H
II o
But as compounds in which two or more atoms of oxygen are directly united together (for example, hydrogen dioxide, the higher oxides of chlorine, and chloric acid, as to the constitution of which there can be little doubt as they contain only monad elements associated with oxygen) readily decompose with separation of oxygen, the first of these formulae for sulphuric acid appears improbable on account of the stability of the acid and of most of the compounds derived from it. If, however, the second formula be adopted, it follows that the formula which represents the three atoms of oxygen in the trioxide as directly associated with the sulphur atom is the more probable. It is probable that sulphur dioxide and sulphurous acid are analogous in constitution to sulphur trioxide and sulphuric acid, and from the manner in which hyposulphurous acid is formed from sulphurous acid we may infer that it has the constitution represented by the second of the following formulae—
0 0
II II
H—O—S—O—H H—S—Ο—H
Sulphurous acid. Hyposulphurous acid.
The following formulae are the most probable expressions of the constitution of the remaining acids of sulphur; the constitution of the SO 2 group is the same as in sulphuric acid and the dots are used instead of lines—
HO.SO 2 .OH HO.SO 2 .SH so 2 ∙oh 2 2 SO 2 .OH
Sulphuric acid. Thiosulphuric acid. Dithionic acid.
SO 2 . OH s . so 2 . OH qZ¶ -s ° 2 ∙ 0h ? 5 S.SO 2 .OH ’ ς ∩ ∩ ττ
SO 2 .OH 2 ^S-SO 2 .OH
Trithionic acid. Tetrathionic acid. Pentathionic acid.
We have previously pointed out (p. 474) the value which is to be attached to constitutional formulae such as are here assigned to the sulphur compounds.
Nitrogen.
Symbol, N; Atomic wt., 14; Molecular wt., 28.
Nitrogen in the free state constitutes about one-fifth by volume of the atmosphere; in combination it occurs in nitrates and ammoniacal salts, and it enters into the composition of all animal and vegetable tissues.
It is best obtained from air by removing the oxygen by means of copper heated to redness; the air being first led through a solution of potassic hydrate to free it from carbon dioxide, and then through concentrated sulphuric acid to remove moisture, and when thus purified, passed through a tube containing metallic copper heated to dull redness. By passing a mixture of air and ammonia over heated copper, the copper oxide is reduced as fast as it is formed, and a short length of copper suffices for the preparation of an indefinite quantity of nitrogen; thus—
2Cu + O 2 = 2CuO, 3CuO + 2NH 3 = N 2 + 3H 2 O + 3Cu.
The supply of air is easily regulated, as an excess shows itself by tarnishing the surface of the copper.
The easiest method of obtaining pure nitrogen is to heat a solution of ammonium nitrate, which splits up into nitrogen and water: NO 2 NH 4 = N 2 + 2H 2 O. But as this salt is difficult to prepare, it is better to substitute for it a mixture of potassium nitrate and ammonium chloride, which together produce ammonium nitrate and potassium chloride—
KNO 2 + NH 4 C1 = NH 4 NO 2 + KC1.
Nitrogen is a colourless, inodorous, tasteless, incondensable gas; it is only very sparingly soluble in water, 100 volumes of water at 15° C. dissolving about one and a half volumes of nitrogen. It is incombustible, and does not support the combustion of ordinary combustibles; it is not poisonous, but an animal immersed in it dies simply for want of oxygen. Nitrogen evinces but little tendency to enter into reaction with other elements; titanium, tungsten, vanadium, and probably a few others combine directly with it, however, but its compounds are mostly produced by indirect means.
Atmospheric air consists not only of nitrogen and oxygen, which are its chief constituents, but besides these contains carbon dioxide, ammonia, water vapour, &c.; solid substances, such as common salt, are also frequently held in suspension by it, especially in the neighbourhood of the sea and of towns. Air from which all other constituents are removed does not always exhibit the same composition, however, although the variations are very slight; usually in pure air the proportion of oxygen is from 20 ’ 9 to 21 volumes in 100 of air, but considerably less oxygen has been found in air from confined spaces and in a few samples collected in warm countries. Full information on this subject may be obtained from Dr Angus Smith’s work on Air and Rain.
The fact, however, that oxygen and nitrogen are not always contained in air in the same proportions is alone a sufficient proof that they are not combined, but only mixed together, as the constituents of a compound always occur in invariable proportions. This conclusion is confirmed in many ways. Thus, the proportions in which nitrogen and oxygen are present in air are not those of their atomic weights, nor do they bear any simple relation to them. We know also that when gases enter into reaction to form new compounds their combination is usually attended with an alteration of volume, and heat is developed; moreover, the resulting compound possesses properties which differ strikingly from those of its constituents. But when oxygen and nitrogen are mixed together in the proportions in which they are present in air, neither is any alteration in volume, observed, nor is heat developed; and the properties of the resulting mixture are precisely those of air, and just such as we should expect to result from the admixture of a gas which very readily supports combustion with one in which combustion is impossible. Again, were air a compound it should dissolve in water as such, or in other words, the proportion of oxygen and nitrogen in the dissolved air should· be the same as in the undissolved air; but if a mixture, the more soluble constituent should dissolve the more readily, and relatively more oxygen than nitrogen should dissolve, since oxygen is more soluble than nitrogen. Experiment proves that the latter is the case, for if water which has been recently boiled to free it from dissolved gases and allowed to cool out of contact with air be shaken with air, and the dissolved air be then expelled from it by heating, and collected, it is found on analysis to contain 32 instead of only 21 per cent, by volume of oxygen. Lastly, we have seen (p. 481) that the oxygen may be to a great extent separated from the nitrogen by a mechanical process by submitting air to filtration through a thin caoutchouc membrane.
It is impossible for animals to live for any length of time in pure oxygen, apparently because oxidation takes place so rapidly that the animal is incapable of assimilating sufficient food to supply the waste; but by admixture with the perfectly neutral nitrogen the activity of the oxygen becomes greatly diminished.
Air from open places contains usually from 3 to 6 volumes of carbon dioxide in 10,000 volumes, but the amount of carbon dioxide in the atmosphere is subject to continual change, although within narrow limits. It does not continually increase in amount, notwithstanding that animals expire carbon dioxide, and that large quantities are produced by the combustion of wood and coal, because plants exercise a power which is the reverse of that of which animals are possessed, viz., that of decomposing carbon dioxide and restoring its oxygen to the atmosphere.
Ammonia, NH 3 ; Mol. wt., 17.
Ammonia is the only compound of nitrogen and hydrogen which is known to us. It may be formed directly from its elements by submitting a mixture of the two gases to the action of the silent electric discharge; the combination is very imperfect, however. It is also obtained by the action of nascent hydrogen on nitric acid and many of its salts; but it is always prepared by heating an ammonium salt with an alkaline hydroxide, ammonium chloride and calcium hydroxide or slaked lime being the substances which are usually employed—
2NH 4 Cl + Ca(OH) 2 = 2NH 3 + 2OH 2 + CaCl 2 .
Ammonium calcium A mm r>r>i<> Calcium chloride hydroxide. Ammonia. chloride.
Ammonia is a product of the decay of all nitrogenous animal and vegetable substances, and the ammonia present in the atmosphere is chiefly if not entirely derived from this source; but it appears not improbable that ammonia is directly produced in the atmosphere by the decomposition of water into its elements by the electric discharge, and the combination also under the influence of the discharge of the hydrogen thus produced with nitrogen. Ammonia is present in the atmosphere apparently as carbonate, and in rain-water, especially in that of thunder showers, as nitrate and nitrite. Ammonia salts are sometimes found as minerals, chiefly in volcanic districts. The source from which ammonia salts are now obtained, however, is the watery liquid which distils over in the manufacture of coal gas; the ammonia is liberated from this liquid by heating it with slaked lime, and by receiving it in sulphuric acid ammonium sulphate is produced,—a salt which is largely employed as an artificial manure.
Ammonia is a transparent colourless gas, of a very pungent peculiar odour, and a burning taste. It may be reduced to the liquid state by a pressure of about 17 atmospheres at the ordinary temperature, or by cold alone at about - 50° to - 40° C.; by exposing the dry gas to a cold of — 75° C. and a pressure of 20 atmospheres, Faraday obtained ammonia as a white transparent crystalline body, melting at — 75° C. It does not support combustion, and is only feebly combustible. It is decomposed into its elements by a succession of electric sparks.
Ammonia is dissolved by water with great avidity, much heat being developed and great expansion taking place; according to Roscoe and Dittmar, 1 gramme of water at 0° C. dissolves no less than ’ 875 gramme of ammonia. The solution has the smell and taste of the gas, and a powerfully alkaline reaction; it loses almost all its ammonia below 100° C.
Ammonia completely neutralizes acids, forming definite crystalline salts, known as ammonium salts, which are formed by the direct combination of ammonia with the acids; ammonia and hydrochloric acid, for example, form ammonium chloride, NH 4 C1 = NH 3 + HC1, whilst ammonia and sulphuric acid furnish ammonium sulphate, (NH 4 ) 2 SO 4 = 2NH 3 + H 2 SO 4 . The constitution of these salts, it will be evident, is analogous to that of the salts of the metals generally if we regard them as derived from the acids by the displacement of the hydrogen of the latter by the monad compound radicle ammonium, NH 4 , and this view is confirmed by the observation that the ammonium salts are isomorphous with the corresponding potassium salts. They are all soluble in water, and are readily decomposed by the alkaline hydroxides, and by most basic oxides, with evolution of ammonia.
The solution of ammonia in water is frequently regarded as the hydroxide of the hypothetical radicle ammonium, that is to say, as a solution of ammonium hydroxide, NH 4 .OH, the analogue of potassium hydroxide, KOH. But Thomsen has shown that considerably less heat is developed when an acid is neutralized by a solution of ammonia than when it is neutralized by a solution of an alkaline hydroxide, such as potassium hydroxide, for example, the reaction KHOAq, HCLAq being accompanied by the development of 27,500 units of heat, whilst only 24,500 are developed in the reaction NH 3 Aq, HClAq. Thomsen is inclined to regard this result as evidence that ammonia exists as such iu its aqueous solution, and not as the hydroxide NH 4 .OH, but the difference observed appears scarcely sufficient to warrant this conclusion iu the face of the chemical evidence which points to the existence of an ammonium hydroxide; it is more probable, perhaps, that an aqueous solution of ammonia consists in part of the hydroxide and in part of free ammonia.
A very large number of derivatives may be obtained from ammonia by displacing one or more atoms of hydrogen in it by positive or negative radicles. Those which are formed by the introduction of monad positive radicles are distinguished by the name amines, whilst those containing monad negative radicles are called amides ; when two atoms of hydrogen in a single molecule of ammonia are displaced by dyad negative radicles so-called imides are produced. A simple instance of the formation of an amine derivative is afforded by the action which takes place when potassium is heated in an atmosphere of ammonia—
2NH 3 + 2K = 2NH 2 K + H 2 . Ammonia. Potassaπιine,
By digesting iodine in an excess of aqueous ammonia a black explosive compound is produced, which apparently is formed by the displacement of two of the atoms of hydrogen by iodine—
3NH 3 + 2I 2 = NHI 2 + 2NH 4 I. Ammonia. Diiodamide.
It may be exploded by friction evèn under water, and in the dry state can scarcely be touched without exploding. When chlorine is passed into an aqueous solution of ammonia, nitrogen is evolved and ammonium chloride is produced; thus—
2NH 3 + 3C1 2 = N 2 +6HC1 6NH 3 + 6HC1 = 6NH 4 Cl .∙. 8NH 3 + 3C1 2 = N 2 + 6NH 4 C1.
If the action of the chlorine be continued after the whole of the ammonia is thus acted upon, the ammonium chloride becomes attacked, and yellow oily drops of the so-called chloride of nitrogen are formed. The composition of this substance has not yet been satisfactorily ascertained; it probably contains hydrogen, and its formula is supposed to be NHC1 2 , but it is not unlikely that the compound NC1 3 is also produced. It explodes with extreme violence when heated, or when brought in contact with fatty matters, or with turpentine, phosphorus, and many other substances.
The Oxides and Acids of Nitrogen.
No less than five oxides of nitrogen are known, viz.—
N s O; NO; N 2 O 3 ; N 2 O 4 ; N 2 O 5 .
Nitrous Nitric Nitrous Nitric Nitric oxide. oxide anhydride. peroxide. anhydride.
Nitrous and nitric oxide are merely dissolved by water, but the remaining oxides enter into reaction with it, producing acids,—nitrous and nitric anhydride being converted into the corresponding acids, nitrous and nitric acid, and a mixture of those two acids being formed from the intermediate oxide, nitric peroxide; thus—
N 2 O 3 + H 2 O = 2HNO 2 Nitrous anhydride. Nitrous acid.
N 2 O 4 + H 2 O = HNO 2 + HNO 3 Nitric peroxide. “Nitrous acid. Nitric acid.
N 2 O 5 + H 2 O = 2HNO 3 Nitric anhydride. Nitric acid.
As these various oxides of nitrogen are all prepared from nitric acid, we may conveniently describe this compound first.
Nitric Acid, HNU 3 .
This acid is now usually prepared by distilling sodium nitrate or Chili saltpetre, a salt which occurs abundantly in South America, with concentrated sulphuric acid in the proportions indicated by the equation—
2NaNO 3 + H 2 SO 4 = 2HNO 3 + Na 2 SO 4 .
When potassic nitrate is employed, it is advantageous to use double the quantity of sulphuric acid; thus —
KNO 3 + H 2 SO 4 = HNO 3 + KHSO 4 , as the heat required to produce the reaction—
KNO 3 + KHSO 4 = HNO 3 + K 2 SO 4 is so great that a considerable amount of the acid is decomposed.
The acid passes over as an almost colourless liquid in the middle of the process, but is coloured at the beginning and end by decomposition products. To obtain the pure acid, it is mixed with an equal bulk of concentrated sulphuric acid, and the mixture distilled; the first portions of the distillate are collected, and a current of dry air passed for several hours through the liquid, which is gently warmed and sheltered from strong daylight.
Pure nitric acid is a colourless, mobile, fuming liquid, of the specific gravity 1 *53 at 15° C.; it solidifies at about 55° C. It is an extremely unstable substance, and cannot be distilled without experiencing partial decomposition; it is also decomposed when exposed to sunlight, becoming yellow, and oxygen being evolved. It begins to boil at 86° C., but the temperature rises gradually, and oxygen and red fumes of oxides of nitrogen are evolved; when the boiling point reaches about 123° C., an aqueous acid, having a specific gravity of about 1 ’ 42, and containing about 68 per cent, of nitric acid, distils unchanged, and weaker and stronger acids may alike be reduced to this strength by boiling. As in the case of other aqueous acids, however, the composition of nitric acid of constant boiling point varies with the pressure under which ebullition takes place.
The addition of water to nitric acid causes the development of heat. The following table shows the extent to which this occurs, the amount of heat developed on adding n molecules of water to a molecule of the acid being given in the second column of the table, whilst the third exhibits the amount developed on adding a quantity of water equal to that already present :— [table]
These numbers do not appear to furnish any evidence of the formation of a distinct hydrate on adding water to nitric acid.
Nitric acid is a monobasic acid, and furnishes an important class of salts called nitrates, such as potassium nitrate, KNO 3 , copper nitrate, Cu(NO 3 ) 2 , and bismuth nitrate, Bi(NO 3 ) 3 . In addition to these normal nitrates, a number of so-called basic nitrates are known which may be regarded as normal salts of a distinct acid, formed by the combination of a molecule of nitric acid with a molecule of water; such are the basic bismuth nitrate BiNO 4 and the basic lead nitrates Pb 2 H 2 N 2 O 8 and Pb 3 N 2 O 8 . The normal nitrates are best prepared by dissolving the metallic oxides, hydroxides, or carbonates in diluted nitric acid; they are all soluble in water. The most important nitrate is that of potassium, which is employed in the manufacture of gunpowder.
Nitric acid is an extremely powerful oxidizing agent, and one of the most corrosive substances known; it rapidly destroys all animal textures and most vegetable products, and even if diluted it stains the skin, wool, and all albuminous substances a bright yellow colour. Many hydrogenized carbon compounds are converted by it into so-called nitro-derivatives, one or more atoms of hydrogen being displaced by the monad compound radicle NO 2 ; the hydrocarbon benzene, for example, when added to the concentrated acid yields nitrobenzene—
C 6 H 6 + HNO 3 = C 6 H 5 .NO 2 + H 2 O.
Most metals, excepting gold, platinum, rhodium, iridium, titanium, and perhaps a few others, are more or less readily acted on by nitric acid and converted into nitrates; but the non-metallic elements and metalloids—iodine, sulphur, selenium, tellurium, phosphorus, arsenic, antimony, amorphous boron and carbon, and tungsten—are oxidized by it, and furnish iodic, sulphuric, selenious, tellurous, phosphoric, arsenic, antimonic, boric, carbonic, and tungstic acids.
The action of nitric acid on metals, however, is much influenced by temperature and concentration. An acid of the specific gravity 1 ’ 25 to 1 ’ 35 is usually the most active. The pure concentrated acid is without action on bismuth, iron, tin, and many other metals at ordinary temperatures. Thus, a piece of bright iron is at once attacked by an acid of about the specific gravity 1'35, but it maybe preserved in acid of the specific gravity I *45 without losing its brilliancy; what is more remarkable, however, is that by plunging it into the more concentrated acid it is rendered passive to the action of the weaker acid, for on removing it and at once introducing it into the weaker acid, no action is observed, although on diluting the acid below 1 ‘35 the iron is attacked.
In order, however, to understand the behaviour of nitric acid with metals, it is necessary that we should first consider the action of metals upon acids generally. There is little doubt that in all cases the metal simply displaces the hydrogen of the acid, forming a salt; and if, under the conditions under which the experiment is made, the acid has no tendency to enter into reaction with the hydrogen which is displaced, whilst it is in the nascent state, hydrogen is also evolved; but if the acid can enter into reaction with the nascent hydrogen the products of this secondary reaction are obtained instead. For example, zinc readily dissolves in cold diluted sulphuric acid, forming zinc sulphate, and hydrogen is evolved since it is without action on sulphuric acid under these conditions—
Zn + H 2 SO 4 = ZnSO 4 + H 2 .
But when zinc and concentrated sulphuric acid are heated together, zinc sulphate and sulphur dioxide are obtained, and no hydrogen is evolved. In this case, the hydrogen is displaced from the sulphuric acid under conditions which are favourable to its action upon the acid, and it deprives the acid of a portion of its oxygen, forming sulphurous acid, which is at once resolved into water and sulphur dioxide.
The behaviour of nitric acid with metals is precisely similar to that of zinc with heated concentrated sulphuric acid, nitric acid being a substance which with the greatest readiness may be deprived of its oxygen by the action of nascent hydrogen. In fact, hydrogen is never evolved by the action of metals on nitric acid, but instead oxides of nitrogen, nitrogen itself, ammonia, and other products of secondary reaction are obtained. The formation of these products is due to a somewhat complex series of reactions, which most probably are as follows. In the first place, by the removal of one of the atoms of oxygen the nitric acid is converted into nitrous acid, HNO 2 by further reduction this acid furnishes the so-called hyponitrous acid, HNO; and by the continued action of the nascent hydrogen the hyponitrous acid is converted into hydroxylamine or oxyammonia, NH 2 (OH), which finally is reduced to ammonia, NH 3 . The reactions which successively occur in the formation of these compounds are represented by the following equations :— [table]
But the products of these reactions enter into reaction with each other, and are decomposed, and thus the gases are produced which are evolved when metals are dissolved in nitric acid. The nitric oxide doubtless results chiefly from the decomposition of the nitrous acid in the manner represented by the equation—
3ΞNO 2 = 2NO + HNO 3 + H 2 O.
The nitrous oxide may be produced in two ways :—from ∙hyponitrous acid, which immediately on formation is resolved into nitrous oxide and water: 2HNO = N 2 O + H 2 O; and by the action of nitrous acid on hydroxylamine—
HNO 2 + NH 2 (OH) = N 2 O + 2H 2 O .
The nitrogen produced is probably formed in a similar manner by the action of the nitrous acid on the ammonia—
HNO 2 + H 3 N = N 2 + 2H 2 O; and perhaps also, together with nitric oxide, by the action of nitric acid on hyponitrous acid.
As the product of the action of nitric acid on a metal is always a mixture, it is evident that several of the reactions pointed out must occur simultaneously. The composition of the product varies, in a manner not yet understood, with the metal, the strength of the acid, and with the temperature.
The approximate percentage composition of the gas obtained by the action of a mixture of nitric acid of the specific gravity 1 ’ 42 with twice its bulk of water on a number of metals is given in the following table :— [table]
It is difficult at present to explain the remarkable difference in the behaviour of the various metals with nitric acid which this table indicates. As we have stated above, the action of the metal probably consists simply in displacing the hydrogen of the acid, and the gases evolved are the more or less direct products of the action of the hydrogen so displaced at the moment of liberation on the acid which is always present in excess. The question, therefore, that we have to consider is, Why does the hydrogen displaced from nitric acid by different metals produce such very different effects î If the behaviour of closely related metals such as magnesium, zinc, and cadmium be compared, it is evident that magnesium is the most active, since it produces a relatively smaller quantity of nitric oxide, and relatively larger quantities of nitrous oxide and nitrogen, than either zinc or cadmium, cadmium being the least active; in other words, the reduction of the nitric acid is most perfect when it is effected by the hydrogen displaced by the aid of magnesium, and. least perfect when it is effected by the aid of cadmium. We know that when these metals act upon acids which are not attacked by nascent hydrogen, such as hydrochloric acid, for example, different amounts of heat are developed, most heat being developed by the action of magnesium, and least by the action of cadmium. But since the mere displacement of hydrogen in hydrochloric acid by different metals is attended with the development of different amounts of heat, it appears probable that the same will be the case with nitric acid, and also that more heat would be developed in the reaction
Mg + 2HNO 3 = Mg(NO 3 ) 2 + 2H
than in the corresponding reactions with either zinc or cadmium; and that on this account the reduction of nitric acid is carried furthest when magnesium is employed, and further with zinc than with cadmium. Whether this explanation is applicable to the action of metals generally there is not sufficient evidence at present to show.
Nitrous Oxide, N 2 O; Mol. wt., 43 96.
This gas is obtained in a pure state by carefully heating ammonium nitrate, which at a temperature between 200° and 250° C. breaks up into water and nitrous oxide—
NH.NO 3 = N 2 O + 2H 2 O. ⅛ o A A
It is the chief constituent of the gas produced on dissolving magnesium, zinc, tin, nickel, or cobalt in nitric acid, and it is also furnished, but to a less extent, by other metals. . .. ,
Nitrous oxide is a transparent colourless gas, possessing a faint sweetish smell and taste; it may be condensed by a pressure of about 50 atmospheres at 7° C. to a colourless liquid, and even frozen by the cold produced by its own evaporation. By exposing a mixture of liquid nitrous oxide and carbon disulphide to evaporation in vacuo, Natterer obtained a reduction of temperature which he estimated at - 140° C., which is lower than has been obtained by any other means. 100 volumes of water at 0° C. dissolve 130 of nitrous oxide, but the solubility diminishes rapidly as the temperature rises. It supports the combustion of bodies almost, if not quite, as well as oxygen, but is readily distinguished from that gas by its solubility in water, and by not forming red fumes when mixed with nitric oxide gas.
The most remarkable property which nitrous oxide gas possesses is that of causing loss of sensibility in animals. When respired for a short time it produces a singular species of transient intoxication, attended in many instances with an irresistible tendency to muscular exertion, and often to uncontrollable laughter; hence it has acquired the name of laughing gas. It is now largely used as an anaesthetic for producing insensibility to pain during surgical operations, and especially in the extraction of teeth.
Nitrous oxide is a perfectly neutral substance, and does not exhibit any tendency to enter into combination with other bodies.
Nitric Oxide, NO; Moh wt., 29’96.
This gas is one of the products of the action of nitric acid on metals; the gas obtained on dissolving copper in moderately concentrated nitric acid contains at first about 97 per cent, of nitric oxide, but as the action proceeds and cupric nitrate accumulates in the solution, the proportion of nitric oxide evolved diminishes, whilst that of nitrous oxide increases. This appears to be due to the alternate reduction of the cupric nitrate to cuprous nitrate—
2Cu(NO 3 ) 2 + H 2 = Cu 2 (NO 3 ) 2 + 2HNO 3 cupric nitrate. Cuprous nitrate.
and reconversion of the latter into cupric nitrate, as by the action of nitric acid on cuprous nitrate nitric oxide mixed with a large proportion of nitrous oxide is produced.
Pure nitric oxide may be prepared by gently heating a mixture of a ferrous salt, such as ferrous chloride or ferrous sulphate, with hydrochloric or dilute sulphuric acid and potassium or sodium nitrate. The nitric acid liberated from the nitrate is reduced by the ferrous salt, which is oxidized to a ferric salt, and nitric oxide is evolved—
6FeCl 2 + 6HCl + 2HNO 3 = 2NO + 3Fe 2 Cl 6 + 4H 2 O. Ferrous chloride. Nitric acid. Nitric oxide. Ferric chloride.
Nitric oxide is a colourless, transparent, uncondensable gas, almost insoluble in water; it is the most stable of the oxides of nitrogen, and may even be exposed to a red heat without undergoing decomposition. A lighted taper and phosphorus just kindled are extinguished by it, but if the phosphorus be burning vigorously when introduced into the gas the temperature is then sufficiently high to enable it to decompose the gas into its elements, and it continues to burn as brilliantly as in pure oxygen.
Nitric oxide immediately combines with oxygen when mixed with it, forming deep red fumes of higher oxides of nitrogen. It is perfectly absorbed by a solution of ferrous chloride or sulphate, which it turns black; when the solution is heated, most of the nitric oxide is expelled from it unchanged, and in this way nitric oxide may be separated from other gases insoluble in solutions of the ferrous salts. Nitric oxide is absorbed by nitric acid, and apparently nitrous acid is produced—
HNO 3 + 2NO + H 2 O = 3HNO 9 . o Δ Δ
It combines directly with chlorine, forming the compound NOC1, or nitrosyl chloride, which is also produced together with chlorine when a mixture of concentrated nitric and hydrochloric acids—the so-called aqua regia—is heated—
HNO 3 + 3HC1 = NOC1 + Cl 2 + 2H 2 O.
Nitrosyl chloride is most readily prepared, however, by gently heating a mixture of dry sodium chloride and the lead chamber crystal compound (p. 504)—
SO 2 1° + NaCl = NOC1 + SO 2 1 θg a
It is an orange-yellow coloured gas, which readily liquefies; the liquid has a deep orange colour, and boils at about - 8° C. It is decomposed by water, forming hydrochloric and nitrous acids—
NOC1 + HOH = NO. OH + HC1.
Nitroβyl chloride. Nitrous acid.
Nitrous Anhydride, N 2 O 3 .
The properties of this compound have not yet been satisfactorily ascertained. Apparently, it is produced, together with nitric peroxide, when nitric oxide is mixed with oxygen, and when nitric acid is heated with arsenious anhydride; it is stated that it may be prepared in a pure state by passing a mixture of mtric peroxide with an excess of nitric oxide through a heated tube, and that it may be condensed to a deep blue coloured liquid, which begins to boil at about 2° C., but is decomposed into nitric oxide and peroxide. It combines readily at ordinary temperatures with oxygen, forming nitric peroxide.
Nitric Peroxide, N 2 O 4 j Mol. wt., 91’86.
Nitric peroxide may be obtained by combining nitric oxide with oxygen, and by heating lead nitrate—
2Pb(NO 3 ) 2 = 2N 2 O 4 + 2PbO + O 2 . Lead nitrate."" Nitric peroxide. Lead oxide.
At low temperatures it forms colourless prismatic crystals, which melt at - 9° C.; at this temperature the liquid oxide is almost colourless, but it assumes a yellow colour, increasing in depth up to its boiling point. It boils at about 25° C., yielding a brownish-red vapour, the depth of which also increases with the temperature, until at 40° C. it is so dark as to be almost opaque. The vapour of nitric peroxide, however, probably even at temperatures below its boiling point, is a mixture of the compound N 2 O 4 with simpler molecules of the composition NO 2 ; the proportion of the latter increases as the temperature rises, and at temperatures above 150° C. the decomposition into NO 2 is apparently complete. The compound NO 2 is not decomposed at a dull red heat.
Nitric peroxide has a pungent suffocating odour, and is quite irrespirable. It is decomposed by water with production of nitrous and nitric acids: N 2 O 4 + H 2 O = HNO 2 + HNO 3 . Chlorine is without action on nitric peroxide in the cold, but when a mixture of the two gases is passed through a heated tube the compound NO 2 C1 or nitrylic chloride is formed. It is a pale yellow liquid, which is decomposed by water into nitric and hydrochloric acids.
Nitric Anhydride, N j O 5 .
This compound may be obtained directly from nitric acid by the action of phosphoric anhydride—
2HNO s + P 2 O 5 = N 2 O 5 + 2HPO 3 . xτiM. -- jλ Phosphoric Nitric Metaphosphoric
Nitπc aeιd. anhydride . anh ydride. acid.
It may also be obtained by the action of chlorine on silver nitrate—
2AgNO 3 + 2C1 2 = 2NO 2 C1 + 2A g Cl + O 2 AgNO 3 + NO 2 Cl =N 2 O 5 + AgCl.
Nitric anhydride crystallizes in transparent prisms, which melt at about 30° C.; it is very unstable, particularly in the fused state, and gradually decomposes even at ordinary temperatures into oxygen and nitric peroxide. Water converts it into nitric acid: N 2 O 5 + H 2 O = 2HNO 3 .
Nitrous Acid, HNO 2 .
In the pure state this compound is unknown to us, but its salts, the nitrites, may readily be prepared. Thus, on heating sodium or potassium nitrate to redness until the gas which is evolved begins to contain nitrogen, a residue which consists chiefly of sodium or potassium nitrite is obtained; these nitrites are soluble in alcohol, and may by its aid be separated from the nitrates, which are insoluble. A mixture of nitrate and nitrite is also produced on passing the gas evolved when nitric acid is heated with arsenious anhydride, starch, or sugar into a solution of potassium or sodium hydroxide, and when oxygen is added to nitric oxide which is in contact with a solution of an alkali. On adding silver nitrate solution to a concentrated solution of the impure nitrite, silver nitrite, AgNO 2 , which is difficultly soluble in water, is precipitated; this salt may be purified by crystallization from a large quantity of hot water, and from it other nitrites may be prepared in a pure state by double decomposition with the corresponding chlorides; thus—
AgNO 2 + KC1 = AgCl + KNO 2 .
Silver Potassium Silver Potassium nitrite. chloride. chloride. nitrite.
Although oxygen and nitrogen do not directly combine, when a succession of electric sparks is passed through a mixture of the two gases confined over water, red fumes are produced which dissolve in the water, forming nitric and nitrous acids. The nitrite and nitrate present in the atmosphere and in rain-water are doubtless formed in this manner.
On adding an acid to a nitrite, nitrous acid is liberated, but at once decomposes with evolution of nitric oxide—
KNO 2 + HC1 - HNO 2 + KC1 3HNO 2 = 2NO + HNO 3 + OH 2 .
Nitrous acid can only exist, in fact, in presence of a large quantity of water, or of nitric acid. It parts with its oxygen more readily, and is therefore a more powerful oxidizing agent, than nitric acid. It also appears to act more readily than nitric acid upon many metals. Thus, pure nitric acid is almost without action upon silver, but on passing a few bubbles of nitric oxide gas into the acid, and thus producing a minute quantity of nitrous acid, the metal is at once attacked, and is dissolved more and more rapidly the longer the action continues, doubtless because the hydrogen displaced from the acid by the silver reduces nitric acid to nitrous acid, so that the more silver is dissolved the richer the solution becomes in nitrous acid. Platinum also, which is insoluble in nitric acid, is dissolved by nitrous acid.
On some hydrogenized carbon compounds nitrous acid exerts an action similar to that of nitric acid, and causes the displacement of hydrogen by the monad compound radicle NO, producing mïroso-compounds. Thus—
C 6 H 5 .OH + NO. OH = C 6 H 4 (NO).OH + H 2 O. Phenol. Nitrous acid. Nitrosophenol.
Nitrous acid readily absorbs oxygen, and is converted into nitric acid. The nitrites are also converted into nitrates when exposed in the moist state to air. Nitrous acid is easily distinguished from nitric acid by its power of liberating iodine from iodides, by the readiness with which it destroys the blue colour of a solution of indigo at ordinary temperatures, and by its decolorizing a solution of potassium permanganate,—nitric acid being without the power of liberating iodine from iodides, or of bleaching permanganate solution.
Hyponitrous Acid, HNO.
When sodium is added to a solution of sodium nitrate, the hydrogen displaced from the water by the sodium reduces the nitrate to nitrite, which in its turn undergoes further reduction to the hyponitrite, NaNO. On rendering the solution slightly acid with acetic acid, and adding silver nitrate, a yellow pulverulent precipitate of silver hyponitrite, AgNO, is produced. It is insoluble in water, and almost insoluble in acetic acid, but is soluble in either dilute nitric or sulphuric acid, and without immediate decomposition. Moderately diluted nitric, sulphuric, or hydrochloric acid decomposes it with the evolution of nitrogen, and the production of apparently both nitrous and nitric acids in the solution. It is immediately oxidized by concentrated nitric acid. A solution of the sodium salt acidified with acetic or hydrochloric acid decolorizes potassium permanganate; it does not liberate iodine from iodides, however, but on the contrary decolorizes a solution of iodine. When the solution acidified with acetic acid is heated, nitrous oxide is evolved.
Hydroxylamine, NH 2 (OH).
This compound is formed when nitric acid is added to a mixture of tin and hydrochloric acid, the hydrogen produced by the action of the tin on the hydrochloric acid reducing the nitric acid. It may also be formed by the direct union of nitric oxide with hydrogen, namely, by passing a stream of nitric oxide gas through a series of glass cylinders containing tin and hydrochloric acid, together with a little platinum chloride solution, whereby hydrogen is produced at ordinary temperatures.
It is a very unstable substance, and can be obtained only in solution, but well-crystallized hydroxylammonium salts are formed by its union with acids. Thus, the hydrochloride, NH 3 C1(OH), or NH 2 (OH). HC1, crystallizes from alcohol in long spicular crystals, and from water in large irregular monoclinic prisms; it melts at 100° C., but decomposes, with violent evolution of gas, into nitrogen, ammonium chloride, water, and hydrochloric acid. A solution of hydroxylamine has an alkaline reaction, and precipitates many metallic salts; it decomposes quickly if concentrated, and gradually if dilute, with evolution of nitrogen and formation of ammonia—
3NH 2 (OH) = N 2 + NH 3 + 3OH 2 .
Hydroxylamine is readily reduced tc ammonia by the action of the nascent hydrogen from sodium amalgam and water.
The results of Thomsen’s thermochemical investigation of various nitrogen compounds are collected in the following table :— [table]
It will be observed that a considerable amount of heat would be absorbed in the formation of nitrous oxide from its elements; consequently, when this gas is decomposed into its elements, heat is developed, and on this account it is readily decomposed by burning bodies. This fact also explains the non-formation of nitrous oxide from its elements. Probably, heat would also be absorbed in the formation of nitric oxide from its elements. The absorption of heat is perhaps necessary, because the amount of energy in the form of heat which can be developed by the combination of nitrogen with oxygen to form nitrous and (1) nitric oxides is less than must be expended in separating the atoms of oxygen of the oxygen molecules from each other, and the atoms of nitrogen of the nitrogen molecules from each other.
Phosphorus.
Symbol, P; At. wt., 30'96; Mol. wt. of gas, 123'84.
This element never occurs in the free state, but is always found in combination as a salt of phosphoric acid. Considerable deposits of more or less pure calcium phosphate occur in a few places, and phosphates are found in minute proportions in most rocks and in soils, and in river and spring waters. Phosphates are necessary to the life of all plants and animals. In plants they accumulate chiefly in the seed; in animals they accumulate in the bones, of which calcium phosphate is the chief earthy constituent, but they are also an important element of blood and nervous tissue.
Phosphorus was discovered in 1669 by Brandt of Hamburg, who obtained it by distilling the residue of evaporated urine with charcoal. Scheele in 1775, however, was the first who devised a process for its extraction from bones, and it has always since been prepared from this source. The bones are burnt to a white ash, which is finely powdered and mixed with a sufficient quantity of diluted sulphuric acid to displace by hydrogen two-thirds of the calcium in the tricalcium phosphate, which is the main constituent of bone ash, in the manner represented by the equation—
Ca 3 (PO 4 ) 2 + 2H 2 SO 4 = H 4 Ca(PO 4 ) 2 + 2CaSO 4 .
Tricalcium Sulphuric Monocalcium Calcium phosphate. acid. phosphate. sulphate.
The solution of monocalcium phosphate, or superphosphate of lime as it is ordinarily termed, which is obtained, is separated from the insoluble calcium sulphate, and evaporated to a syrup; this is mixed with about onefourth its weight of charcoal powder, and heated gradually to dull redness in an iron pot with constant stirring. By this means the elements of two molecules of water are removed, the monocalcium phosphate being converted into calcium metaphosphate—
CaH 4 (PO 4 ) 2 = Ca(PO 3 ) 2 + 2H 2 O.
The porous mixture of calcium metaphosphate and charcoal thus obtained is introduced into earthen retorts, and distilled at a bright red heat. The calcium metaphosphate is then converted into tricalcium phosphate and phosphoric anhydride, and the latter is decomposed by the charcoal into carbonic oxide and phosphorus, which passes over in vapour and is condensed in water—
3Ca(PO 3 ) 2 = 2P 2 O 5 + Ca 3 (PO 4 ) 2 P 2 O δ + 5C = 2P + 5CO.
The crude phosphorus which is produced is usually purified by adding a mixture of potassium chromate and sulphuric acid to it when in a melted state. The impurities are thus oxidized, and rise to the surface in the form of a scum, and the pure phosphorus remains colourless and transparent at the bottom of the vessel.
Freshly-prepared phosphorus is almost perfectly transparent and colourless, or has, at most, a faint yellow tinge. It melts at 44° C., forming a viscid oily liquid, and boils at 290° C. Its vapour at a temperature of 1040° C. is about 62 times as heavy as hydrogen gas, and hence, as the atomic weight of phosphorus is about 31, it appears that the molecule of gaseous phosphorus is tetratomic. Phosphorus has a specific gravity of about 1 ’ 82. It is a non-conductor of heat and electricity. It is insoluble in water, but freely soluble in carbon disulphide, phosphorus trichloride, and sulphur chloride, S 2 C1 2 . It crystallizes in forms of the regular system.
Phosphorus is extremely inflammable, taking fire in the open air at a temperature very little above its melting point; if it contains impurities, it inflames still more easily. It gradually absorbs oxygen when exposed to the air at ordinary temperatures, giving off a white vapour, which has a peculiar garlic odour; in presence of moisture phosphorous acid, H 3 PO 3 , and phosphoric acid, H 3 PO 4 , are produced.
Phosphorus, like sulphur, can exist in several allotropic modifications, the most remarkable being that produced by exposing ordinary phosphorus to light, or by heating it for some hours to a temperature of about 240° C. in an atmosphere free from oxygen. It is thus converted into a red amorphous substance, which is insoluble in carbon disulphide, and may be heated to 250° C. without alteration, but at 260° C., under the ordinary pressure, it is reconverted into ordinary phosphorus. The red modification has a much higher specific gravity (2 , 14) than ordinary phosphorus, and it is also distinguished by its inertness as compared with the latter; thus, it is not oxidized in the air at common temperatures, and it emits no odour. The transformation of ordinary phosphorus into this variety is attended with development of heat. Troost and Hautefeuille have shown that the formation of the red modification is governed by different laws according as the phosphorus is in the state of gas or liquid, and that it takes place much more rapidly in the latter case. When the tension of the vapour becomes diminished to a certain minimum value, varying for each temperature, the transformation ceases; this tension of transformation is established only after some time. Liquid phosphorus at 280° C., for instance, becomes wholly transformed into red phosphorus. The vapour given off at 260° C. is stable, but that formed at higher temperatures becomes slowly and partially converted into red phosphorus, the production of which ceases when the tension attains a given minimum; the rapidity with which this change takes place is greater the higher the temperature. At temperatures up to 520° C., however, the maximum tensions of phosphorus vapour are much higher than the tensions of transformation; thus, at 360° C. the tension of transformation is '6 atmosphere, but to prevent phosphorus boiling at this temperature, a pressure of 3'2 atmospheres must be exerted. Above 520° C. the maximum tension is not established on account of the rapidity with which the transformation takes place; and at temperatures above 550° C. the tension observed is no greater than the tension of transformation, because the transformation of the liquid phosphorus takes place more rapidly than its evaporation. The observed maximum tension and tension of transformation at different temperatures are given in the following table :— [table]
The red phosphorus produced by heating ordinary phosphorus in closed vessels differs, however, in appearance according to the temperature at which it is prepared. That obtained at 265° C. has a bright red colour and glassy fracture; that produced at 440° C. is orange-yellow in colour, and exhibits a dull fracture; at 500° C. it is denser, and has a violet grey colour; and, lastly, that prepared at 580° C. has a conchoidal fracture, exhibits signs of commencing fusion, and often encloses ruby red crystals. These modifications differ also in specific gravity, and when oxidized to phosphoric acid by iodic acid solution they develop different amounts of heat. Thus—
[table]
Hydrogen, except in the nascent state, appears to be without action on phosphorus. Oxygen forms the two oxides P 2 O 3 and P 2 O 5 . Chlorine, bromine, and iodine enter directly into reaction with ordinary phosphorus, the combination being attended with inflammation; red phosphorus also enters into reaction with these elements at ordinary temperatures, and heat is developed, although to a much less extent. When sulphur and ordinary phosphorus are melted together, combination takes place with great violence, and the experiment of bringing the two bodies together is attended with much danger. If dissolved in carbon disulphide, they are without action upon each other. Red phosphorus readily enters into reaction with sulphur when a mixture of the two substances is gently warmed, and although much heat is developed by their combination, the reaction is not explosively violent. Selenium also unites with phosphorus when the two are heated together nearly to the boiling point of the latter. Most metals combine with phosphorus when they are heated in its vapour, or when it is thrown upon them whilst they are in a state of ignition.
Compounds of Phosphorus with Hydrogen.
Three of these compounds are known, viz., phosphine or phosphoretted hydrogen gas, PH 3 , and a liquid and a solid hydrogen phosphide.
Phosphine is obtained in a pure state by decomposing phosphonium iodide by a solution of potassium or sodium hydroxide: PH 4 I + KHO = PH 3 + KI + H 2 O. It is a colourless gas, possessing a most intolerable garlic-like odour; it is sparingly soluble in water, the solution being neutral to test paper. Phosphine is extremely inflammable, igniting at a temperature a little above 100’ C.; contact with a drop of fuming nitric acid, or with chlorine or bromine, also causes it to inflame. A mixture of the gas with oxygen standing over water is gradually absorbed, and phosphorous acid produced. It is a powerful reducing agent, withdrawing oxygen with great readiness from bodies like nitric oxide, sulphur dioxide, and sulphuric acid. It precipitates many metallic salts, the precipitate in most cases consisting either of metallic phosphide, as in the case of copper, or of reduced metal, as in the case of gold and silver. It combines with hydriodic acid, forming phosphonium iodide, PH 4 I, and with hydrobromic acid forming phosphonium bromide, PH 4 Br, but not with hydrochloric acid or other acids; these salts are crystalline bodies, which are decomposed by water into phosphine and hydriodic or hydrobromic acid.
Phosphine mixed with more or less hydrogen is obtained by the action of water on the calcium phosphide, prepared by heating phosphorus with lime; by heating phosphorus with a solution of an alkaline hydroxide—
3Ca(OH) 2 + 8P + 6OH 2 = 2PH 3 + 3Ca(PH 2 O 2 ) 2 j Calcium hydroxide. calcium hypophosphite, and when phosphorous, or hypophosphorous, acid is heated—
2H 3 PO 2 = PH 3 + H 3 PO 4 .
∏ypophosphorous acid. Phosphoric acid.
4H 3 PO j = PH 3 + 3H 3 PO 4 .
Phosphorous acid. Phosphoric acid.
The gas obtained by the first and second methods is spontaneously inflammable, owing to the presence of liquid hydrogen phosphide, which may be separated by passing the gas through a U-tube cooled by a mixture of ice and salt. Liquid hydrogen phosphide probably has the composition P 2 H 4 ; by contact with various substances, and especially by hydrochloric acid, and by exposure to light, it is converted into phosphine and a yellowish solid hydrogen phosphide, supposed to have the composition P 2 H or P 4 H 2 -
5P 2 H 4 = 6PH 3 + P 4 H 2 .
Compounds of Phosphorus with the Halogens.
A gaseous phosphorus pentafluoride may be obtained by the action of phosphorus pentachloride on arsenic trifluoride: 5AsFo + 3PCL = 3PFκ + 0AsCL. It is not de- ō Ο O ' ð composed by the passage of electric sparks, even when mixed with oxygen or hydrogen. Two chlorides of phosphorus are produced by the action of chlorine on phosphorus,—a liquid trichloride, PC1 3 , and a solid pentachloride, PC1 5 . Bromine in like manner forms a liquid tribromide, PBr 3 , and a solid pentabromide, PBr 5 . Bromine also combines with phosphorus trichloride, forming the chlorobromide PCl 3 Br 2 , and this compound is capable of combining with further quantities of bromine, forming the compounds PCl 3 Br 2 . Br 2 and PCl 3 Br 2 .3Br 2 ; all these chlorobromides of phosphorus are crystalline bodies. Iodine forms the two iodides PI 3 and P 2 T 4 , which are both crystalline.
Phosphorus trichloride boils at 76° C., and the tribromide is also volatile without decomposition, but the remaining chlorine and bromine compounds of phosphorus are decomposed by heat into the trihaloid compounds and halogen: PC1 5 = PC1 3 + Cl 2 ∙ The iodides are also decomposed by heat, apparently into iodine and amorphous phosphorus.
The haloid compounds of phosphorus are all readily decomposed by water. The pentachloride and pentabromide are first converted into the oxychloride and oxybromide, two of the atoms of halogen being displaced by a single atom of oxygen—
PCl δ + H 2 O = POC1 3 + 2HC1; and by the continued action of water these compounds are converted into phosphoric acid; thus—
POBr 3 + 3H 2 O = PO 4 H 3 + 3HBr.
The behaviour of the trihaloid compounds of phosphorus is analogous to that of the oxychloride and oxybromide, phosphorous acid being produced—
PI 3 + 3H 2 O = PO 3 H 3 + 3HI.
The compounds containing more than five atoms of halogen behave as mixtures of the pentahaloid compounds with halogens—that is to say, they furnish the products of the decomposition by water of the pentahaloid compound, and also the free halogen. The iodide of phosphorus, P 2 I 4 , is peculiar in its behaviour with water; it is stated that when it is decomposed by a small quantity of water only phosphorous and hypophosphorŏus acids are formed—
P 2 I 4 + 5H 2 O = H 3 PO 3 + H 3 PO 2 + 4HI; but that when a large quantity of water is employed, a yellow insoluble substance of the composition P δ H 3 O is produced, the reaction taking place in the manner represented by the equation—
6P 2 I 4 + 24H 2 O = P δ H 3 O + PO 2 H 3 + 3PO 3 H 3 Hypophos-Phosphorone phorous acid. acid.
+ 3PO 4 H 3 + 24HL Phosphoric acid.
Oxides of Phosphorus.
The combustion of phosphorus in air produces both phosphorous and phosphoric anhydrides, P 2 O 3 and P 2 O δ ; the latter may be readily obtained by burning phosphorus in an excess of dry air, but it is almost impossible to obtain the former free from the latter. The best mode of preparing the trioxide is to bum phosphorus in a very slow current of dry air; it then condenses as a bulky white amorphous sublimate.
Phosphorous anhydride is readily soluble; it absorbs moisture with avidity, and dissolves in water, producing phosphorous acid: P 2 O 3 + 3H 2 O = 2H 3 PO 3 ∙
Phosphoric anhydride is a snow-white, flocculent, amorphous substance. It sublimes below a red heat. It has an attraction for water which probably is only exceeded by that of sulphuric anhydride, deliquescing quickly in moist air, and dissolving in water with a hissing noise and great development of heat; the solution contains metaphosphoric acid: P 2 O 5 + II 2 O = 2HPO 3 ; but this acid gradually combines with a further quantity of water, forming phosphoric acid: HPO 3 + H 2 O = H 3 PO 4 By virtue of its affinity to water, phosphoric anhydride causes the separation of the elements of water from many compounds; sulphuric acid, for example, when heated with phosphoric anhydride is converted into sulphuric anhydride—
H 2 SO 4 + P 2 O 5 = SO 3 + 2HPO 3 .
c - . . . 1 Phosphoric Sulphuric Metaphosphoric
Sulphuric acid. anh ∕ dride , an ⅛ de , p ati Z
Oxyacids of Phosphorus.
The following oxyacids of phosphorus are known :— [table]
Hgpophosphorous acid, HPH 2 O 2 ∙—To prepare this acid a solution of barium hypophosphite is treated with exactly sufficient sulphuric acid to precipitate the barium, and concentrated until its boiling point rises to 130° C.; on the liquid cooling, a mass of crystals of the acid is obtained. The solution of the barium salt is obtained by warming (not boiling) phosphorus with an excess of a solution of barium hydroxide, and afterwards removing the excess of hydroxide by passing carbon dioxide through the liquid :—
3Ba(OH) 2 + 8P + 6OH 2 = 3Ba(PH 2 O 2 ) 2 + 2PH 3 .
Hypophosphorous acid melts at 17° C.; it is perfectly stable at ordinary temperatures, but its solution is slowly oxidized to phosphorous and phosphoric acids when exposed to the air. It is a powerful reducing agent, precipitating silver and gold from solutions of their respective salts, and when heated to about 60° C. with a solution of copper sulphate it forms a precipitate of copper hydride, Cu 2 H 2 . When zinc is dissolvedin the acid, the hydrogen produced reduces a portion of the acid to phosphine, which is given off as gas.
Hypophosphorous acid is a monobasic acid. Its salts are crystallizable and soluble in water. They are permanent when dry, but their solutions are gradually oxidized on exposure to the air, especially if heated. When boiled with alkaline hydroxides they are decomposed into phosphate and hydrogen, thus—
KPH 2 O 2 + 2KHO = K 3 PO 4 + 2H 2 .
The dry salts are decomposed by heat, and in most cases furnish a residue of pyrophosphate and metaphosphate, phosphine and hydrogen being evolved, and sometimes water produced; the proportions in which these products are formed is not constant, and depends on the nature of the salt, thus— [table]
The behaviour of the nickel and cobalt, and uranyl salts is peculiar, a phosphide being also formed— [table]
Phosphorous acid, H 2 PHO 3 .—This acid is most readily prepared by decomposing phosphorus trichloride with water. By concentrating the solution by boiling until the temperature rises to 180° C., and then cooling, it is obtained in the crystalline state. The crystals melt at 70° C. Like hypophosphorous acid, it is resolved into phosphine and phosphoric acid when heated. Its solution is oxidized on exposure to the air. It is a powerful reducing agent, precipitating silver, gold, and mercury from their salts, but it does not form copper hydride with copper sulphate. Iron and zinc dissolve in a solution of phosphorous acid with evolution of phosphine, a portion of the acid being reduced by the hydrogen resulting from the conversion of another portion into the metallic salt.
Phosphorous acid is a dibasic acid, the salts produced on neutralizing its solution with alkaline hydroxide being formed by the displacement of at most two atoms of hydrogen by metals; thermochemical investigation also indicates that it is a dibasic acid. It is possible, however, to displace a third atom of hydrogen by metals, aud to produce such a compound as Na 3 PO 3 , for example, but this cannot exist in presence of water. Many of the phosphites apparently must be regarded not as derived from the acid H 2 PHO 3 , but as derivatives of a distinct-acid of the composition H 4 P 2 H 4 O 7 , or 2H 2 PHO 3 + H 2 O; barium phosphite dried at 200°-250° C., for instance, has the composition Ba 2 P 2 H 4 O 7 .
The phosphites are much more stable than the hypophosphites, but are all decomposed by heat. Salts, such as the barium salt Ba 2 P 2 H 4 O 7 , furnish hydrogen and pyrophosphate on ignition—
Ba 2 P 2 H 4 O 7 = 2H 2 + Ba 2 P 2 O 7 ; but the normal phosphites, such as zinc phosphite ZnPHO 3 , furnish hydrogen, a pyrophosphate, and a phosphide—
14Z∏PHO 3 = 7H 2 + CZn 2 P 2 O 7 + Zn 2 P 2 .
Phosphoric acid, H 3 PO 4 .—This acid may be produced by oxidizing phosphorus with nitric acid, by the oxidation of hypophosphorous and phosphorous acids, by the action of water on phosphoric anhydride and on phosphorus pentachloride, and by decomposing its salts with acids,— tricalcium phosphate or bone earth, for example, with sulphuric acid.
It may be obtained in crystals by concentrating an aqueous solution by boiling until the temperature rises to 215° C., and when the liquidis cold adding a few crystals, which cause it to solidify.
The crystals melt at about 38° C. They are readily soluble in water, and furnish a strongly acid solution, which at a boiling heat decomposes the salts of most volatile acids.
Phosphoric acid is a tribasic acid, and furnishes three classes of salts, of which the three sodium salts
NaH 2 PO 4 , Na. 2 HPO 4 , Na 3 PO 4
Sodium dihydιogen Disodium hydrogen τrisodium phosphate, phosphate. phosphate. r r may serve as examples. It has a great tendency, however, to furnish dimetallic salts, such as Na 2 HPO 4 , which are always produced when a solution of phosphoric acid is neutralized with a metallic carbonate (see p. 488).
Metaphosphoric acid, HPO 3 .—This acid is the product of the action of water on phosphoric anhydride, and is also obtained by heating phosphoric acid to redness, and by decomposing the metaphosphates with another acid.
Metaplιosphoric acid in the solid state, as obtained by evaporating its solution and heating the residue to redness, is a colourless, glassy, uncrystallizable mass, which dissolves slowly though somewhat abundantly in water, forming a strongly acid liquid; but the solution is very unstable, and is converted gradually at ordinary temperatures, and rapidly on heating, into phosphoric acid. Metaphosphoric acfd is an extremely stable body; it volatilizes completely at a bright red heat, and apparently is only in part decomposed into water and phosphoric anhydride.
The metaphosphates, or salts of metaphosphoric acid, which may be produced by heating the monometallic salts of phosphoric acid, for example—
NaH 2 PO 4 = NaPO 3 + H 2 O, Sodium phosphate. Sodium metaphosphate.
and in various other ways, are remarkable for exhibiting very different properties according to the manner in which they are prepared. Thus, ordinary sodium metaphosphate, obtained by igniting sodium phosphate and sudden cooling, forms a vitreous mass, the aqueous solution of which gives gelatinous precipitates with the salts of the heavy metals. But when a considerable quantity of ordinary fused sodium metaphosphate is allowed to cool very slowly, a crystalline mass results, and on digesting this in a slight excess of warm water the liquid separates into two layers, one containing a crystalline sodium metaphosphate, and the other the ordinary vitreous salt. By heating copper oxide and a slight excess of phosphoric acid together to 350° C., a crystalline powder is formed, insoluble in water; and by treating this copper salt with sodium sulphide another crystalline sodium metaphosphate is produced. If oxide of lead be employed, and the resulting lead salt decomposed with sodium sulphide, a sodium metaphosphate is obtained which forms with water a gummy mass, which will not pass through a filter. Lastly, a fifth variety of metaphosphates, remarkable for their insolubility in water, are formed by adding phosphoric acid in excess to solutions of sulphates or nitrates, evaporating to dryness, and heating the residue to 316° C. or upwards. They are crystalline powders.
These different metaphosphates are generally regarded as polymeric compounds, that is to say, as compounds having different molecular weights although of the same empirical composition; and the attempt has been made to infer their formulae from the relative number of atoms of the two metals contained in the mixed salts derived from them. Thus, the sodium in ordinary sodium metaphosphate may be partially displaced by another metal, and a mixed salt produced containing the two metals in the ratio of 5 atoms of the one monad metal to 1 atom of the other; hence it is concluded that ordinary sodium metaphosphate is a hexmetaphosphate, Na 6 P 6 O 18 . The mixed salts derived from the second variety of sodium metaphosphate above described contain the two metals in the ratio of 2 atoms of the one monad metal to 1 atom of another, and are therefore regarded as trimetaphosphates, the sodium salt being represented by the formula Na 3 P 3 O 9 . The mixed salts formed from the third and fourth varieties contain equal numbers of atoms of the two metals, and it is therefore supposed that they are respectively di-and tetra-metaphosphates, and that their sodium salts, for example, have the composition Na 2 P 2 O 6 and Na 4 P 4 O 12 . The insoluble crystalline metaphosphates are regarded as monometaphosphates.
1 ’ yrophosphoric acid, H 4 P 2 O 7 .—The normal salts of this acid, which is tetrabasic, may be produced by igniting the dimetallic phosphates; for example—
2Na 2 HPO 4 = Na 4 P 2 O 7 + H 2 O .
A solution of the acid may be prepared by decomposing lead pyrophosphate suspended in water by hydrogen sulphide. Apparently the acid has not yet been obtained in a pure state.
Pyrophosphoric acidis converted into metaphosphoric acid when heated to redness, and into phosphoric acid when boiled with water; the latter change, it is stated, takes place also at ordinary temperatures, but very slowly. A solution of pyrophosphoric acid does not precipitate albumen or silver nitrate, but after neutralization it gives a white precipitate with the latter. Metaphosphoric acid coagulates albumen, and gives a white precipitate with silver nitrate; and phosphoric acid does not coagulate albumen, and when neutralized gives a yellow precipitate with silver nitrate.
In addition to the normal pyrophosphates, acid salts may be obtained, formed by the displacement of only one, two, or three of the four atoms of hydrogen in the acid by metals. By fusing together sodium pyrophosphate and sodium metaphosphate the salts Na β P 4 O 13 and Νa 12 P 10 O 31 have been obtained; these salts may be regarded as formed from acids derived respectively from four and ten molecules of phosphoric acid by the abstraction of the elements of three and nine molecules of water in the same way that pyrophosphoric acid is derived from two molecules of phosphoric acid, by the abstraction of the elements of a single molecule of water— [table]
These acids are terms in a series, of which pyrophosphoric acid is the first member, formed by the withdrawal of the elements of n — 1 molecules of water from n molecules of phosphoric acid.
Very little attention has been paid as yet to the thermochemical investigation of the phosphorus compounds. The results obtained by Andrews and by Thomsen are collected in the following table; the numbers all refer to phosphorus in its ordinary state :— [table]
Sulphides and Sulpho-Acids of Phosplwrus.
The compounds P 4 S 3 , P 2 S 3 , and P 2 S 5 are readily produced by carefully heating together sulphur and red amorphous phosphorus in the required proportions. It is stated that the lower sulphides P 4 S and P 2 S may also be obtained by melting together ordinary phosphorus and sulphur under hot water in the proportions indicated by these formulae, and that both arθ liquid compounds. The sulphides P 4 S 3 , P 2 S 3 , and P 2 S 5 may be crystallized from their solutions in carbon disulphide; they all exhibit acid properties, dissolving readily in solutions of the metallic sulphides, forming salts of sulpho-acids of phosphorus, which have been little studied, however, on account of their instability. They also dissolve in solutions of the alkaline hydroxides and carbonates, forming salts of oxysulpho-acids of phosphorus; but these are also very unstable bodies, and readily decompose in contact with water. The sulphides of phosphorus are readily decomposed by water alone, their sulphur being displaced by oxygen and corresponding oxyacids of phosphorus produced, the sulphur being separated as hydrogen sulphide; thus—
P 2 S 3 + 6H 2 O = 3H 2 S + 2H 3 PO 3 .
By heating amorphous phosphorus with selenium in various proportions, the selenides P 4 Se, P 2 Se, P 2 Se 3 , and P 2 Se 5 are produced. They resemble the corresponding sulphides, and also exhibit acid characters, combining with metallic selenides to form salts of seleniophosphoric acids, which are even less stable than the corresponding sulphosalts.
Oxy-and Sulpho-Haloid Compounds of Phosphorus.
We have already pointed out that phosphorus pentachloride and bromide are converted by water into phosphorus oxychloride, POC1 3 , and phosphorus oxybromide, POBγ 3 . Phosphorus oxychloride may also be obtained by the direct combination of phosphorus trichloride with oxygen, being produced on passing oxygen through the boiling trichloride; it is a colourless mobile liquid at ordinary temperatures, but solidifies on cooling to a mass of crystals, which melt at 2° C.; it boils at 110° C. The oxybromide is a similar compound; its crystals melt at 46° C., and it boils at 195° C.
The oxides of phosphorus do not combine with the halogens, but by the action of nitric peroxide on phosphorus trichloride an oxychloride of phosphorus is formed, which bears the same relation to pyrophosphoric acid that phosphorus oxychloride or phosphoric chloride bears to phosphoric acid—
4PC1 3 + 3N 2 O 4 = 2P 2 O 3 C1 4 + 2ΝΟ + 4NOC1.
Pyrophosphoric chloride, as this compound may be termed, is a colourless liquid, which boils between 210° and 215° C., but decomposes partially into phosphoric chloride and phosphoric anhydride: 3P 2 O 3 C1 4 = 4POC1 3 + P 2 O δ . It is immediately decomposed by water, producing hydrochloric and phosphoric acids.
Phosphorus sulphochloride or sulphophosphoric chloride, PSC1 3 , and sulphophosphoric bromide, PSBr 3 , are obtained on heating phosphorus trichloride and tribromide with sulphur, and by the action of hydrogen sulphide on the pentachloride and pentabromide of phosphorus: PC1 5 + H 2 S = PSC1 3 + 2HC1. The sulphochloride is an oily liquid, having an intensely pungent odour, and boils at 126° C. The sulphobromide crystallizes; it is partially decomposed on distillation into sulphur and the compound PSBr 3 , PBr 3 . They are only slowly decomposed by water, the sulphobromide even forming a crystalline hydrate with water, PSBr 3 , H 2 O; the decomposition of the sulphochloride is represented by the following equations—
PSC1 3 +H 2 O =POC1 3 +H 2 S POC1 3 + 3H 2 O = PO 4 H 3 + 3HC1.
The sulphobromide apparently is first converted into monosulphophosphoric acid: PSBr 3 + 3H 2 O = PSO 3 H 3 + 3HBr, which is then resolved partly into sulphur and phosphorous acid, and partly, by the action of the water, into phosphoric acid and hydrogen sulphide. Phosphorus sulphochloride behaves in a similar manner when heated with a solution of sodium hydroxide; thus—
PSC1 3 + 6NaOH = FSO 3 Na 3 + 3NaCl + 3H 2 O.
A sulphopyrophosphoric bromide, P 2 S 3 Br 4 , corresponding to pyrophosphoric chloride, is formed by directly combining bromine with the sulphide P 2 S 3 . It is an oily liquid, which cannot be distilled, being resolved by heat into phosphorus pentasulphide and phosphoric bromide: 3P 2 S 3 Br 4 = P 2 S 5 + 4PSBr 3 .
Constitution of the Phosphorus Compounds.
The formulae of the two chlorides and of the tetriodide of phosphorus may be given in illustration of the constitution of the compounds of phosphorus with monad elements—
Cl Cl
I I I—P—I
Cl—P—Cl Cl—P—Cl
Z∖ i— p— i
Cl Cl
Phosphorus trichloride. Phosphorus pentachloride. Phosphorus tetriodide.
The two oxychlorides of phosphorus are represented by the following formulae, the corresponding bromine and sulphur compounds, of course, being similarly constituted:—
Cl Cl Cl
I I I
Cl—P—Cl O=P-ο—P=O
II II
ο α ci
Phosphoric chloride. Pyrophosphoric chloride.
The constitution of the oxides is expressed by the formulae—
0 0
X Z
O=P—0—P=O P—O—P
Cf ^O
Phosphorous anhydride. Phosphoric anhydride.
Similar formulae are assigned to the corresponding sulphides. Hypophosphorous, phosphorous, and phosphoric acids may be regarded as derived from the compound POH 3 , or phosphine oxide, by the displacement of one, two, and three atoms of hydrogen by the monad radicle OH; although this compound is not known at present, analogous bodies are readily obtained on oxidizing the substituted phosphines formed by displacing the hydrogen in phosphine by positive monad compound radicles, such as methyl, CH 3 —
0 0 0
II II II
H—P—H H—P—O—H H—0—P—0—H
À
H H H
Hypophosphorous Phosphorous Phosphoric acid, acid. acid. x
The relation of meta-and pyrophosphoric acid to phosphoric acid will be evident from the following formulae—
0
II
H—0—P—Ο—H
o I
Il?
0=P— O—Ξ I
H—0—P—0—H
II
0
Metaphosphoric acid. Pyrophosphoric acid.
Boron.
Symbol, B; At. wt., 11; Valency, '".
This element always occurs in the combined state as boric acid, or as a salt of boric acid. Two modifications of boron may be obtained, viz., an amorphous and a crystalline modification; the former is produced by reducing boric anhydride, B 2 O 3 , by heating it with sodium, and the latter is formed when the reduction is effected by the aid of aluminium at a very high temperature.
Amorphous boron is a dark brown powder; it does not oxidize in the air at ordinary temperatures, but when heated it burns in air or oxygen, in the latter with dazzling brightness, forming the oxide B 2 O 3 . By ignition in an atmosphere of nitrogen it is converted into a white amorphous boron nitride, BN. It does not decompose water, even at the boiling heat, but readily dissolves in nitric acid, producing .boric acid; when heated with potassium hydroxide it forms a potassium borate, hydrogen being evolved.
On heating amorphous boron with aluminium it is dissolved, and crystallizes from the fused metal on cooling; it may be separated from the latter by the aid of hydrochloric acid. The crystals are usually of a more or less brown colour, but in lustre, refracting power, and hardness they are nearly equal to the diamond. Crystalline boron is only slightly oxidized at the temperature at which diamond burns.
Boron forms a trichloride, BC1 3 , a tribromide, BBr 3 , and a trifluoride, BF 3 . The chloride may be obtained by the direct action of chlorine on amorphous boron, the combination taking place at ordinary temperatures; it is also produced by strongly heating a mixture of boric anhydride and charcoal in an atmosphere of chlorine, and by heating boric anhydride with phosphorus pentachloride to 150° C. The bromide may be obtained in a similar manner by the direct combination of its elements. Both are colourless, mobile liquids; the chloride boils at 17’ C., and the bromide at 90° C. They are readily decomposed by water; thus—
BC1 3 + 3H 2 O = H 3 BO 3 + 3HC1. Boron chloride. Boric acid.
Boron chloride is not easily deprived of its chlorine by the action of metals, which is apparently accounted for by the fact that much heat is evolved in its formation from its elements. According to Troost and Hautefeuille, the amounts of heat disengaged in the formation of the chloride and oxide of boron from their elements, and in the decomposition of the former by a large excess of water, are as follows— [table]
Boron fluoride is formed by the action of hydrofluoric acid on boric anhydride, viz., by heating a mixture of boric anhydride and calcium fluoride with concentrated sulphuric acid. It is a colourless gas, of pungent, suffocating odour; when passed into water, which dissolves about 700 times its volume of the gas, it is partly decomposed into boric and hydrofluoric acids, and partly combines with the hydrofluoric acid thus produced, forming fluoboric acid, HBF 4 . This acid can only be obtained in a state of dilute solution, but many salts formed from it are known, such as potassium borofluoride, KBF 4 , for example. Its existence appears to indicate that boron is capable of acting as a pentad element.
Boron oxide, B 2 O 3 , is most readily obtained by strongly heating boric acid; it forms a colourless, brittle, glassy mass, which dissolves readily in water, producing boric acid, and is one of the most stable oxides known. It unites with metallic oxides, when fused with them, forming borates, and at high temperatures it is capable of decomposing carbonates, sulphates, and indeed the salts of all acids the anhydrides of which are more volatile than itself.
Boric acid, H 3 BO 3 or B(OH) 3 , occurs native in many volcanic districts, especially in Tuscany, where it issues from the earth together with vapour of water; borax, N^a 2 B 4 O 7 + 1OH 2 O, and a few other of its salts are also found in nature. It crystallizes from water in white nacreous laminae, which are easily soluble; when heated to about 100° C., it furnishes a residue of the composition H 6 B 4 O 9 , which, on heating to 160° C., becomes H 2 B 4 O 7 , aud at a stronger heat it is converted into the anhydride, B 2 O 3 . A large number of borates are known, but the nature of their relation to boric acid is not well understood. A comparatively small number are derived from the acid H 3 BO 3 , the majority apparently being formed from an acid of the empirical composition HBO 2 , to which the name metaboric acid is given. Salts derived from the acids H 6 B 4 O 9 and H 2 B 4 O 7 are also known, and ordinary borax, the most important of the borates, may be regarded as formed from the latter acid.
Carbon.
Symbol, C; At. wt., 11 ’ 97; Valency, lτ .
The properties of carbon itself and its oxides have already been described (p. 86), and need not therefore be again discussed.
On comparing the compounds of carbon with those of other elements we find that, whereas there is reason to believe that the number of atoms of any of the other elements directly associated together in a molecule of their compounds is very small, probably never exceeding five, carbon compounds frequently contain a relatively very large number of carbon atoms, which from the behaviour of the compounds we are led to suppose are in direct association with each other; and while none of the remaining elements are known to furnish more than a single stable compound with hydrogen, the number of stable compounds of carbon with hydrogen which have been obtained may be counted by hundreds. We are thus led to the conclusion that carbon possesses two distinctive properties :—firstly, that of uniting with itself to an almost unlimited extent in comparison with other elements; and secondly, that of combining with hydrogen in numerous proportions. It is in consequence capable of uniting with the same elements in a multiplicity of proportions, and of furnishing a great variety of compounds, which probably exceed in number those of all the remaining, elements together.
The study of the compounds of carbon constitutes that branch of our science which is termed organic chemistry, which is treated separately below, pp. 54Asgg. ;— the study of the remaining elements and of their compounds constituting inorganic or, as it is also termed, mineral chemistry. The division is both useful and necessary, on account of the great number of carbon compounds which exist, and because, generally speaking, the compounds of other elements are distinguished from those of carbon by their comparative simplicity, and also by the comparative want of stability which all but the simplest exhibit; this we may attribute to the possession by carbon of the distinctive properties above mentioned, but otherwise there is no essential difference between the compounds of carbon and those of the remaining elements.
Organic chemistry originally dealt only with substances more or less directly derived from the animal or vegetable kingdom, and it was long believed that the chemist was powerless to produce organic substances from their elements as they were formed in the animal or plant under the influence of life, it being supposed that, therefore, the interposition of a special force, termed the vital force, was requisite. The first step towards the disproval of this hypothesis was made by Wöhler, who in 1828 succeeded in artificially producing urea, the characteristic crystalline constituent of the urine; but its final overthrow was not accomplished until 1845, when Kolbe showed that it was possible to produce acetic acid from carbon by a comparatively very simple series of reactions. Berthelot’s discovery in 1862 of the formation of acetylene, C 2 H 2 , from its elements, however, affords the simplest solution of the problem of the artificial production of organic substances from their elements, as acetylene may be without difficulty converted into ordinary alcohol, and from this body all the carbon compounds which have been artificially prepared may be more or less directly derived. Chemists have already succeeded in preparing a large number of substances which occur as products of animal or vegetable life,—for example, oxalic, tartaric, and salicylic acids; coumarin, the crystalline substance from which the Tonka bean derives its pleasant odour; and alizarin, the colouring matter derived from madder root. From the insight we have recently gained into their constitution, there is little doubt that eventually the synthesis of even the most complex organic bodies, such as albumen, will be possible.
Although carbon combines with hydrogen in a great variety of proportions, it furnishes only a very limited number of compounds with other elements; thus, only one sulphide of carbon, CS 2 , and one nitride of carbon, C 2 N 2 , are known, and no well characterized compounds of carbon with the metals have been obtained. The compounds of carbon with the halogens are more numerous. The only compound that calls for consideration here is the sulphide; the others will be noticed in the section on Organic Chemistry.
Carbon disulphide, CS 2 , the analogue of carbon dioxide, is obtained by passing sulphur vapour over charcoal heated to redness. It is a colourless, very mobile, highly refracting liquid of faint unpleasant odour, insoluble in water; it boils at 46° C. It is extremely inflammable, and its vapour when mixed with air is highly explosive; the products of its combustion are carbon dioxide and sulphur dioxide. The formation of carbon disulphide from its elements is attended with the absorption of no less than 22,000 units of heat, which explains how it is that it so readily inflames, and also the fact that it is necessary to apply heat continuously in preparing it; whereas the combustion of carbon in oxygen, when once it commences, proceeds spontaneously, owing to the large amount of heat developed in the process; we may suppose that the expenditure of energy is necessary because less heat is developed by the combination of the atoms of carbon and sulphur than is requisite to convert the carbon and sulphur from the state in which they ordinarily exist into that in which they exist in carbon disulphide, or, in other words, than is requisite to effect the separation from each other of the atoms of carbon and of sulphur in the carbon and sulphur molecules.
Carbon disulphide belongs to the class of acid sulphides or sulphur-anhydrides. It readily dissolves in solutions of alkaline hydroxides and of metallic sulphides, forming salts analogous to the metallic carbonates, and which may be regarded as derived from the carbonates by the partial or entire displacement of the oxygen in the latter by sulphur. The stability of these salts and of the corresponding acids increases with the amount of sulphur; thus, sulphocarbonic acid, H 2 CS 3 , which is formed from the salt produced by combining carbon disulphide with metallic sulphides, may be obtained as a yellow oily liquid, whereas carbonic acid, H 2 CO 3 , can only exist in extremely dilute solution.
Carbon disulphide readily dissolves sulphur and phosphorus, and also oils and fatty matters; on this account it meets with many practical applications.
By gently heating a mixture of carbon disulphide and sulphuricanhydride,carbonoxysulphideis produced,one-half the sulphur in the former being displaced by oxygen: CS 2 + SO 3 = COS + SO 2 + S; this compound may also be obtained by combining carbon monoxide with sulphur, by passing a mixture of the gas with sulphur vapour through a red hot tube. It is a colourless gas, possessing an odour like that of carbon disulphide; in properties, as in composition, it is intermediate between carbon dioxide and disulphide.
Silicon.
Symbol, Si; At. wt., 28; Valency, iv .
This element always occurs in combination either with oxygen alone as silicon dioxide or silica, or with oxygen and metals as silicates,—constituting, in fact, in these forms of combination, the greater part of the earth’s crust.
Silicon may be obtained from its chloride or fluoride by theaction of metals such as potassium, sodium, or aluminium; like carbon it exists in three distinct modifications. Amorphous silicon produced by heating potassium silicofluoride, K 2 SiF 6 , with potassium, or the corresponding sodium salt with sodium, is a dull brown powder, heavier than water; it is not affected by nitric or sulphuric acid, but is readily dissolved by hydrofluoric acid, and by a warm aqueous solution of potassium hydroxide. It fuses at a temperature below that at which steel melts; and when heated in air or oxygen it burns brilliantly and is converted into silicon dioxide. When strongly heated in a platinum crucible it becomes much denser and darker in colour, and much less oxidizable, being converted into graphitoidal silicon. On heatinga mixture of aluminium with potassium silicofluoride to the melting point of silver, a metallic button is obtained, which, when treated successively with hydrochloric and hydrofluoric acids, yields graphitoidal silicon partly in isolated hexagonal tables. This modification has the specific gravity 2 ’ 49, and may be heated to whiteness in oxygen without burning; it is not attacked by any acid excepting a mixture of nitric and sulphuric acids, and is only slowly dissolved by a solution of potassium hydroxide.
When the vapour of silicon tetrachloride mixed with hydrogen is passed over fused aluminium, the chloride is reduced, and the silicon dissolves in the aluminium; after a time a point is reached at which the silicon separates from the fused metal in large beautiful needles, having a dark iron-grey colour. These crystals constitute the adamantine variety of silicon.
Silicon appears to be capable of combining with hydrogen in the nascent state, as when a plate or wire of aluminium containing silicon is connected with the positive pole of a galvanic battery, and made to decompose a solution of sodium chloride, gas is evolved which spontaneously inflames. Silicon hydride mixed with much hydrogen is also obtained on decomposing with concentrated hydrochloric acid the magnesium silicide produced by heating a mixture of magnesium chloride, sodium silicofluoride, and sodium chloride with sodium. The pure gas is produced by the decomposition of triethylsiliconorthofoιmate in contact with sodium—
4SiH(OC 2 H 5 ) 3 = SiH 4 + 3Si(OC 2 H 5 ) 4 .
Triethylsiliconorthoformate. Silicon hydride. Tetrethylorthosili<ute.
It is a colourless gas, not spontaneously inflammable under the ordinary temperature or pressure, but only when gently heated under reduced pressure, or when mixed with hydrogen. When heated by itself, it is decomposed into amorphous silicon and free hydrogen. It is insoluble in water, and is not altered by dilute hydrochloric or sulphuric acid; potassium hydroxide, however, decomposes it readily—
SiH 4 + 2K0H + H 2 O = K 2 SiO 3 + 4H 2 .
Silicon hydride. Potassium silicate.
It precipitates silicide of copper from a solution of copper sulphate, and metallic silver mixed with a black substance, which is, perhaps, silver silicide, from a solution of silver nitrate.
By heating silicon in chlorine gas, or in an atmosphere of bromine, or iodine, it is converted into the compounds SiCl 4 , SiBr 4 , and Sil 4 ; the chloride and bromide are best prepared, however, by passing chlorine or bromine vapour over an intimate mixture of silica and charcoal heated to redness. Silicon tetrachloride and tetrabromide are colourless liquids, boiling at 50° C. and 153° C. respectively; the tetriodide crystallizes from its solution in carbon disulphide in colourless octahedra (which melt at 120 o, 5 C.), and boils at about 290° C., distilling without alteration in a stream of carbon dioxide. They are quickly decomposed by water; thus—
SiCl 4 + 2H 2 O = SiO 2 + 4HC1.
When silicon tetriodide is heated to near its boiling point with finely divided metallic silver it is converted into disilicon hexiodide, Si 2 I 6 ; this crystallizes from carbon disulphide in colourless hexagonal prisms which are decomposed on fusion into the tetriodide and a body which is perhaps the subiodide, Si 2 I 4 . By the action of bromine, this compound is converted into the corresponding bromide, Si 2 Br c , and by gently heating it with mercuric chloride the chloride Si 2 Cl 6 is obtained. The bromide forms crystals which distil without decomposition at about 240° C.; disilicon hexachloride is a colourless mobile liquid, which boils at 146°-148° C., the corresponding carbon chloride, C 2 C1 6 , being a crystalline solid which boils at 330° C.; it is very slowly decomposed when heated in closed vessels to 350° C. into the tetrachloride and silicon, and is entirely converted into these substances at 800° C. A lower liquid chloride, which probably has the composition Si 2 Cl 4 , is obtained in small quantity together with the hexachloride when the vapour of the tetrachloride is passed over fused silicon heated, to a very high temperature.
By passing the vapour of silicon tetrachloride through an exhausted porcelain tube containing fragments of felspar heated nearly to the fusing point of the latter, it is partially converted into the oxychloride (SiCl 3 ) 2 Oj this is a colourless liquid boiling at 136°-139° C., which is decomposed by water in the same manner as silicon tetrachloride. When the vapour of this compound mixed with oxygen is passed several times through a heated tube, silicon tetrachloride is reproduced, and a complex mixture of higher liquid oxychlorides is obtained containing the compounds Si 4 O 3 Cl 10 (b. p. 152°-154° C.), Si 4 O 4 Cl 8 (b. p. about 200° C.), Si 8 O 10 Clj 2 (b. p. about 300° C.), together with still higher oxychlorides.
By passing hydrochloric acid over crystallized silicon heated to barely visible redness, a mixture is produced of silicon tetrachloride with the compound SiHCl 3 , which from its relation in composition to trichloromethane or chloroform, CHC1 3 , is termed silicon-chloroform. The corresponding iodo-derivative, SiHI 3 , is formed in a similar manner by the action of hydriodic acid gas mixed with hydrogen. Silicon chloroform is a colourless extremely mobile liquid, and boils at about 36° C.; it is converted by chlorine at ordinary temperatures into the tetrachloride, and when heated with bromine it furnishes silicon trichlorobromide, SiCl 3 Br. which boils at 80° C. Silicon iodoform, SiHI 3 , is a colourless, strongly refracting liquid; it boils at about 220° C.
When a mixture of silicon tetrachloride vapour and hydrogen sulphide gas is passed through a tube heated to redness, the compound SiCl 3 .SH is obtained; it is a colourless liquid boiling at 96° C.
Hydrofluoric acid acts readily on silicon dioxide and silicious substances generally, forming silicon tetrafluoride . ’ SiO 2 + 4HF = SiF 4 + 2H 2 O. Silicon tetrafluoride is a colourless gas, which liquefies under very strong pressure, and solidifies at-140° C.; it is not inflammable. It is readily absorbed by water, which decomposes it into silica and hydrofluosilicic acid: 3SiF 4 + 2H 2 O = 2H 2 SiF 6 + SiO 2 . This acid is not known except in solution, but its salts, which are readily produced by neutralizing its solution with metallic oxides, hydroxides, or carbonates, are stable compounds, and are mostly soluble in water. By prolonged ignition they are decomposed into silicon fluoride which escapes, and a metallic fluoride which remains behind. Potassium fluosilicate, K 2 SiF 6 , is one of the least soluble of the potassium salts, and consequently this acid is often used to precipitate potassium from solutions of its salts.
By passing the tetrafluoride over silicon heated nearly to the softening point of porcelain, it is converted into a subfluoride, which probably has the composition Si 2 F β ; it is a white, very light powder.
Only one oxide of silicon is known, viz., the dioxide or silica, SiO 2 . According to Troost and Hautefeuille, its formation from amorphous silicon and oxygen is attended with the development of a very large amount of heat, viz.,—
(Si, O 2 ) = 219,240 units of heat.
According to the same observers, in the formation of the tetrachloride from amorphous silicon and chlorine, and in its decomposition by 140 times its weight of water, the following amounts of heat are developed :— [table]
and from the amount of heat developed on dissolving amorphous and crystallized (? adamantine) silicon in a mixture of nitric and hydrofluoric acids they calculate that the transformation of the former into the latter is attended with the development of 8120 units of heat.
When obtained by igniting the gelatinous silica from the decomposition of the haloid silicon compounds by water, or of the soluble silicates by acids, silica is a snow-white amorphous powder; but it occurs native as quartz in large transparent crystals of the hexagonal system, of the specific gravity 2 , 5 to 2 ’ 8, and of extreme hardness. Native silica of all kinds is insoluble in water and in all acids except hydrofluoric acid; this is also true of artificial ignited silica. The gelatinous hydrate which separates on the addition of acids to solutions of alkaline silicates is also almost insoluble in water and in acids, but that which results from the decomposition of silicon fluoride is stated to be soluble in a considerable quantity of water. Gelatinous silica dissolves without difficulty in solutions of potassium or sodium hydroxide at the ordinary temperature, and all forms of silica are more or less readily dissolved by solutions of these alkalies, especially when heated with them under pressure; in this manner, according to the proportions of silica and alkali employed, various alkaline silicates are formed which are the more soluble the greater the proportion of alkali metal they contain. Those which contain a larger proportion than is indicated by the formula Na 2 SiO 3 , but still a sufficient quantity to make them completely soluble in water, are termed “water-glasses;” water-glass is now largely manufactured and employed for hardening and preserving stone, in the preparation of paints, especially for mural painting, and in the manufacture of soap.
Although silica is an extremely weak acid oxide, on account of its non-volatility, it is capable of decomposing the salts of all volatile acids at more or less elevated temperatures, and the salts formed from it are stable at the highest temperatures. A great variety of artificial and natural silicates are known,—some of which may be regarded as derived from silicic acid, H 2 SiO 3 , and others from an acid which may be termed orthosilicic acid, II 4 SiO 4 , but the majority are far more complex, those which occur as natural minerals being formed by the association of silica with basic oxides in proportions included between the following limits :— [table]
On adding acids to a solution of an alkaline silicate, a gelatinous “hydrate of silica” is precipitated; the whole of the silica may be retained in solution, however, if a dilute solution of an alkaline silicate be poured into a considerable excess of hydrochloric acid, but may be precipitated by gradually neutralizing the acid. The retention of the silica in solution is, perhaps, due to the formation of a soluble chlorhydrin (p. 565). If a stratum τ 4 σ of an inch in depth of the hydrochloric acid solution be placed in a “dialyzer” formed by stretching wetted parchment paper across a light hoop of wood or gutta-percha, and the dialyzer be floated in a vessel of water, the water in the outer vessel being changed at intervals, after four or five days the hydrochloric acid and the soluble chlorides are found to have diffused completely into the water; but the solution in the dialyzer still contains the silica, and a 5 per cent, solution prepared in this way may be concentrated by boiling down in an open flask until it contains nearly 14 per cent, of silica. The solution is tasteless, limpid, and colourless; it has an acid reaction rather greater than that of carbonic acid; for every 100 parts of silica present, however, only 2 Ι 3 parts of potassium hydroxide are requisite to neutralize the acid reaction. The solution is not easily preserved for many days, as it becomes converted into a solid transparent jelly; coagulation is retarded by hydrochloric acid, and by small quantities of potassium or sodium hydroxide, but is effected in a few minutes by the addition of 10 1 00 part of any alkaline carbonate, although ammonia and its salts are without effect. Sulphuric, nitric, and acetic acids are also without action, but a few bubbles of carbon dioxide gas slowly cause coagulation. The condition in which the silica is present in the solution having these very remarkable properties is not known, but on the assumption that a chlorhydrin is produced on decomposing the alkaline silicate by an excess of hydrochloric acid, it appears not improbable that, as the excess of the latter is removed by dialysis, the chlorhydrin is gradually decomposed by the water and, perhaps, converted into the hydroxide H 4 SiO 4 , or some other soluble hydroxide, and that the subsequent coagulation may be due to the conversion of this hydroxide into an insoluble hydroxide of different molecular composition. The solution evaporated at 15° C. in a vacuum leaves the silica in the form of a transparent, glassy, very lustrous hydrate, containing after two days’ exposure over sulphuric acid an amount of water which agrees very nearly with the formula H 2 SiO 3 .
When silicon chloroform is added to water at the ordinary temperature, hydrogen is evolved, and hvdrated silica is produced: SiHCl 3 + 2H 2 O = H 2 + SiO 2 + 3HCl; but if water at zero is employed, a precipitate is obtained which, after drying over sulphuric acid, and then at 150° C., has the composition (HSiO) 2 O. This compound is analogous in composition to the unknown anhydride of formic acid: (HCO) 2 O = 2HCO(OH) - OH 2 ; its formation is probably preceded by that of the hydroxide SiH(OH) 3 . It is a snow-white, voluminous, amorphous substance, slightly soluble in water; it is decomposed by the least heat when in a moist state, with evolution of hydrogen, and is dissolved by ammonia, caustic alkalies, and alkaline carbonates, with evolution of hydrogen and production of an alkaline silicate. It is not acted upon by any acids except hydrofluoric acid, even concentrated nitric acid being without action. It decomposes when heated above 300° C.
Silicon hexachloride and hexiodide are decomposed by water in a similar manner, being first converted into the hydroxide Si 2 (OH) 6 , which then furnishes siliconoxalic acid Si 2 O 2 (OH) 2 , This is a white substance, and when heated is decomposed, like the preceding compound, with evolution of hydrogen, leaving a residue of silica; it is also decomposed with evolution of hydrogen by even the weakest bases. It rapidly reduces potassium permanganate in the cold, but is only slowly oxidized by chromic acid, and is without action on solutions of gold chloride or selenious acid. On decomposing the chloride, Si 2 Cl 4 , however, by water an hydrated oxide is obtained which not only reduces potassium permanganate and chromic acid, but also gold chloride and selenious acid. The behaviour of siliconoxalic acid on oxidation is precisely similar to that of its carbon analogue oxalic acid, C 2 O 2 (OH) 2 , which, however, is a crystalline, strongly acid body, which dissolves readily in water and furnishes stable salts.
By the action of anhydrous alcohol, silicon tetrachloride is converted into tetrethylorthosilicate, Si(OC 2 H 5 ) 4 , which is a colourless liquid boiling at 166° C.; it is gradually decomposed by water, being converted into alcohol and gelatinous silica. By the joint action of zinc ethyl and sodium on this compound, it is possible to displace in succession each of the groups (OC 9 H.) by ethyl, thus:— [table]
In this manner the following compounds have been prepared :— [table]
These compounds are all colourless liquids, insoluble in water, and stable in the air; the first only is decomposed when allowed to remain in contact with water. Silicon ethyl, Si(C 2 H 5 ) 4 , may also be obtained by the action of zinc ethyl on silicon tetrachloride; and from the latter compound and zinc methyl silicon methyl, Si(CH 3 ) 4 , has been prepared; similarly, the hexiodide is converted into the corresponding ethyl derivative Si 2 (C 2 H 5 ) 6 by the action of zinc ethyl. These silicon hydrocarbons exhibit the closest resemblance to their analogues the paraffins. Thus by the action of bromine on silicon triethylhydride, Si(C 2 H δ ) 3 H, the bromide Si(C 2 H 5 ) 3 Br is produced, which may be converted by the action of aqueous ammonia into the silicon alcohol triethylsilicol, Si(C 2 H 5 ) 3 .OH; this method of formation is precisely analogous to that by which the hydrocarbon C(CH 3 ) 3 H is converted into the alcohol C(CH 3 ) 3 .OH. Triethylsilicol also resembles the corresponding alcohol triethylcarbinol, C(C 2 H 5 ) 3 .OH, in properties; thus, both are colourless, viscid liquids, having a strong camphor-like odour, difficultly soluble in water; the former boils at 154° C., the latter at 141° C.; they dissolve sodium with evolution of hydrogen, being converted into corresponding sodium derivatives Si(C 2 H 5 ) 3 .ONa and C(C 2 H 5 ) 3 .ONa; and both are acted upon by acetic chloride in a similar manner— [table]
The most important difference is in their behaviour on oxidation,—triethylsilicol remaining unattacked, while triethylcarbinol is readily converted into acids containing fewer atoms of carbon.
Silicon ethyl is converted by the action of chlorine into a chlorinated derivative Si(C 2 H 5 ) 3 (C 2 H 4 Cl), from which the silicon alcohol SiC 8 H 19 . OH may be prepared, corresponding to the yet unknown carbon alcohol C 9 H 1 q.OH. It is a liquid insoluble in water, smelling like camphor, and boiling at 190° C
By heating the compound Si(C 2 H 5 )(OC 2 H 5 ) 3 with hydriodic acid solution it is converted into siliconpropionic acid—∙
C 2 H 5 .Si(OC 2 H 5 ) 3 + 3HI = C 2 H 5 .SiO(OH) + 3C 2 H 5 I + H 2 O.
fr " ,lh p⅛≡' th - Oo.Bpwpi.nidd.
In a similar manner, siliconacetic acid, CH 3 .SiO(OH)j may be prepared from the corresponding methyl compound. But although these silicon acids correspond in composition to acetic acid, CH 3 .CO(OH), and propionic acid, C 2 H 5 .CO(OH), they exhibit very different properties; thus, they are white amorphous substances, insoluble in water, although soluble in alkaline solutions, from which they are precipitated by the addition of acids, whereas acetic and propionic acids are colourless liquids, soluble in water, and boil respectively at 119° C. and 140° C.
When the vapour of carbon disulphide is passed over a heated mixture of silica and carbon, silicon disulphide, SiS 2 , is produced; it crystallizes in white silky needles, which quickly decompose in moist air into hydrogen sulphide and amorphous silica.
From the foregoing description of the silicon compounds, it will be evident that while closely allied both in composition and in many of their properties to the carbon compounds, they nevertheless differ from them in numerous important particulars. Thus, carbon dioxide is gaseous, and silicon dioxide is a non-volatile solid; the chlorides of carbon are stable in presence of water except perhaps at relatively very high temperatures, but the chlorides of silicon are with the greatest readiness decomposed by water; carbon disulphide is a volatile liquid not affected by water, while silicon disulphide is a solid which cannot exist in presence of water; and obviously the representatives of the carbon compounds oxalic acid, acetic acid, and propionic acid in the silicon series possess very different properties. In many respects silicon bears considerable resemblance to boron, the resemblance being especially noticeable between the elements themselves, and in the behaviour of their haloid compounds with water, and also in the property which the fluorides of both elements possess of combining with hydrogen fluoride. It is of interest to note that much more heat is developed in the formation of the oxides of boron and silicon than in the formation of carbon dioxide, which alone is gaseous; thus— [table]
In the case of the corresponding chlorides the order of volatility is reversed: thus— [table]
In discussing the remaining elements it will suffice to indicate the general nature of their relations to each other, as a full description of the more important will be given under other headings. It will be convenient in the first instance to consider those elements together which are most closely related in properties, and afterwards to indicate the manner in which the elements generally are related to each other.
Metals of the Alkalies. [table]
The elements of this class are white metals, volatile at high temperatures; lithium is softer than lead but harder than sodium, while sodium is harder than potassium, and potassium harder than rubidium, the last mentioned being as soft as wax. They may all be separated from their chlorides by electrolysis, and apparently also by strongly heating mixtures of their carbonates with charcoal in iron retorts; the latter method is employed in the manufacture of sodium and potassium, and rubidium has been prepared by it. Caesium has not yet been obtained in a pure state, but an amalgam of caesium may be procured by submitting its chloride to electrolysis, employing a globule of mercury as the negative electrode.
Caesium is the most electro-positive element yet discovered; the remaining members of the group follow it in this respect in the order of their atomic weights. They are easily fusible (see table above), and their compounds with other elements are all fusible.
The metals of this group and their compounds furnish characteristic spectra, which are distinguished from those of most other elements by their simplicity. Lithium and its salts communicate a beautiful red colour to flame, sodium salts an intense yellow, and potassium, caesium, and rubidium salts a violet colour.
According to Troost and Hautefeuille, when potassium is heated to 350°-400° C. in an atmosphere of hydrogen, it is converted into a hydride of the composition K 2 H; and the corresponding hydride, N^a 2 H, may be prepared in a similar manner from sodium. Lithium, however, manifests but little tendency to combine with hydrogen, absorbing only 17 times its volume of the gas at 500° C. The hydrides are white bodies resembling silver in appearance; potassium hydride is very brittle, but sodium hydride is as soft as sodium, although it becomes brittle when heated; the former takes fire spontaneously in air, but the latter is much more stable. The compounds of lithium, sodium, and potassium with hydrocarbon radicles are only known in combination with the zinc compounds; thus, the body obtained by the action of sodium on zinc ethyl has the composition ZnNa(C 2 H 5 ) 3 .
The haloid compounds of the elements of this group may be formed by the direct combination of the metals with halogens; their affinity for halogens, however, appears to be inversely proportional to their atomic weights. Thus, it is requisite to heat sodium to a moderately high temperature in an atmosphere of chlorine in order to secure its conversion into the chloride NaCl; but potassium inflames in chlorine at the ordinary temperature. Similarly, sodium may be preserved unchanged in contact with bromine, and is scarcely affected even when heated with it to 200° C.; potassium, however, causes a violent explosion when thrown on bromine. Sodium may also be fused with iodine without appreciable reaction occurring, but potassium at once combines with it with explosive violence.
The metals of this group all decompose water in the cold with evolution of hydrogen and production of a solution of the hydroxide: for example—
2Na + 2OH 2 = H 2 + 2Na,OH.
Lithium does not fuse when thrown on water, and sodium fuses but does not inflame; potassium and rubidium, however, not only fuse but also take fire and burn with a violet coloured flame. According to Thomsen’s determinations, more heat is developed in the decomposition of water by lithium or sodium than by potassium; thus— [table]
The difference in their behaviour when thrown on water is, perhaps, due to the circumstance that the action proceeds more rapidly with potassium than with sodium.
All the metals of this group attract oxygen with avidity on exposure to air, but lithium is much less oxidizable than sodium or potassium, rubidium takes fire spontaneously in air, and caesium amalgam is much more readily oxidized than rubidium amalgam; so that their affinity for oxygen, as for the halogens, is proportional to their atomic weights. Only one oxide of lithium, Li 2 O, has been obtained, but two oxides of sodium, Na 2 O, and Na 2 O 2 , and three oxides of potassium, K 2 O, K 2 O 2 , and K 2 O 4 , are known; the oxides of the remaining elements have not been investigated. The oxides of the form M' 2 O are white deliquescent solids, which readily dissolve in water with the development cf much heat, forming solutions of the corresponding hydroxides, Μ'.OH; they are the most powerfully basic oxides known. The oxides of the form M 2 O 2 dissolve in water, butcorrespondingsalts are unknown; potassium tetroxide dissolves in water with evolution of oxygen. The hydroxides M'.OH, or alkalies, are white fusible solids of remarkable stability, volatilizing at high temperatures apparently without undergoing decomposition; they are very soluble in water, furnishing strongly alkaline solutions, soapy to the touch.
The metals of the alkalies appear to act uniformly as monads, and the salts derived from them are all formed on the types M'C1, M 2 'SO 4 , <fcc. Their salts are colourless excepting those derived from coloured acids; and the corresponding salts of the several metals are isomorphous. Their chlorides, carbonates, sulphates, and phosphates are soluble in water, and the chlorides and sulphates furnish neutral solutions. Their sulphates form characteristic double salts called “alums” with sulphate of aluminium and the sulphates of allied elements. The alums all crystallize with the same number of molecules of water, and correspond in composition to ordinary alum, K 2 A1 2 (SO 4 ) 4 , 24H 2 O; they are also isomorphous.
Although the general resemblance between the metals of this group is very great, it will be obvious that they differ in many important respects. The properties of potassium, rubidium, and caesium, which are very closely related, are especially different from those of lithium and sodium, which are also closely related. For instance, the normal lithium and sodium salts of chloroρlatinic acid, H 2 PtCl 6 , the acid lithium and sodium salts of tartaric acid, H 2 .C 4 H 4 O 6 , and the lithium and sodium alums are readily soluble in water, whereas the corresponding salts of potassium, rubidium, and caesium are difficultly soluble. Lithium, however, in many of its properties is more closely allied to magnesium and calcium than to sodium; and it may be noticed that its chloride, like that of magnesium, dissolves in water with considerable development of heat.
The reactions involved in the formation of a few compounds of lithium, sodium, and potassium have been investigated by Thomsen, with the following results :— [table]
Metals of the Alkaline Earths. [table]
The members of this group are strongly positive elements, which act uniformly as dyads; the relations between them are similar to those which obtain in the preceding group, the element with the highest atomic weight being the most positive. Strontium, however, in some respects appears to be more closely related to barium than to calcium. They are best prepared by the voltaic decomposition of their fused chlorides; their chlorides are not reduced by sodium, but calcium may be obtained by heating calcium iodide with sodium.
Calcium is a yellow metal of the colour of gold largely alloyed with silver; in hardness it is intermediate between lead and gold, and it is very ductile. It melts at a red heat. Strontium closely resembles calcium, but has a deeper colour; little is known of barium, but it appears to resemble strontium and calcium. The behaviour of these elements with halogens is apparently similar to that of the metals of the preceding group with low atomic weights.
They decompose cold water with evolution of hydrogen, but less readily than the alkali metals; it is remarkable that concentrated nitric acid is almost without action on calcium and strontium even when heated to boiling, although the diluted acid dissolves them rapidly.
They are less easily oxidized than the alkali metals, calcium being the least, and barium apparently the most oxidizable; in dry air calcium and strontium remain untarnished for a considerable period. They burn with great brilliancy when ignited in air.
Each metal furnishes two oxides,—a monoxide such as barium oxide, BaO, and a dioxide such as barium peroxide, BaO 2 ; these oxides are white, earthy, infusible substances; the dioxides are decomposed into the monoxides and oxygen at a red heat. Their monoxides are powerfully basic, but apparently the dioxides are almost destitute of basic properties; the action of acids on the two classes of oxides is illustrated by the following equations :— [table]
Barium monoxide is converted into the dioxide when heated to dull redness in an atmosphere of oxygen, but the dioxides of strontium and calcium cannot be prepared in this manner. By the action of water the monoxides are converted with development of much heat into the correspondinghydroxides, which are strongly alkaline. Barium hydroxide, Ba(OH) 2 , is dissolved by water, although to a less extent than tho hydroxides of the alkali metals; it separates from water in crystals of the composition Ba(OH) 2 , 8H 2 O; it is not decomposed even by prolonged ignition. Strontium hydroxide, which also crystallizes with 8 molecules of water, is much less soluble than the barium compound; it is not decomposed by ignition. ¼ Calcium hydroxide, Ca(OH) 2 , however, requires about 700 parts of cold water to dissolve it, the crystals of strontium hydroxide requiring only 50 parts, and is decomposed into calcium oxide and water by prolonged ignition. When barium hydroxide solution is added to hydrogen dioxide a hydrated barium peroxide, BaO 2 , 6H 2 O, is precipitated in crystalline scales; the corresponding strontium and calcium compounds, SrO 2 ,8H 2 O and CaO 2 , 8H 2 O, may be prepared in a similar manner, and are converted into the anhydrous oxides by heating to 100° C.
Barium and strontium carbonates are practically insoluble in water, and are not decomposed on ignition; calcium carbonate, CaCO 3 , is only very slightly soluble in water, and is slowly converted into calcium oxide and carbon dioxide on ignition. In this respect lithium carbonate, Li 2 CO 3 , closely resembles calcium carbonate, as it is only sparingly soluble in water, and decomposes on ignition; whereas the carbonates of the other alkali metals are very soluble, and are not decomposed by heat.
Calcium sulphate, CaSO 4 , crystallizes with two molecules of water, and is slightly soluble in water; but strontium and barium sulphates form anhydrous crystals, and are practically insoluble in water. Similarly, calcium nitrate crystallizes in deliquescent, very soluble monoclinic prisms of the composition Ca(NO 3 ) 2 , 4H 2 O, and an apparently isomorphous salt of similar composition, but efflorescent, may be obtained by crystallizing strontium nitrate at a low temperature; but from a hot solution strontium nitrate separates in anhydrous octahedra isomorphous with barium nitrate, which crystallizes without water. The chlorides of calcium and strontium are deliquescent crystalline substances, very soluble in water, and also soluble in alcohol; barium chloride is very much less soluble in water, and is insoluble in alcohol, and does not deliquesce. The chlorides of barium, strontium, and calcium furnish perfectly neutral solutions. Barium salts communicate an apple-green colour to flame, strontium salts a brilliant crimson, and calcium salts an orange-red; the spectra of the three elements are comparatively simple.
Glucinum — Magnesium — Zino — Cadmium — Mercury. [table]
Gludnum, or beryllium, obtained by the action of sodium on its chloride, is a white, malleable metal fusible below the melting-point of silver; it does not burn in air or oxygen, but becomes coated with a layer of oxide which seems to protect it from further change. When heated in an atmosphere of chlorine it is converted into the chloride GC1 2 , It does not decompose water even when heated to whiteness. It is easily dissolved by dilute hydrochloric and sulphuric acids; but nitric acid, even when concentrated, has little action on it. It is readily dissolved by a solution of potassium hydroxide.
Glucinum hydroxide, G(OH) 2 , separates as a white bulky precipitate on the addition of an alkaline hydroxide to a solution of a glucinum salt; it exhibits pronounced acid properties, dissolving in solutions of the alkalies and even in those of alkaline carbonates. Glucinum oxide, GO, is a light, very bulky, white powder; when strongly heated it volatilizes without fusing. Glucinum chloride i GC1 2 , crystallizes in silky needles; it is very deliquescent, and dissolves in water with a hissing noise and considerable rise of temperature. Glucinum hydroxide readily dissolves in acids, but the ignited oxide slowly; the glucinum salts are mostly soluble, and have a sweet, slightly astringent taste. Basic glucinum salts are obtained with great facility; thus by digestion with the hydroxide, glucinum nitrate, G(NO 3 ) 2 , is converted into the basic nitrate, G 3 N 2 O 8 , and the basic sulphate, G 3 SO 6 , may in a similar manner be obtained from the normal sulphate, GSO 4 . Glucinum sulphate, GSO 4 , 4H 2 O, is extremely soluble in water; it forms a double salt with potassium sulphate of the composition GK 2 (SO 4 ) 2 ,2H 2 O.
Magnesium, Zinc, Cadmium. — Magnesium is a white, malleable ductile metal about as hard as calcspar; it melts and volatilizes nearly as readily as zinc. It readily decomposes water below its boiling-point, but not in the cold, unless associated with a less positive metal such as copper. It is easily dissolved by dilute acids.
Zinc is a hard, white metal with a shade of blue; when pure it is very malleable, but it is usually brittle owing to the presence of impurities such as lead and iron. Commercial zinc, however, becomes malleable at 100°-150° C., and may then be rolled; at 210° C. it again becomes brittle, and may easily be reduced to powder at that temperature. Zinc decomposes vapour of water at a temperature below dull redness, and readily, even at ordinary temperatures, when it is associated with coper or other less positive metals. The pure metal dissolves very slowly in dilute acids, but ordinary zinc dissolves quickly, because the lead contained in it acts as an electro-negative element. Zinc dissolves in aqueous alkalies with evolution of hydrogen.
Cadmium is a bluish-white metal like zinc; it is soft, though harder and more tenacious than tin, and is very malleable and ductile. When heated to about 80° C. it becomes very brittle, and may be powdered with facility. It appears to decompose water at a high temperature, and dissolves when heated with slightly diluted hydrochloric or sulphuric acid, but its best solvent is dilute nitric acid.
Magnesium is usually prepared by heating its chloride with sodium; zinc and cadmium are readily obtained by reducing their oxides with charcoal. Magnesium, zinc, and cadmium are capable of being polished, and then assume a bright metallic lustre; they retain their lustre in dry air, but when exposed to moist air a film of oxide forms on the surface which protects them from further oxidation. Magnesium is the most positive metal of the three, as it precipitates zinc and cadmium from their salts, and zinc is more positive than cadmium. When heated in chlorine, or in the vapour of bromine or of iodine, they burn brilliantly, forming corresponding chlorides, bromides, or iodides, such as MgCl 2 , ZnBr 2 , Cdl 2 ; if moistened with water, zinc and cadmium readily combine with the halogens at the ordinary temperature.
They are easily combustible in air or oxygen; the light emitted by burning magnesium is intensely brilliant and very rich in actinic rays; zinc also burns with a brilliant flame. They are thus converted into the oxides MgO, ZnO, CdO; magnesium and zinc oxides are white, but the latter becomes lemon yellow when heated; cadmium oxide is yellowish brown. These oxides are unalterable in the fire; they dissolve readily in acids, forming corresponding salts, e.g. : MgO + H 2 SO 4 = MgSO 4 + H 2 O. Magnesium oxide, when mixed with water, gradually combines with it, and forms the hydroxide Mg(OH) 2 , but no sensible development of heat occurs during the hydration. Cadmium oxide also unites with water.
The hydroxides of magnesium, zinc, and cadmium separate as white amorphous precipitates, insoluble in water, on the addition of alkalies to solutions of the salts of these metals. Magnesium hydroxide alone exhibits a faint alkaline reaction; zinc hydroxide readily dissolves in alkalies; they are all dissolved by solutions of ammonium salts; they are easily resolved by heat into water and the oxide.
Cadmium also furnishes a suboxide, Cd 2 O, which is a green powder; on treatment with acids it is resolved into metallic cadmium and the oxide CdO; in other words, a salt of the oxide CdO is produced, and cadmium remains.
The chlorides of these metals, like glucinum chloride, are volatile deliquescent solids; magnesium and zinc chlorides are extremely soluble in water, and their solutions are strongly acid, but cadmium chloride is less soluble; their dissolution in water is attended with development of heat, thus :— [table]
Their bromides and iodides are also soluble in water. The chlorides of magnesium, zinc, and cadmium form crystalline compounds with the chlorides of many other metals, the tendency to combine with other chlorides being especially marked in the case of cadmium chloride.
Their sulphates are soluble in water, especially those of magnesium and zinc, which are isomorphous; the sulphates of the three metals form isomorphous double salts with potassium sulphate. The composition of their sulphates and double sulphates with potassium is as follows :— [table]
Zinc, like glucinum, manifests a great tendency to form basic sulphates.
The carbonates of magnesium, zinc, and cadmium are white and insoluble in water ; they are without difficulty decomposed by heat into the oxide and carbon dioxide.
Glucinum sulphide is soluble in water without decomposition, but is decomposed by dilute acids; magnesium sulphide is converted by water into magnesium hydroxide with evolution of hydrogen sulphide: MgS + 2H 2 O = Mg(OH) 2 + H 2 S; zinc sulphide is insoluble in water, but is dissolved and decomposed by dilute mineral acids; lastly, cadmium sulphide is insoluble in water and dissolves with difficulty even in boiling dilute hydrochloric acid, though it is readily decomposed by concentrated hydrochloric acid even at ordinary temperatures. Cadmium sulphide is a brilliant yellow substance, the remaining sulphides are white.
Mercury. — This element is the only metal that is liquid at common temperatures, with the exception, possibly, of caesium, and the recently discovered element gallium. The latter element, indeed, appears to be closely allied to zinc in properties, but is less positive, and is probably a member of the aluminium group; its specific gravity is 5 ’ 95.
Mercury occurs native chiefly in the form of sulphide, from which it is separated either by distillation with slaked lime (calcium hydroxide), or by burning off the sulphur. It is a highly lustrous metal, and in the solid state is malleable. It slowly combines with the halogens at ordinary temperatures, and also with sulphur if triturated with it; it does not tarnish, even in moist air, but slowly absorbs oxygen when heated to about 400° C., being converted into the red oxide HgO; this oxide is decomposed on ignition. Hydrochloric acid, whether cold or hot, is without action on mercury, but it is slowly dissolved by hydriodic acid; concentrated nitric acid dissolves it readily; it is also dissolved by hot concentrated sulphuric acid. It has no action upon water at any temperature. It is capable of uniting with most metals, forming liquid or solid compounds called amalgams; the solid amalgams appear for the most part to be of definite composition, but the two metals are held together by very feeble affinities.
Mercury unites with the halogens, with oxygen, and with sulphur in two proportions, forming the two series distinguished as mercurous and mercuric compounds; thus :— [table]
The compounds of mercury with halogens, and mercury compounds generally, are difficultly soluble, or insoluble, in water. But the compounds with halogens dissolve readily in solutions of the haloid compounds of the alkali metals in consequence of the formation of double salts; the haloid compounds of mercury, in fact, exhibit a more marked tendency than those of any other metal of the group to form double salts with the haloid compounds of other metals, cadmium being most nearly allied to mercury in this respect ; they unite also with the haloid acids, forming the compounds HHgCl 3 , HHgBr 3 , and HHgI 3 .
Mercurous chloride or calomel, Hg 2 Cl 2 , is a white crystalline solid, insoluble in water; it blackens slowly on exposure to light, and is decomposed by heat into metallic mercury and mercuric chloride, HgCl 2 . Mercuric chloride or corrosive sublimate is also white and crystalline, but soluble in water; it boils without decomposition at 295° C.; the density of its vapour corresponds with the formula HgCl 2 . By the action of ammonia on the chlorides of mer cury, one-half the chlorine is removed and displaced by the monad radicle amidogen: Hg 2 Cl 2 + 2NH 3 = Hg 2 Cl(NH 2 ) + NH 4 C1; the lower chloride is thus converted itìto black mercurous amidochloride, the higher chloride furnishing mercuric amidochloride or white precipitate, HgCl(NH 2 ).
Mercury hydroxides are unknown, the chlorides of mercury being at once converted into corresponding oxides by the action of alkalies; but when mercury is associated with positive hydrocarbon radicles, it forms powerfully basic hydroxides such as Hg(C 2 H 5 )OH, which is a colourless oil. Mercurous oxide, Hg 2 O, is a black substance which by mere exposure to light, or a very gentle heat, is converted into the metal and the red oxide. Mercuric oxide, HgO, when prepared by heating mercury in oxgyen, forms red crystalline scales, but it is precipitated as a yellow powder on the addition of an alkali to a solution of mercuric chloride. These two forms of the oxide differ even in their chemical behaviour, and are perhaps polymeric; thus, when the yellow oxide is boiled with a solution of potassium dichromate a basic mercuric chromate, Hg 3 CrO 6 , is formed, but the crystallized oxide forms a more basic salt, viz., Hg 4 CrO 7 , under similar circumstances. Mercuric oxide is dissolved by acids, forming mercuric salts such as mercuric nitrate, Hg(NO 3 ) 2 ; by digesting solutions of these salts with mercury they are converted into corresponding mercurous salts such as mercurous nitrate Hg 2 (NO 3 ) 2 , But mercuric oxide also exhibits feeble acid properties; thus it forms the compound K 2 HgO 3 when dissolved in fused potassium hydroxide. Basic mercuric salts are also very readily produced; mercuric sulphate, HgSO 4 , for example, is decomposed by water into a soluble acid salt and the insoluble basic sulphate Hg 3 SO d . Mercuric sulphide, HgS, is thrown down as a black precipitate on passing hydrogen sulphide through solutions of mercuric salts; it is not decomposed by hydrochloric acid; when heated it becomes red. Native cinnabar and vermilion consist of the red modification of mercuric oxide.
No tendency to combine with hydrogen has been observed in the case of the metals of this group, but they have all been obtained in combination with hydrocarbon radicles. Their organo-metallic derivatives correspond in composition to zinc ethyl, Zn(C 2 H 5 ) 2 , the attempt to convert mercurous chloride into a corresponding organo-metallic compound by the action of zinc ethyl has been unsuccessful, mercury ethyl, Hg(C 2 H 5 ) 2 , being formed instead and mercury separated; thus :—
Hg 2 Cl 2 + Zn(C 2 H δ ) 2 = Hg(C 2 H 5 ) 2 + Hg + ZnCl 2 .
It would appear, in fact, that only mercurous compounds can exist in which mercury is associated with negative radicles.
It has been shown that the density of the vapour of each of the metals cadmium and mercury is the half of its atomic weight, in other words, the gaseous molecules of these elements are monatomic (p. 471 ). As we are not acquainted with the density of any other metal in the state of gas, it is impossible to say whether these metals are peculiar in this respect.
On comparing the properties thus briefly sketched of the elements of this group, the general resemblance will be at once noticed, but it will be remarked also, that, although in respect of many properties the variations are gradational in the same degree as the atomic weights, certain properties are not continuous, but recur periodically. Thus glucinum, zinc, and mercury, the first, third, and fifth terms of the group, have many characteristics in common which are apparently wanting in the case of magnesium and cadmium, the second and fourth terms. This is evidenced especially by certain of their physical properties, by the solubility of their hydroxides in alkalies, by the formation of basic salts, and by the behaviour of their sulphides.
The elements of this group are in many respects closely related to the metals of the previous group, of which they may be regarded as forming a sub-group, the formulae of their compounds being similar in most cases; the general behaviour especially of magnesium and its compounds, it will be evident, is in close accordance with that of calcium and its compounds. But many important distinctions characterize the two groups. Thus the metals of the one group decompose cold water; their oxides are powerfully basic, and readily combine with water; their carbonates are extremely stable; and their sulphates are insoluble or very slightly soluble in water, and do not readily form double salts. The metals of the other group, however, either are without action on water, or decompose it only when heated; their oxides, although basic, exhibit little tendency to combine with water; their carbonates are comparatively unstable; and their sulphates are mostly very soluble in water, and readily form double salts. In the one group the positive character becomes more pronounced as the atomic weight increases, whereas in the other the positive character diminishes with increase of atomic weight.
Glucinum, apparently, is much less closely related to magnesium than is the latter to zinc, and mercury much less closely to cadmium than cadmium to zinc; but zinc and cadmium are more closely related than zinc and magnesium. These conclusions, which result chiefly from the comparison of what may be termed chemical properties, are entirely confirmed by Thomsen’s thermochemical investigation of the reactions involved in the formation of a number of compounds of the metals of this group, as will be evident from the following tables :— [table] [table]
Copper — Silver —G old. [table]
These metals may be regarded as forming a sub-group to the metals of the alkalies.
Copper is a very tenacious, ductile, malleable metal, of red colour, more easily fused than gold, but less fusible than silver. It is without action on water at a red heat, and does not oxidize in the air at ordinary temperatures, but when heated to redness it rapidly combines with oxygen forming a black oxide CuO; when in a finely divided state, it readily unites with chlorine at ordinary temperatures. Moderately diluted nitric acid dissolves copper with great facility, and it is also dissolved by heated concentrated sulphuric acid; hydrochloric acid, even when boiling, has little action on the massive metal, but slowly dissolves the finely divided metal.
Copper forms two chlorides,—cuprous chloride, Cu 2 Cl 2 , and cupric chloride, CuCl 2 , but only one iodide Cu 2 I 2 . Cuprous chloride is a white crystalline substance insoluble in water, but it dissolves in hydrochloric acid or ammonia, forming colourless solutions. Cupric chloride is readily soluble; it is liver-coloured, but its solution in water is blue; it forms crystalline double salts with the chlorides of the alkali metals. It is converted into cuprous chloride and chlorine at a red heat.
Cupric hydroxide, Cu(OH) 2 , separates as a pale blue precipitate on the addition of an alkali to a solution of a cupric salt; in the dry state it is stable at 100° C., but at a slightly higher temperature it is converted into the black oxide. Cupric hydroxide, although insoluble in alkalies, is dissolved by ammonia to a deep blue solution; the production of deep blue-coloured liquids on dissolution in ammonia is, in fact, characteristic of copper compounds. Cupric oxide, CuO, dissolves in acids, forming cupric salts, which are white when anhydrous, but furnish green or blue solutions; they are mostly soluble. Cupric sulphate crystallizes with 5 molecules of water, but it forms a double sulphate with potassium sulphate, CuK 2 (SO 4 ) 2 , 6H 2 O, isomorphous with the corresponding zinc and magnesium salts. Basic cupric salts are readily produced by digesting the normal salts with cupric hydroxide.
By igniting cupric oxide with metallic copper it is converted into red cuprous oxide, Cu 2 O; this oxide may also be prepared by heating a solution of a cupric salt with an easily oxidizable substance, such as grape sugar, and an alkali. Most acids decompose cuprous oxide, forming a cupric salt and separating metallic copper; but it is converted into cuprous chloride by hydrochloric acid. A number of cuprous salts exist, but they are extremely unstable, and mostly absorb oxygen readily, becoming converted into cupric salts. Both oxides of copper are easily reduced by ignition with hydrogen or charcoal.
A black hydrated cupric sulphide is precipitated by hydrogen sulphide even from strongly acid solutions of cupric salts.
Cuprous hydride, Cu 2 H 2 , is deposited as a yellow precipitate when a solution of cupric sulphate mixed with hypophosphorous acid is heated to about 60° C.; it soon turns brown, and when heated to 70° C. suddenly decomposes, with evolution of hydrogen; hydrochloric acid converts it into cuprous chloride :—
Cu 2 H 2 + 2HC1 = Cu 2 Cl 2 + 2H 2 .
Silver is a white metal, inferior in malleablity only to gold among the heavy metals; it is harder than gold, and softer than copper. It is the best-known conductor of heat and electricity. It may be distilled by the aid of the oxy-hydrogen flame. It is not oxidized by ordinary oxygen, and is incapable of decomposing water, but like mercury it is oxidized by ozone. Moderately diluted nitric acid dissolves it readily, and it is dissolved by heated concentrated sulphuric acid; it also resembles copper and mercury in its behaviour with hydrochloric acid, and mercury in its behaviour with hydriodic acid. It unites with the halogens at ordinary temperatures, and, like mercury, readily combines with sulphur, and is precipitated by hydrogen sulphide from strongly acid solutions of its salts as black sulphide, Ag 2 S.
Silver forms two chlorides,—argentic chloride, AgCl, which is white, and insoluble in water, but readily soluble in ammonia, and a black argentous chloride, the composition of which is not known with certainty, but is probably Ag 2 Cl 2 . Corresponding bromides exist, but only the one iodide, Agi, is known. Argentic chloride becomes violetcoloured on exposure to light, apparently owing to the formation of the lower chloride; the iodide does not alter in appearance, but acquires the property of attracting metallic silver, and hence its employment in photography. The subchloride is not changed by nitric acid, but is converted by ammonia into metallic silver and argentic chloride.
A brown argentous hydroxide, Ag (OH), is precipitated from solutions of argentic salts on the addition of an alkali; it is slightly soluble in water and its solution is faintly alkaline; it loses water at a temperature above 60° C., and is converted into argentic oxide, AgO. Argentic oxide is a brown powder; it gives off a certain amount of oxygen, even at 100° C., and also when exposed to sunshine; it neutralizes the strongest acids, forming argentic salts such as argentic nitrate, AgN^O 3 . Argentic sulphate, Ag 2 SO 4 , is isomorphous with anhγdrous sodium sulphate, and it also forms an alum, Ag 2 Al 2 (SO 4 ) 4 , 24H 2 O, which, however, is extremely unstable.
Silver also forms a suboxide probably of the composition Ag 4 O 2 , and a peroxide of which the composition is uncertain, but probably is expressed by the formula Ag 2 O 2 . Both are extremely unstable; the former is decomposed by ammonia into metallic silver and argentic oxide, which dissolves, and acids act upon it in a similar manner.
Argentic carbonate is insoluble in water, and decomposes at 200 C o . into the oxide and carbon dioxide.
Metallic silver is slowly precipitated from solutions of its salts by hydrogen at ordinary temperatures, so that it is a less positive element than hydrogen.
Gold is a metal of rich yellow colour nearly as soft as lead when in a pure state. It exhibits no tendency to combine directly with oxygen, and is not dissolved by any single acid, except selenic acid, but it readily combines with the halogens, and therefore dissolves in a mixture of nitric and hydrochloric acids, which liberates chlorine (p. 466).
Gold is reduced from its compounds with extreme readiness, but varies in properties according to the nature of the compound from which it is separated; in fact, it appears to exist in several allotropic modifications. Thus, according to Thomsen, when a solution of auric chloride is reduced by sulphurous acid, the gold separates as a lightcoloured powder which readily cakes together, but from a solution of auric bromide it separates as an extremely fine dark-coloured powder which exhibits no tendency to agglomerate; a third modification, also finely pulverulent, but lustrous, is obtained by reducing aurous bromide or iodide with sulphurous acid. Thomsen finds that the conversion of the second (Auα) and third (Auj8) of these modifications into that obtained from auric chloride (Au) would involve development of heat, and of different amounts in the two cases; thus :—
AuØ = Au + 4700 units of heat Auα=Au + 3210 „ „
By the action of chlorine, reduced gold is converted into the chloride, Au 2 C1 4 ; this is decomposed by water into aurous chloride, Au 2 C1 2 , and auric chloride, AuCl 3 . Auric chloride is soluble in water, but aurous chloride is insoluble; in contact with water, especially on warming, the latter decomposes into metallic gold and auric chloride. Auric chloride is readily converted into aurous chloride when heated to 185° C. The gold bromides may be obtained in a similar manner, and exhibit similar properties, but are less stable. Aurous and auric iodides, Au 2 I 2 and Aul 3 , are both insoluble, and still less stable. The auric compounds combine with the haloid acids, and with the haloid compounds of most other metals, forming crystalline so-called double salts; these double salts, however, may be regarded as salts of the acids formed by the union of the haloid acids with the auric halogen compounds; the double chloride of gold and potassium, for example, KCLAuCl 3 , as the potassium salt of the acid HAuCl 4 .
On treating aurous chloride with a solution of potassium hydroxide, a dark-green substance is obtained, which is probably the corresponding hydroxide; it decomposes spontaneously into metallic gold and auric hydroxide. Auric hydroxide, Au(OH) 3 , is precipitated by alkalies from a solution of auric chloride as a dark-brown powder, which is quickly reduced on exposure to light. A third gold hydroxide, derived from the oxide Au 2 O 2 , appears also to exist.
The hydroxides of gold exhibit very feeble basic properties; in fact, the salts formed on dissolving auric hydroxide in nitric and sulphuric acids are so unstable that they have not been isolated. Auric hydroxide, however, possesses marked acid properties, and is dissolved by a solution of potassium hydroxide,—potassium aurate, a crystalline salt of the composition KAuO 2 , 3H 2 O, separating from the solution on concentration. On digesting auric hydroxide with a mm onia it is converted into the so-called fulminating gold, which is obtained in the form of a yellowish-brown powder easily exploded by percussion or heat; it is somewhat uncertain what is the composition of this body, but it is not improbable that it is represented by the formula Au (NH 2 ) j θyjq -
The distinctions between cupper, silver, and gold are obviously very marked; yet they resemble each other in many respects, the relation being especially evident when the cuprous and aurons compounds are compared with the argentic compounds. It is to be noted that their hydroxides and many other of their compounds are acted on by ammonia iu a characteristic manner, both silver and gold hydroxides being converted into highly explosive bodies; the nature of the products has not been ascertained with certainty, but apparently they are allied in composition. The ammoniacal derivatives of copper correspond in composition to those formed from nickel, a metal which in many respects is allied to copper.
As in the magnesium group of metals, the positive character diminishes with increase of atomic weight, but to a much greater extent.
Silver is the only element of the group which exhibits marked analogy with the metals of the alkalies,—being connected with them by the isomorphism of its sulphate with that of sodium, by the formation of an alum, by the alkalinity of its hydroxide, and by its colour. Thermochemical investigation also indicates a marked analogy between silver and the metals of the alkalies, the sulphates of silver, sodium, and potassium dissolving in water with absorption of heat, whereas nearly as much heat is developed on dissolving copper sulphate as when the equivalent amount of sulphuric acid is added to water, copper sulphate resembling in this respect the sulphates of magnesium and zinc; thus :— [table]
The existence of double salts such as CuK 2 (SO 4 ) 2 , 6H 2 O, isomorphous with those derived from the sulphates of magnesium and zinc, and the readiness with which basic salts of copper are formed, are confirmatory of the conclusion that copper is allied to zinc and magnesium.
The results of the thermo-chemical investigation of copper, silver, and gold are of considerable interest, as will be evident from the following tables. To illustrate the application of the values in these tables, it may be pointed out that the stability of the oxides of copper and the instability of silver oxide and of auric hydroxide are explained by the fact that much heat is developed in the formation of the oxides of copper from their elements, whereas the formation of silver oxide would involve but a slight development of heat, and the formation of auric hydroxide would even involve a large expenditure of energy. Then it will be noticed that, in the formation of cuprous iodide, heat is developed to the extent of 32,520 units per molecule, and that the formation of an aqueous solution of cupric iodide would involve the development of only 10,410 units per molecule; hence it follows that the conversion of two molecules of cupric iodide into a molecule of cuprous iodide and a molecule of iodine would be attended by the development of no less than 32520 — 2×10410 = 11700 units of heat. This explains the fact that a mixture of cuprous iodide and iodine, and not cupric iodide, is obtained on adding potassium iodide to a solution of cupric sulphate :—
2CuSO 4 + 4KI = Cu 2 I 2 +1 2 + 2K 2 SO 4 .
Again, the superior affinity of auric bromide for hydrogen bromide as compared with that of auric chloride for hydrogen chloride is shown by the development of 7700 units of heat in the formation of the compound AuBr 4 H, and of only 4530 units in the formation of the corresponding chlorine compound AuC1 4 H ; and that the latter is almost completely decomposed and converted into the former on treatment with hydrobromic acid is evidenced by the development of 13,800 units of heat in the reaction, complete decomposition requiring the development of 14,300 units. [table] [table]
The values in this table have reference to the modification of gold obtained by reducing a solution of auric chloride with sulphurous acid.
Thallium — Lead. [table]
Thallium is a heavy crystalline metal, resembling cadmium in appearance; it is extremely soft, even lead scratching it readily. It melts at 294° C., and boils below a white heat; the liquid metal can hardly be distinguished from mercury. It is reduced from solutions of its salts by zinc, but not by cadmium. It readily dissolves in concentrated nitric acid, and is also dissolved by dilute sulphuric acid, especially if it be in contact with a piece of platinum, but it is only slowly attacked even by boiling hydrochloric acid; it decomposes water at a red heat. It combines with the halogens at ordinary temperatures. When exposed to the air it oxidizes almost as readily as the more positive alkali metals, chiefly thallious oxide, T1 2 O, being formed; and when heated to redness and plunged into oxygen it burns brilliantly with a pure green light, chiefly thallic oxide, r Π 2 O 3 , being produced.
Four chlorides of thallium have been described, viz., T1C1 or T1 2 C1 2 , T1 2 C1 4 , Tl 4 Cl 6 , and T1C1 3 or TLCL; corre-spending bromides, and the two iodides, Til or T1 2 I 2 , and T1I 3 or T1 2 I 6 , have also been obtained. Thallious chloride, T1C1 or T1 2 C1 2 , separates as a white precipitate scarcely distinguishable at first sight from silver chloride, on the addition of hydrochloric acid to a solution of a thallious salt; but in properties it more nearly resembles lead chloride, PbCl 2 , being slightly soluble in boiling water, from which it separates again on cooling in crystals, and scarcely more soluble in aqueous ammonia than in water. Thallium sesquichloride, T1 4 C1 6 , formed by melting thallious chloride in a stream of chlorine, crystallizes in brilliant orange-yellow plates, moderately soluble without decomposition in water slightly acidulated with hydrochloric acid. Thallic chloride, T1C1 3 or T1 2 C1 6 , is formed on dissolving thallic oxide in hydrochloric acid; it crystallizes from hydrochloric acid in colourless prisms of the composition T1C1 3 , H 2 O or T1 2 C1 6 , 2H 2 O. Little is known of the chloride T1 2 C1 4 , but the corresponding bromide is obtained on adding thalíious bromide to a hot solution of thallic bromide, separating from the hot liquid in yellow shining needles. Thaliious bromide is nearly insoluble in water, but thallic bromide is very soluble and deliquescent; the bromide Tl 2 Br 4 is decomposed by water into thaliious and thallic bromide and the bromide Tl 4 Br 0 , which crystallizes in orange-red laminae. This last compound is also decomposed by water, with separation of thaliious bromide. Thaliious iodide, TH or T1 2 I 2 , has a brilliant yellow colour, and is almost insoluble in water, and scarcely more soluble in a solution of potassium iodide; when heated it becomes scarlet, but its yellow colour is restored by friction. It is therefore opposite in character to mercuric iodide, Hgl 2 , which is scarlet, but becomes yellow when heated, and is readily soluble in potassium iodide solution, although insoluble in water. Thallic iodide is an extremely unstable compound, and, like cupric iodide, readily decomposes into iodine and thaliious iodide.
Thaliious oxide is almost black in colour; like litharge, PbO, the corresponding lead oxide, it fuses readily; and the fused oxide is absorbed by bone-ash so that a silver thallium alloy may be cupelled like a silver-lead alloy. It dissolves readily in water, forming a colourless, powerfully alkaline solution, greasy to the touch, of thaliious hydroxide, T1(OH) or T1 2 (OH) 2 . The latter compound may be obtained in pale-yellow prismatic crystals; it is reconverted into the oxide by mere exposure over sulphuric acid in a vacuum at the ordinary temperature. The solution of thaliious hydroxide has nearly all the properties characteristic of sodium hydroxide; but it has a slight tendency to absorb oxygen. Thallic hydroxide, T1(OH) 3 or T1 2 (OH) 6 , separates as a brown gelatinous precipitate on the addition of alkalies to a solution of a thallic salt; by heating to about 260° C. it is converted into thallic oxide, T1 2 O 3 . This oxide is a dark brown powder fusible with difficulty; it is reduced to the lower 'oxide at a red heat. It has much less marked basic properties than thaliious oxide, but dissolves readily in acids, forming thallic salts; it is insoluble in water. Thaliious salts are not precipitated by hydrogen sulphide in presence of strong acids, but the thallium is completely precipitated as a brown sulphide from solutions of the salts of weak acids, such as thaliious acetate, and by alkaline sulphides.
Thaliious hydroxide readily absorbs carbon dioxide, formmg thaliious carbonate, T1 2 CO 3 . This salt crystallizes in prisms, and is soluble in about four times its weight of boiling water, forming an alkaline solution; it is fusible, but decomposes when heated to dull redness. Thaliious sulphate, T1 2 SO 4 , is isomorphous with potassium sulphate; it also forms an alum, T1 2 A1 2 (SO 4 ) 4 , 24H 2 O, and double salts such as Tl 2 Mg(SO 4 ) 2 , 6H 2 O, isomorphous with the corresponding potassium salts. A large number of thaliious salts, in fact, are isomorphous with the corresponding sodium, potassium, and ammonium salts. Thallic sulphate, T1 2 (SO 4 ) 3 , 7H 2 O, is decomposed even by cold water; it forms an anhydrous double sulphate with potassium sulphate quite different from the alums in composition, viz., T1 2 (SO 4 ) 3 , 2K 2 SO 4 .
Lead is a bluish-white, extremely soft metal; it fuses at 325° C. It is readily dissolved by nitric acid, and is slowly acted on when boiled with hydrochloric or dilute sulphuric acid, hydrogen being evolved. Although it tarnishes in the air, the oxidation is only superficial, but it readily absorbs oxygen when heated. Its oxides are very easily reduced.
Only one stable chloride of lead, PbCl 2 is known; when, however, the oxides Pb 2 O 3 and PbO 2 are dissolved in hydrochloric acid at a low temperature, solutions are formed which apparently contain the corresponding chlorides Pb 2 Cl 0 and PbCl 4 , as the oxides may be reprecipitated by alkalies, but the slightest heat causes chlorine to be evolved. Lead cldoride is difficultly soluble in water, from which it crystallizes in delicate white needles. Lead iodide, Pbl 2 , crystallizes in brilliant yellow plates sparingly soluble in water.
Lead forms a suboxide, Pb 2 O, a monoxide, PbO, a dioxide or peroxide, PbO 2 , and a fourth oxide, Pb 3 O 4 , which may be regarded as a compound of the monoxide and dioxide. The suboxide is black; like the corresponding oxides of mercury and silver, it furnishes the metal and a salt derived from the monoxide on treatment with acids. The monoxide, PbO, ordinarily known as litharge, is prepared on the large scale by the oxidation of lead in air; the pure oxide has a lemon-yellow colour, and is insoluble in water; it fuses at a heat below redness. The corresponding hydroxide, Pb(OH) 2 , is obtained on adding ammonia to a solution of a lead salt such as lead nitrate, Pb(NO 3 ) 2 , as a white amorphous precipitate, which absorbs carbon dioxide on exposure to air; it is converted into the oxide when heated above 100° C. The oxide and hydroxide are readily acted on by acids, and form stable salts; but they are also soluble in alkalies. Lead nitrate is easily soluble, but lead sulphate and carbonate are insoluble in water; they may be obtained in crystals isomorphous with those of the corresponding barium salts. The carbonate is readily decomposed when heated, and on heating the sulphate with lead sulphide metallic lead is obtained: PbSO 4 + PbS = 2Pb + 2SO 2 .
Lead sesquioxide, Pb 2 O 3 , is little known; it has a brown colour, and is converted into the monoxide and oxygen when heated. Lead dioxide, PbO 2 , also has a brown colour; it may be obtained by passing chlorine into water in which lead oxide is suspended. It unites directly with sulphur dioxide, forming lead sulphate, sufficient heat being developed to cause the mass to glow; it is converted into the monoxide when heated; it is insoluble in most acids, but is dissolved by acetic acid, and apparently converted into a corresponding acetate; by fusing it with potassium hydroxide, potassium plumbate is produced, which may be obtained in crystals of the composition K 2 PbO 3 , 3H 2 O.
By carefully heating lead monoxide to low redness in air it is converted into the red oxide or minium, the most usual composition of which is represented by the formula Pb 3 O 4 . On treating this oxide with acids—nitric acid, for example—it furnishes lead nitrate and lead dioxide.
Lead is precipitated by hydrogen sulphide, as a black sulphide, even from strongly acid solutions of its salts.
Although no stable chloride of lead exists containing more than two atoms of chlorine, a stable tetraethyl derivative^
Pb(C 2 H 5 ) 4 , the existence of which affords conclusive evidence that lead may function as a tetrad element, is obtained by the action of zinc ethyl on the chloride PbCl 2 (1Pb 2 Cl 4 ). The hydroxide, Pb(C 2 H 5 ) 3 .OH, is powerfully basic.
Compounds of thallium with hydrocarbon radicles only have not yet been obtained; a thallium diethylchloride, Tl(C 2 H δ ) 2 Cl, is known, however. It is a crystalline body Of remarkable stability, and is readily converted into the hydroxide, Tl(C 2 H 5 ) 2 ∙OH, which is also crystalline and very soluble in water; the solution is strongly alkaline, nevertheless the hydroxide does not exhibit the slightest tendency to attract carbon dioxide, although it readily forms salts with other acids.
The relation of thallium to lead will be sufficiently evident'from the above sketch of the properties of some of the more important compounds of the two elements; it will be noticed also that thallium has strong points of resemblance with the metals of the alkalies, and that it is in some respects related to mercury. Its relation to the alkali metals is especially indicated by the extreme readiness with which it is oxidized, by the properties of thallious hydroxide, and by the isomorphism of many thallious compounds with the corresponding potassium and sodium -compounds. Thermochemical investigation entirely confirms these conclusions, as will be evident from the following tables, the heat of neutralization of thallium hydroxide being equal to that of the hydroxides of potassium and sodium, but the affinity of thallium for oxygen being only about equal to that of mercury and copper, and somewhat less than that of lead. The heat of neutralization of thallic hydroxide exceeds that of aluminic hydroxide. [table] [table]
Boron — Aluminium — Indium. [table]
The first of these elements has already been described. Aluminium is related not only to boron but also to silicon among the non-metals, and to magnesium among the metals. Its oxide is one of the most stable known, and, like silicon and boron oxides, cannot be reduced by charcoal alone. The metal is usually obtained by heating the double chloride of sodium and aluminium, 2NaCl,Al 2 Cl β , with sodium. Aluminium is a white malleable metal, nearly resembling zinc in colour and hardness; it fuses at a lower temperature than silver, but does not volatilize. It is rapidly dissolved by hydrochloric acid, but only slowly acted on by heated nitric and sulphuric acids; it is readily dissolved by solutions of the alkalies. In a finely-divided state it appears to decompose water almost as easily as magnesium, but when in mass, even if heated to full redness in steam, it decomposes water but slowly, owing to the formation of a coating of oxide on its surface which protects the metal; on this account also it may be heated intensely in air without undergoing more than a superficial oxidation, but in the form of powder it burns brightly when heated to redness in air or oxygen. Aluminium takes fire when heated in chlorine gas, and is converted into the chloride, which is also produced on passing chlorine over an ignited mixture of alumina and charcoal.
Aluminium forms only one chloride, the density of which in the state of vapour corresponds with the formula Al 2 Clθ; it also forms only a single oxide, A1 2 O 3 . Aluminium chloride is a colourless, crystalline, volatile substance; it is deliquescent and dissolves readily in water, much heat being developed, forming a strongly acid solution. Aluminium hydroxide, Al 2 (OH) β , separates as an almost colourless gelatinous precipitate on the addition of alkalies to a solution of an aluminium salt; it is readily soluble in acids, forming salts such as aluminium sulphate, A1 2 (SO 4 ) 3 . Aluminium sulphate is extremely soluble in water; it is very readily converted into basic salts; it combines with the sulphates of the alkali metals forming characteristic double salts—the so-called alums.
Aluminium hydroxide also possesses marked acid properties, being soluble in alkalies, and furnishing salts such as sodium aluminate, Na β Al 2 O 6 . On ignition aluminium hydroxide is converted into aluminium oxide or alumina, A1 2 O 3 ; if it has been exposed only to a moderate red heat the alumina is a soft, white, light powder like magnesium oxide, and is soluble in acids; but after strong ignition it becomes very hard and insoluble, or difficultly soluble, in acids. Alumina occurs native as corundum in crystals isomorphous with the corresponding oxides of chromium and iron; in the crystalline state it is the hardest substance known next to the diamond, and insoluble in all acids, but like silica it is dissolved on fusion with alkalies.
The only reactions which, in the case of aluminium, have been submitted to thermochemical investigation are— [table]
The resemblance of aluminium to boron, silicon, and magnesium is most evident when the oxides and chlorides are compared. Thus, their oxides are all bodies of great stability, and, excepting magnesium oxide, possess marked acid properties; and their chlorides behave similarly when added to water, boron and silicon chlorides being entirely decomposed, and aluminium and magnesium chlorides at least partially, the reaction in each case being attended by the development of a very considerable amount of heat; thus, according as quantities are taken which are as the molecular weights of the chlorides, or which contain the same amount of chlorine, we have— [table]
It may be remarked also that magnesium and aluminium both form very soluble sulphates, which readily combine with other sulphates.
Indium is easily reduced from its oxide by hydrogen, and is precipitated from solutions of its salts by cadmium and zinc. It is an extremely rare metal, and occurs always associated with zinc. It is a soft, ductile, white metal, destitute of crystalline structure, much resembling lead in appearance. It rapidly dissolves in concentrated hydrochloric acid, slowly in dilute nitric and sulphuric acids. It melts at 176’ C., but is much less volatile than cadmium, and may be melted in the air without oxidizing; at a bright red heat it burns with a violet flame forming a yellow oxide, In 2 O 3 .
Indium chloride, In 2 Cl 6 , sublimes without fusion at an incipient white heat, condensing in soft white laminae. Indium oxide, In 2 O 3 , is not volatile even at a white heat; it is readily soluble in acids. It is stated that by heating in a stream of hydrogen the oxides In 7 O 9 , In 4 O 5 , andInO 2 are in succession formed from it, but their existence as definite compounds has not been satisfactorily established. Indium hydroxide separates as a white gelatinous precipitate on the addition of ammonia to cold solutions of indie salts; it is soluble in alkalies, and on ignition is converted into the oxide In 2 O 3 . Indium sulphate forms with ammonium sulphate a double sulphate In 2 (NH 4 ) 2 (SO 4 ) 4 ,24H 2 O, isomorphous with the alums; when heated only to 36° C., however, this salt is converted into one containing only 4 molecules of water of crystallization.
On passing hydrogen sulphide into solutions of indium a yellow precipitate is obtained, which is decomposed by hydrochloric acid; no precipitate is produced when aluminium salts are similarly treated, aluminium sulphide being at once decomposed by water and converted into the hydroxide.
Chromium — Manganese — Iron — Cobalt — Nickei λ [table]
The general resemblance between these metals is very great, but each exhibits properties which serve to distinguish it readily from the others. They are also related to many other elements; thus, chromium and iron are closely allied to aluminium; and chromium is also allied to sulphur and its congeners; manganese manifests some analogy with the halogens; nickel has properties in common with copper: and iron, cobalt, and nickel are closely related to the so-called platinum metals.
The oxides of chromium and manganese are reduced by charcoal only at a white heat; the oxides of the remaining metals are easily reduced on ignition with hydrogen or carbon.
The properties of chromium appear to differ considerably according to the manner in which it is prepared. Thus, by heating its chloride with potassium it is obtained as a dark grey powder, which oxidizes readily, taking fire in the air at a heat below redness, and dissolving in dilute sulphuric and hydrochloric acids. By heating its chloride with sodium, however, it is obtained in hard, shining crystals, insoluble in all acids. The metal prepared by reducing the oxide with charcoal is shining, has a steel-grey colour, and is hard enough to scratch glass.
Manganese is a soft, brittle, greyish-white metal, which oxidizes quickly on exposure to the air, decomposes water slowly at ordinary temperatures, and dissolves easily in acids; it is feebly magnetic.
The properties of iron are too well known to need description. It dissolves readily in acids, and decomposes water rapidly at a red heat, it does not alter in dry air, but rapidly oxidizes in moist air; especially in presence of carbonic acid.
Cobalt has a steel-grey colour; nickel is silver-white. Both are hard, ductile, malleable metals, of great tenacity, somewhat more fusible than iron; they decompose water at a red heat, and like iron dissolve without difficulty in acids. They are not altered by the combined action of air and water at ordinary temperatures, but if previously heated burn in oxygen like iron. Nickel appears to be less oxidizable than cobalt. They are both magnetic metals.
Each of these metals forms several oxides; and hydroxides and chlorides, or fluorides, corresponding to most of their oxides are known, as will be evident from the following list :— [table]
[table]
Their chlorides, and, in fact, their halogen compounds generally, are all easily soluble in water, with the exception of sublimed chromic chloride, Cr 2 Cl 6 , which is insoluble in cold, and is only slowly dissolved by boiling water; but if the cold solution contains a minute quantity of chromous chloride, Cr 2 Cl 4 , the chromic chloride dissolves immediately with development of heat, forming a green solution identical with that produced by dissolving chromic hydroxide in hydrochloric acid; this effect is, perhaps, due to the formation in the first instance of an intermediate chloride, or combination of the two chlorides, which is decomposable by water. The chlorides which correspond in composition to ferrous chloride, Fe 2 Cl 4 , are all perfectly stable, and volatilize without decomposition; chromic and ferric chlorides are also highly stable bodies, but the remaining chlorides are exceedingly unstable. Chromium hexafluoride, CrF 6 , is a very volatile blood-red liquid which is decomposed by water.
Their oxides, with the exception of chromium trioxide or chromic anhydride, CrO 3 , and hydroxides are all insoluble in water, but readily dissolve in acids. Chromic oxide, Cr 2 O 3 , is green, and chromium trioxide has a magnificent scarlet colour. Ferric oxide, Fe 2 O 3 , has a reddish-brown colour; the oxides of manganese are brownish black; and the oxides of nickel and cobalt are all intensely black like cupric oxide. Chromic and ferric oxides form crystals isomorphous with those of aluminic oxide, A1 2 O 3 .
The affinity of chromium to oxygen is so great that chromous oxide, Cr 2 O 2 , is apparently incapable of existing, and it is difficult to obtain ferrous oxide, Fe 2 O 2 , on account of the readiness with which it is oxidized to ferric oxide, Fe 2 O 3 . Ferric oxide, like chromic oxide, Cr 2 O 3 , is a body of great stability, but unlike the latter it is converted into a lower oxide, ferroso-ferric oxide, Fe 3 O 4 , on ignition. The most stable oxide of manganese is the corresponding oxide, Mn 3 O 4 ; but manganous oxide, Mn 2 O 2 , is less readily oxidized than ferrous oxide, and manganic oxide is less stable than ferric oxide; manganese is the only metal of the iron group which forms a dioxide, or, as it is usually termed, a peroxide. Cobaltous oxide, Co 2 O 2 , absorbs oxygen when heated to dull redness in air, and is converted into the oxide Co 3 O 4 , but when more strongly heated this oxide is reconverted into cobaltous oxide, which is also obtained on igniting cobaltic oxide, Co 2 O 3 . Nickelous oxide, Ni 2 O 2 , exhibits no tendency to absorb oxygen, and the higher oxide is very readily decomposed by heat. From this it will be evident that the affinity to oxygen diminishes rapidly from chromium to nickel.
Similar differences are noticeable between their hydroxides. Thus, the tendency of chromous hydroxide, Cr 2 (OH) 4 , to become converted into chromic hydroxide, Cγ 2 (OH) 6 , is so great that it decomposes water even at ordinary temperatures with evolution of hydrogen: Cr 2 (OH) 4 + 2OH 2 = Cr 2 (OH) 6 + H 2 . Ferrous hydroxide does not decompose water in this manner, but in presence of both air and water it is quickly converted into ferric hydroxide: 2Fe 2 (OH) 4 + 4OH 2 + O 2 = 2Fe 2 (OH) 6 + 2OH 2 ; manganous hydroxide behaves similarly, but oxidizes much less rapidly; cobaltous hydroxide manifests but little tendency to form the higher hydroxide, and nickelous hydroxide is unaffected, when placed in contact with air and water. The two last named hydroxides, however, may be converted into the higher hydroxides by treatment with a solution of sodium hvDOchlorite.
Their hydroxides are without difficulty converted into the corresponding oxides, giving off water when very gently heated; the behaviour of chromous hydroxide is peculiar, as, when ignited, it is converted into chromic oxide with evolution of hydrogen: Cr 2 (OH) 4 = Cr 2 O 3 + OH 2 + H 2 . Many hydroxides are known intermediate in composition between the oxides and the highest hydroxides,—such, for example, as are represented by the formulae Fe 2 O(OH) 4 and Fe 2 O 2 (OH) 2 , which obviously are intermediate in composition between ferric oxide, Fe 2 O 3 , and ferric hydroxide, Fe 2 (OH) 6 .
With the exception of chromium trioxide, the oxides and hydroxides of the metals now under consideration all exhibit marked basic properties; but chromic hydroxide, like aluminie hydroxide, also has feeble acid properties, dissolving in solutions of the alkalies; and cobaltous and nickelυus hydroxides, like cupric hydroxide, readily dissolve in ammonia, the solution of the former being brownish-red, and that of the latter blue. The oxides and hydroxides which correspond in composition with ferrous oxide and hydroxide furnish stable salts such as ferrous sulphate, Fe 2 (SO 4 ) 2 . The chromous salts, however, are but little known; they form either red or blue solutions, which absorb oxygen with avidity, and dissolve nitric oxide gas, forming dark brown solutions. The ferrous salts form pale green solutions, which slowly absorb oxygen, and readily dissolve nitric oxide (p. 513). The manganous salts are of a pale pink colour, and their solutions are scarcely altered in contact with oxygen. Solutions of cobaltous salts have a rose-red colour, unless they are very concentrated or mixed with a strong acid, in which case they are blue. Nickelous salts are green. The solutions of cobaltous and nickelous salts are permanent in air.
A second series of chromium and iron salts corresponding to the oxides Cr 2 O 3 and Fe 2 O 3 are readily obtained, as they are of great stability. Solutions of chromic salts are green or violet; solutions of ferric salts are yellow. Chromic salts such as chromic sulphate, Cr 2 (SO 4 ) 3 , for example, are reduced to chromous salts by the action of the nascent hydrogen produced when zinc is immersed in their solutions, but they cannot be thus reduced by the aid of iron; ferric salts, however, are readily reduced to ferrous salts even by feeble reducing agents such as hydrogen sulphide. Manganic and cobaltic salts corresponding to the chromic and ferric salts may be obtained, but they are extremely unstable, and nickelic salts appear to be incapable of existing, so that, on treating the higher oxides and hydroxides of manganese, &c., with acids, salts derived from manganous oxide and its congeners are usually obtained. Thus, nickelic oxide, Ni 2 O 3 , is converted by the action of sulphuric acid into nickelous sulphate, with evolution of oxygen; similarly, when manganic oxide, Mn 2 O 3 , manganic peroxide, MnO 2 , or cobaltic oxide, Co 2 O 3 , is carefully dissolved in hydrochloric acid, solutions are obtained which apparently contain the corresponding chlorides, Mn 2 Cl 6 , MnCl 4 , and Co 2 Cl β , but these rapidly decompose with evolution of chlorine, leaving manganous chloride, Mn 2 Cl 4 , or cobaltous chloride Co 2 Cl 4 . The oxides which correspond in composition to ferroso-ferric oxide, Fe 3 O 4 , behave with acids as if mixtures of the two oxides. Salts corresponding to chromium trioxide are not obtainable, for on treatment with sulphuric acid it is converted into chromic sulphate, with evolution of oxygen; and by the action of hydrochloric acid it is converted into chromic chloride, with evolution of chlorine. Like all feebly positive metals, the metals of the iron group form numerous basic salts, iron and chromium being especially characterized by the formation of basic chromic and ferric salts; a remarkable series of ferric and chromic “mixed salts,” such, for example, as ferric tetraceto-dinitrate, Fe 2 (NO 3 ) 2 (C 2 H 3 O 2 ) 4 , and chromic pentacetonitrate, Cr 2 (NO 3 )(C 2 H 3 O 2 ) 5 , have been obtained by dissolving the hydroxides in mixtures of two acids in various proportions. Like aluminic sulphate, chromic, ferric, and manganic sulphates form double salts with the sulphates of the alkali metals, which correspond both in composition and crystalline form with the alums. The chrome alums are almost as stable as the true alums, but the iron alums are much less stable, and the manganese alums are so unstable that water decomposes them into their component salts. Ferrous, manganous, cobaltous, and nickelous sulphates form double sulphates with the sulphates of the alkali metals, corresponding in composition and isomorphous with double sulphates such as K 2 Zn(SO 4 ) 2 , 6H 2 O and K 2 Cu(SO 4 ) 2 , 6H 2 O.
The relation of chromium to sulphur and the allied elements is indicated by the existence of the trioxide CrO 3 ; this oxide, however, is not merely analogous in composition to sulphuric anhydride, but also resembles it in properties. Thus, it dissolves readily in water, forming a strongly acid, yellow solution of chromic acid, H 2 CrO 4 ; and the salts of this acid are not only analogous in composition to the corresponding sulphates, but are also isomorphous with them, and are similar to them in many respects, the chromates of metals which form insoluble or difficultly soluble sulphates being as a rule also insoluble or difficultly soluble. But although the chromates are stable salts, chromic acid is not known except in solution, chromic anhydride, CrO 3 , being obtained when the solution is evaporated; moreover, chromic acid very readily parts with oxygen, and on this account is a most valuable oxidizing agent. Acid salts of chromic acid corresponding to the acid sulphates apparently cannot exist, for when a solution of a chromate is rendered acid the colour changes from pale yellow to orange-yellow, the chromate being converted into an anhydrochromate or dichromate, which is formed from two molecules of the acid chromate by the withdrawal of the elements of a molecule of water; thus — [table]
The anhydrochromates are mostly orange-red, the corresponding chromates being yellow; they are very stable, and unlike the anhydrosulphates (p. 507), to which they correspond in composition, dissolve in water unchanged.
The relation of chromium to sulphur is also indicated by the existence of an oxychloride, CrO 2 Cl 2 , corresponding to sulphuric chloride (p. 505), SO 2 C1 2 ; it is a blood-red coloured liquid, which is violently decomposed by water, and boils at 118 β C. The determination of the vapour density of this compound shows that it is correctly represented by the formula CrO 2 Cl 2 .
On fusing an oxide of manganese with potassium nitrate, a deep green-coloured mass is formed, from which potassium manganate, K 2 MnO 4 , may be separated in crystals isomorphous with potassium sulphate. A concentrated solution of this salt is decomposed merely by dilution, and the solution cannot be boiled; on adding an acid the colour changes from green to purple-red, the manganate being converted into permanganate :— [table]
The permanganates are much more stable than the manganates, and their solutions may be boiled without their decomposing; but they readily part with oxygen, and are therefore most powerful oxidizing agents
As the manganates and sulphates are isomorphous it is probable that the two classes of compounds are analogous in constitution; and from the manner in which the permanganates are formed from the manganates it is probable that the former contain at least two atoms of manganese in the molecule, and that potassium permanganate, for example, is represented by the formula ∙ , r √ z . '. But potassium
0.Mn0 2 (0K) t permanganate is isomorphous with potassium perchlorate and periodate; it therefore appears probable that, as already pointed out (p. 497), the formula K 2 C1 2 O 8 for the former salt is preferable to the simpler formula KC1O 4 , by which it is usually represented. The isomorphism of potassium permanganate with potassium perchlorate and periodate is, however, chiefly of interest as an indication that manganese is related to the halogens.
Iron is also capable of yielding a compound analogous to potassium manganate, but much less stable, termed potassium ferrate, K 2 FeO 4 ; it is produced by fusing an oxide of iron with nitre. The solution has a cherry-red colour, and on the addition of acids is at once decomposed with evolution of oxygen. The most characteristic compounds of iron, however, are the so-called ferrocyanides and ferricyanides. Potassium ferrocyanide, K 8 Fe 2 (CN) 12 , is produced on adding potassium cyanide to a solution of a ferrous salt in sufficient amount to redissolve the precipitate which at first forms; it crystallizes from the concentrated solution in beautiful yellow plates. The tendency to form this salt is so great that metallic iron is rapidly dissolved when heated with an aqueous solution of potassium cyanide: —
12KCN + 2Fe + 4H 2 O = K 8 Fe 2 (CN) 12 + 4KHO + 2H 2 .
. Not only is the presence of iron in this compound not discoverable by the ordinary tests, but on treatment with hydrochloric acid it furnishes hydroferrocyanic acid, H 8 Fe 2 (CN) 12 , and is not, like most double cyanides, decomposed with evolution of hydrocyanic acid. Potassium ferricyanide, K 6 Fe 2 (CN) 12 , is produced by the action of chlorine on the ferrocyanide: K 8 Fe 2 (CN)ι 2 + Cl 2 = K 6 Fe 2 (CN) 12 ÷ 2KC1; it crystallizes in red prisms, and like the ferrocyanide is an extremely stable compound. Other ferro-and ferri-cyanides may be prepared from the potassium compounds by double decomposition.
Cobalticyanides, such as K 6 Co 2 (CN) 12 , isomorphous with the ferricyanides, are readily obtained, and are equally stable; but the cobaltocyanides are very unstable. Chrom i-cyanides and mangano-and mangani-cyanides, isomorphous with the corresponding iron compounds are also known, but they are much less stable than the latter. Nickel does not furnish compounds of this kind, but like copper forms double cyanides, such as Ni(CN) 2 , 2KCN, which are readily decomposed by acids with evolution of hydrocyanic acid.
Nickel, although of the same atomic weight as cobalt, it will be manifest, is possessed of very different properties, being in many respects much more nearly related to copper than to cobalt, iron, &c.; it is noteworthy that nickel and copper are the only metals whose compounds form blue solutions with ammonia. The relation of chromium, iron, manganese, and cobalt to one another, and of the first two especially to aluminium, will be sufficiently evident without further remark; the relation of iron, cobalt, and nickel to the platinum metals will be referred to later on.
Manganese and iron are the only elements of this group which have been submitted to thermochemical investigation. Many of the results are of considerable interest; thus, it will be noticed, on reference to the following tables, that much more heat is developed in the formation of riaugaivíus chloride, MnCl 2 or Mn 2 Cl 4 , than in that of the corresponding chloride of iron; this is of interest in connection with and may perhaps serve to explain, the great instability of the higher chlorides of manganese.
[table] [table]
The only two compounds of the metals of the iron group of which the vapour density has been determined are ferric chloride and chromium oxychloride; the density of the former corresponds to the formula Fe 2 Cl 6 , and that of the latter to the formula CrO 2 Cl 2 . Hence the composition of ferric chloride is similar to that of aluminic chloride, the density of which corresponds to the formula Al 2 CL, and it is therefore probable that chromic chloride, which is in all respects analogous to aluminic and ferric chlorides, is similarly constituted, and that its formula is Cr 2 Cl 6 .
The majority of the formulae employed to represent chemical compounds which cannot be converted into gas are merely the simplest expressions of their composition in terms of the atomic weights of their constituent elements. Thus it is usual to express the composition of the lower chlorides of chromium, iron, manganese, <fcc., by the simple formulae CrCl 2 , MnCl 2 , FeCl 2 , and the corresponding oxides by the similar formulae CrO, MnO, FeO; there are reasons, however, which lead us to suppose that in many cases the less simple formulae, such as Cr 2 Cl 4 , Cr 2 O 2 , are to be preferred, and that the lower ch!orides and oxides of iron, &c., like the higher chlorides and oxides, contain at least two atoms of the metal in the molecule. For example, the lower chlorides of chromium and iron are readily converted into the higher chlorides by the action of chlorine; now if the formula of ferrous chloride is FeCl 2 , that of ferric chloride being Fe 2 Cl β , the conversion of the former into the latter by the action of chlorine involves the extremely improbable assumption that a molecule of the ferric compound is formed from two molecules of the ferrous compound :—
2FeCl 2 + Cl 2 = Fe 2 Cl 6 .
Chlorine is not known ever to produce an effect of this kind, that is to say, directly to cause the formation of a more complex from a less complex molecule; but its action is frequently to produce an effect the very reverse of this, many instances of simplification of the molecule by the action of chlorine being known. By adopting the formula Fe 2 Cl 4 for ferrous chloride a natural interpretation of its conversion into ferric chloride is at once possible, however :—
FeCl 2 FeCl 3
II + cι s = I l'eCl 2 Fed,
Ferrous chloride. Ferric chloride.
Corresponding formulae may be assigned to the lower chlorides of chromium, manganese, cobalt, and nickel, on account of their resemblance to ferrous chloride; and from the general resemblance which the nickel compounds bear to the cupric compounds it appears probable that, if nickel chloride has the formula Ni 2 Cl 4 , cupric chloride has the formula Cu 2 Cl 4 .
The argument may be extended much further, and from the more or less perfect resemblance of silver chloride to cuprous chloride, mercurous chloride, and other chlorides which certainly contain at least two atoms of chlorine in their molecules, we may infer that the formula Ag 2 Cl 2 is probably a more correct expression than that commonly employed. The existence of a subchloride and a suboxide of silver indicates that silver is not uniformly a monad element, and if we regard copper, silver, and gold as triad elements (although the last mentioned is undoubtedly capable of assuming pentad functions) we may represent their lower chlorides by similar formulae, and thus, in a measure, account for the similarity in their properties :—
CuCl AgCl AuCl
II II II
CuCl AgCl AuCl
Cuprous chloride. Argentic chloride. Aurous chloride.
But it has been pointed out that certain of the silver compounds are isomorphous with the corresponding sodium and potassium compounds, and as there is no reason to suppose that the argentic compounds generally are not of similar constitution to argentic chloride, the conclusion that there are two atoms of silver in the molecule of argentic chloride would appear to necessitate the assumption that the compounds of the alkali metals also contain at least two atoms of the metal in their molecules;—that the formula of sodium nitrate, for example, is Na 2 (NO 3 ) 2 , and not NaNO 3 . The existence of a compound of sodium ethyl with zinc ethyl, NaZn(C 2 H 5 ) 3 , from which the former cannot even be separated, appears to show that sodium does not invariably function as a monad.
Speculations of this kind are extremely hazardous, but, as no method is known enabling us to ascertain the molecular composition of compounds which cannot be volatilized, they possess a certain interest and, moreover, point to the necessity of investigation in this direction.
Ruthenium — Rhodium — Palladium — Osmium — Iridium — Platinum. [table]
These metals always occur in the native state, and are usually associated together. The relation between ruthenium, rhodium, and palladium, and between osmium, iridium, and platinum, is similar to that which obtains between iron, cobalt, and nickel; they are indeed very closely related to these last elements,—the resemblance in properties being greatest between iron, ruthenium, and osmium; between cobalt, rhodium, and iridium; and between nickel, palladium, and platinum. Palladium and platinum also manifest much analogy to silver and gold, and ruthenium and osmium exhibit properties which connect them with the elements molybdenum and tungsten.
They are white, or nearly white, hard metals, of high specific gravity, osmium being the heaviest of the elements; they may be melted, but only at a very high temperature with the aid of the oxyhydrogen blowpipe, with the exception of osmium, which does not melt, but volatilizes, palladium being more fusible, and rhodium, ruthenium, and iridium much less fusible, than platinum. Palladium is soluble in concentrated nitric acid, and platinum in nitrous acid, as well as in aqua-regia; the remaining metals are difficultly soluble or are insoluble in aqua-regia. They are all attacked on fusion with alkalies and potassic nitrate. Most of them possess in a high degree the property to which allusion has already been made of condensing gases on their surfaces, especially when in a finely divided state; thus, finely divided platinum at ordinary temperatures absorbs about 240 times its volume of hydrogen, and finely divided palladium at 200° C. absorbs nearly 700 times its volume of hydrogen. W!ιen heated in an atmosphere of the gas platinum and palladium also occlude hydrogen when in a compact state. The latter metal, according to Troost and Hautefeuille, forms a definite compound with hydrogen of the composition Pd 2 H, which is further capable of absorbing hydrogen in considerable quantity; the density of the pure compound is 11 06, that of the melted metal from which it was prepared being 12, and, therefore, if it be supposed that the elements unite without condensation, that is to say, that the hydrogen has the same density as it would have if solidified in the free state, the density of solidified hydrogen is '62, or slightly higher than that of lithium. Iron, cobalt, and nickel also possess the property of occluding hydrogen, although only in a slight degree.
The following chlorides and oxides of the platinum metals are known :— [table]
Ruthenium absorbs oxygen at a red heat, and is converted into the oxide, Ru 2 O 3 , which is its most stable oxide, whereas the dioxide is the most stable oxide of osmium. Ruthenium tetroxide, RuO 4 , is a golden-yellow crystalline substance, sparingly soluble in water, which melts at 58’ C., and boüs a little above 100* C.; hydrochloric acid converts it into the chloride Ru 2 C1 6 , with evolution of chlorine. The solution of this chloride is decomposed by boiling, the hydroxide Ru 2 (OH) 6 being precipitated; like ferric chloride, it is reduced to the chloride Ru 2 C1 4 by hydrogen sulphide, the colour changing from orange-yellow to blue.
Osmium, in the compact state, is slowly converted into the tetroxide, OsO 4 , when roasted in air, but the pulverulent metal obtained by reducing this oxide by hydrogen, oxidizes spontaneously in the air. Osmium tetroxide is a white, crystalline, extremely volatile substance; it melts considerably below 100’ C., and boils at a temperature a little above its melting point. The vapour density determination shows that it is correctly represented by the formula OsO 4 . Its vapour has an intolerably pungent odour, and is excessively poisonous. It is dissolved by water, and is a powerful oxidizing agent, but does not evolve chlorine on treatment with hydrochloric acid; most metals, even silver and mercury, excepting gold and the platinum metals, partially reduce its solution and cause the separation of osmium; it is decomposed by ammonia, being converted into the hydroxide Os(OH) 4 , with evolution of nitrogen: 3OsO 4 + 4NH 3 = 3Os(OH) 4 + 2N 2 . A lower hydroxide of osmium, Os 2 (OH) 4 , is known, which like ferrous hydroxide rapidly oxidizes on exposure to the air.
The tetroxides of ruthenium and osmium are apparently perfectly neutral bodies, possessing neither basic nor acid properties; the remaining oxides exhibit only feeble basic properties, and furnish unstable salts, of which very little is. known. The oxides of ruthenium are readily reduced when heated in hydrogen; osmium tetroxide is reduced at a red heat, but the lower oxides of osmium are reduced by hydrogen at ordinary temperatures.
Ruthenium and osmium furnish ruthenio-and osmio-cyanides, such as K 8 Ru 2 (CN) 12 and K 8 Os 2 (CN) 12 , isomorphous with the corresponding ferrocyanides, from which the acids H 8 Ru 2 (CN) 12 and H 8 Os 2 (CN) 12 may be prepared; the compounds analogous to the ferricyanides are not known. On heating the metals with potassium nitrate, compounds analogous in composition to potassium ferrate, but more stable, are formed; potassium rutheniate, K 2 RuO 4 , is apparently much less stable than potassium osmate, K 2 OsO 4 , which may be obtained in large rose-coloured crystals’; their solutions are at once decomposed by acids, the former yielding the hydroxide Ru 2 (OH) 6 with evolution of oxygen, and the latter the hydroxide Os(OH) 4 and osmium tetroxide.
The chlorides of ruthenium and osmium form numerous double salts, but the most important are the ruthenio-and osmio-chlorides, such as K 2 RuC1 6 and K 2 OsCl 6 , which are isomorphous with the corresponding platinochlorides.
Iridium and rhodium are insoluble in aqua-regia, and even when in a finely divided state are only slowly oxidized at a red heat; the most stable oxide of iridium is the dioxide, and the tendency of iridium is always to form compounds which correspond to this oxide in composition, but compounds corresponding to the oxide Rh 2 O 3 are more often produced from rhodium, this being apparently its most stable oxide. The oxides of rhodium are reduced when heated in an atmosphere of hydrogen, but the oxides of iridium are reduced at ordinary temperatures. The difference between the two metals is therefore similar to that between ruthenium and osmium, which they much resemble in their general behaviour, although neither furnishes a tetroxide, and rhodium forms only one stable chloride, the hydroxide Rh(OH) 4 being dissolved by hydrochloric acid with evolution of chlorine. Iridium tetrachloride forms double chlorides or iridiochlorides isomorphous with the corresponding platinochlorides. Like cobalt, rhodium and iridium form compounds which are not only analogous in composition to the ferricyanides, but also isomorphous with them. A number of rhodium salts corresponding to the oxide Rh 2 O 3 are known, but they crystallize with difficulty; the hydroxides of iridium are soluble in acids, but no definite oxygen salts of iridium have been obtained. The hydroxides Rh 2 (OH) 6 and Ir 0 (OH) β manifest feeble acid properties, dissolving in alkalies. Rhodium and iridium are precipitated from solutions of their salts by most metals, but not by silver or gold.
In many cf their properties palladium and platinum more nearly resemble silver and gold than the preceding metals, just as nickel is more closely allied to copper than to cobalt and iron. Palladium is superficially oxidized at a lower temperature than silver, but its oxide is decomposed at a red heat; platinum is not in the least affected when heated in air or oxygen. The most stable compounds of palladium are the palladious compounds, a solution of palladic chloride, PdCl 4 , being decomposed when heated, with evolution of chlorine, Pd 2 Cl 4 being produced; the most stable compounds of platinum, however, are those which correspond in composition to platinic chloride. Palladium, like silver, is stained by iodine, which is without action on platinum, owing to the formation of palladious iodide, which is also obtained on adding potassium iodide to a solution of a palladious salt; it is a black powder insoluble in. water but soluble in ammonia, and also, like silver iodide, in a solution of potassium iodide. Both platinous and platinic chlorides, like auric chloride, readily unite with hydrochloric acid, forming chloroplatinous and chloroplatinic acids, H 2 PtCl 4 and H 2 PtCl 6 ; salts of these acids are produced by the combination of platinous and platinic chlorides with the chlorides of other metals. Palladious and palladic chlorides, and the corresponding chlorides of the other platinum metals, as already pointed out, form similar compounds, those which correspond in composition being in all cases isomorphous. Although platiuum and palladium do not furnish compounds analogous to the ferro-or ferri-cyanides, they form a distinct series of very beautiful and highly stable salts termed platinocyanides and palladiocyanides, which are analogous in composition to the chloroplatinites, and from which the corresponding acids may be obtained by decomposing the silver salts, for example, with hydrogen sulphide; thus :—
K 2 PtCl 4 K 2 Pt(CN) 4 K 2 Pd(CN) 4
Potassium pïatinochloride. Potassium platinocyanide. Potassium palla<liocyanlde.
The palladiocyanides are isomorphous with the corresponding platinocyanides. These salts combine with the halogens, forming compounds such as potassium platini-chlorocyanide, K 2 PtCl 2 (CN) 4 , which are analogous to the chloroplatinates such as K 2 PtCl 6 . A somewhat similar series of .bodies may be obtained from gold, e.g., potassium auricyanide, KAu(CN) 4 , and potassium aurichlorocyanide, KAuC1 2 (CN) 2 .
The oxides and hydroxides of platinum and palladium are somewhat more stable than the silver and gold compounds, but are entirely decomposed on ignition; they dissolve in acids, but the resulting salts have little stability, although they are considerably more stable than those of gold. The hydroxides appear all to possess acid properties, platinic hydroxide dissolving readily in solutions of alkalies, forming metallic derivatives which may be obtained in crystals. Platinic hydroxide, like auric hydroxide, is converted into a fulminating compound by the action of ammonia.
Cobalt and the platinum metak are especially characterized by the readiness with which they enter into reaction with ammonia, forming salts of cobalt and platinum bases, which are often extremely complex in composition. Thus, a solution of platinous chloride in warm, moderately strong ammonia, concentrated by evaporation, on cooling deposits yellow crystals of platinodiammonium chloride, PtN 4 H 12 Cl 2 , H 2 O; by treating a solution of this salt with silver sulphate it is converted into the sulphate PtN 4 H 12 (SO 4 ), from which the corresponding hydroxide PtN 4 H 12 (OH) 2 may be procured by adding to the solution the equivalent quantity of barium hydroxide. Platinodiammonium hydroxide may be obtained in crystals; it is a powerfully alkaline, caustic substance, aud, according to Thomsen, its heat of neutralization is as great as that of the alkalies and the hydroxides of barium, strontium, and calcium. When platinodiammonium hydroxide is heated as long as it gives off ammonia and water, it is converted into platinoammonium oxide PtN 2 H 6 O, from which salts such as platinoammonium chloride, PtN 2 H 6 Cl 2 , may be prepared by the action of acids; on passing chlorine into water in which the latter is suspended it is converted into chloroplatinoammonium chloride, PtCl 2 N 2 H 6 Cl 2 .
The relation of these compounds to platinous chloride and to one another may be expressed by the following formulae :— [table]
Similarly, when a solution of cobaltous chloride in ammonia is exposed to the air, the liquid assumes first a brown and then a red colour, and if the reddened liquid is mixed with hydrochloric acid a brick-red precipitate is produced; by recrystallizing this from hot water containing hydrochloric acid it is obtained in beautiful ruby-red crystals of the composition Co 2 N 10 H 30 Cl 6 . The relation of this compound, which is termed purpureocobaltic chloride, to cobaltous and cobaltic chlorides may be expressed as follows :— [table]
Nickel and copper appear to form similar compounds, but they are mostly very soluble and far less complex, these elements having little or no tendency to function as hexads.
Titanium — Zirconium — Tin. [table]
Titanium and zirconium are closely related to silicon and, therefore, to carbon, and occupy an intermediate position between these elements and the true metals. Titanium and tin, however, are more nearly related than either titanium and zirconium or tin and zirconium,— zirconium more closely resembling silicon, especially in tho properties of its oxide.
Titanium and zirconium apparently resemble silicon closely in their physical properties, and exist both in the amorphous and crystalline state. In the amorphous state they burn with great brilliancy when heated to redness in oxygen, and they take fire when heated to redness in chlorine; titanium also unites directly with nitrogen at high temperatures, and it is stated forms the three nitrides Ti 3 N 2 , Ti 3 N 4 , TiN 2 . Titanium is said to be soluble in warm hydrochloric acid, but zirconium is only slowly attacked by ordinary acids, although it readily dissolves even in cold hydrofluoric acid. Crystalline zirconium is very hard, and resembles antimony in colour, lustre, and brittleness; it is less fusible than silicon.
Tin is a soft, very malleable, white metal, with a tinge of yellow, and is a tolerably good conductor of electricity, titanium and zirconium being non-conductors of electricity. It fuses at about 230° C. It scarcely tarnishes at ordinary temperatures, but at a red heat it rapidly oxidizes, and readily decomposes water. It is only slowly dissolved by warm hydrochloric acid, which converts it into stannous chloride, Sn 2 Cl 4 , but is rapidly acted on by moderately concentrated nitric acid; chlorine acts readily on the melted metal, converting it into stannic chloride, SnCl 4 . Two chlorides of titanium are known, Ti 2 Cl 6 and TiCl 4 , but only one chloride of zirconium, ZrCl 4 . Titanium and tin tetrachlorides are colourless liquids, which boil respectively at 135° and 115° C.; their lower chlorides are solids, and are powerful reducing agents. Zirconium chloride is a white, crystalline solid; its vapour density corresponds with the formula ZrCl 4 . When titanium and tin tetrachlorides are mixed with a small quantity of water, much heat is developed, and they dissolve, probably forming chlorhydrins; but a large quantity of water causes their conversion into the corresponding hydroxides, which separate as white gelatinous precipitates. Zirconium chloride is dissolved even by a large quantity of water, but the hydroxide Zr(OH) 4 separates as a gelatinous, bulky, white precipitate on the addition of alkalies.
Titanium, zirconium, and tin tetrafluoride unite with metallic fluorides, forming crystalline salts which correspond in composition to the silcofluorides or fluosilicates (p. 522), with which they are isomorphous; thus :— [table]
Titanic oxide, TiO 2 , has a reddish-brown colour; stannic oxide, SnO 2 , is white or yellowish-white, but assumes a transient dark yellow or brown colour when heated; zirconic oxide, ZrO 2 , is white. They may be obtained in crystals which are isomorphous. Zirconic oxide is infusible, and titanic oxide melts only at the heat of the oxyhydrogen blowpipe; stannic oxide is somewhat more fusible, although very difficult to fuse. Titanic oxide is reduced by charcoal at a white heat, and stannic oxide at a full red heat; zirconic oxide, however, appears entirely to resist reduction; by ignition with charcoal in an atmosphere of chlorine they are converted into the corresponding chlorides. After ignition they are insoluble in acids, except concentrated sulphuric acid; they are dissolved on fusion with alkalies. By dissolving the oxides or the corresponding hydroxides in acids, salts such as stannic sulphate, Sn(SO 4 ) 2 , are produced; the titanium and tin salts are so unstable, however, that their solutions are decomposed by boiling; the zirconium salts are somewhat more stable. Stannic hydroxide readily dissolves in alkaline solutions, but titanic and zirconic hydroxides are insoluble. The salts which are formed from the hydroxides or oxides by the action of alkalies are mostly analogous to the silicates of the form M√SiOo; thus :— [table]
When tin is heated with concentrated nitric acid it is converted into white, insoluble, metastannic acid; the air-dry substance has the composition Sn 5 O 10 ,10H 2 O, but on drying at 100° C. it is converted into the compound H 10 Sn 5 O 15 ; metastannic acid has feeble acid properties, and furnishes salts which all correspond in composition to potassium metastannate, K 2 H 8 S 5 O 15 . On boiling a solution of titanic hydroxide in hydrochloric acid an apparently similar compound is precipitated.
Stannous hydroxide, Sn 2 (OH) 4 , like stannic hydroxide, exhibits both basic and acid, properties, dissolving in acids and also in alkalies; it gradually absorbs oxygen on exposure to the air. When a solution of titanic hydroxide in hydrochloric acid is digested with metallic copper, it becomes violet-blue, and on the addition of an alkali a dark brown precipitate separates, which probably is the hydroxide corresponding to the chloride Ti 2 Cl 6 . Tin is the only element of this group which may be precipitated as sulphide by hydrogen sulphide.
V an adi um — Arsenic —N iobium — Antimony — Tantalum — Bismuth . [table]
These elements belong to the same group as nitrogen and phosphorus, but the relation between them is in a marked degree periodic, the resemblance being greatest in many respects between arsenic, antimony, and bismuth, and between vanadium, niobium, and tantalum. Vanadium is the only member of the group, excepting niobium and, perhaps, tantalum, which furnishes coloured compounds, and in this respect resembles titanium and chromium, to both of which it is about equally related. Relationship may also be traced between arsenic aud selenium; between niobium and zirconium, and niobium and molybdenum; between antimony and tin, and antimony and tellurium; and between tantalum and tungsten.
The following chlorides and oxides of the elements of this group are known :— [table]
Arsenic, antimony, and bismuth are readily reduced from their oxides, but the oxides of vanadium, niobium, and tantalum are not reduced to the metallic state by ignition with chareoal; vanadium may be obtained by the reduction of its dichloride in hydrogen at a bright-red heat. Arsenic has a brilliant, dark steel-grey lustre; antimony has a bluish-white, and bismuth a reddish-white colour; they are hard, and extremely brittle, and bad conductors of electricity. Arsenic begins to volatilize at 180° C. without fusing; antimony fuses at about 450° C., and bismuth at about 265° C. They are but little altered on exposure to the air at ordinary temperatures, but readily oxidize when heated. They are readily acted upon by moderately concentrated nitric acid, but are only very slowly dissolved by boiling hydrochloric acid.
Vanadium has been obtained as a whitish-grey powder which under the microscope appears as a brilliant, crystal· line, metallic mass possessing a silver-white lustre. It does not oxidize at ordinary temperatures, but when heated in air it rapidly absorbs oxygen, and is finally converted into the pentoxide. It is not attacked by boiling hydrochloric acid, but nitric acid of all strengths readily dissolves it. It is converted into a mononitride when heated in nitrogen gas.
Niobium and tantalum have been obtained only in an impure state as black powders. Like vanadium, these elements have a great tendency to unite with nitrogen. Vanadium dichloride crystallizes in apple-green coloured plates, and vanadium trichloride forms peach-blossom coloured tables; vanadium tetrachloride is a reddish-brown liquid, which boils at 154° C. An oxychloride of vanadium, VOCl 3 , corresponding to phosphorus oxychloride, POCl 3 , is also known; it is a yellow liquid, which boils at 126° C. The vapour density of the tetrachloride, according to Roscoe, corresponds with the formula VCl 4 , but as it is resolved into the trichloride and free chlorine on boiling, and even undergoes the same decomposition at ordinary temperatures, it may be doubted whether the density observed was really that of the tetrachloride, and not that of a mixture of a lower chloride and chlorine, especially as from the analogy which vanadium presents to chromium the formula V 2 Cl 8 appears to be more probable than the simpler formula.
The vanadium chlorides are dissolved and decomposed by water, the tetrachloride forming a blue, the trichloride a green, the dichloride a lavender-coloured, and the oxychloride a yellow liquid; the solution of the dichloride is a most powerful reducing agent, and the solutions of the tri-and tetra-chloride also exhibit reducing properties, owing to the tendency of the dissolved compounds to combine with oxgyen to form vanadic acid.
Vanadium pentoxide or vanadic anhydride, V 2 O 5 , corresponds in composition and properties to phosphoric anhydride; it has a brownish-red colour, and fuses at a red heat; it dissolves sparingly in water, forming a yellow, strongly acid solution of vanadic acid. By the combination of vanadic anhydride with basic oxides in various proportions salts corresponding to the three series of phosphates, as well as salts of a more complex character, are obtained: thus :—- [table]
Isomorphism has been observed between several corresponding phosphates and vanadates. The tetroxide and trioxide of vanadium also possess feeble acid properties and unite with basic oxides. The vanadium oxides dissolve in acids, mostly forming salts such as the sulphate V 2 O 2 (SO 4 ) 2 , in which the hydrogen of the acid is displaced by the radicle VO; the formation of salts of this kind is characteristic of all the metalloids included in this group, and apparently of the metalloids generally. A solution of vanadium pentoxide in sulphuric acid, which is red, becomes blue on treatment with sulphurous acid or hydrogen sulphide, the salt of the pentoxide being reduced to a salt of the tetroxide; by the aid of magnesium the reduction may be carried a stage further, and a salt of the trioxide is formed, the solution becoming green; and if zinc, cadmium, dγ sodium amalgam is employed, the solution finally becomes lavender-coloured, and contains a salt of the dioxide.
Arsenic trichloride, AsC1 3 , the only known chloride of arsenic, is a heavy, colourless, oily liquid, which boils at 132° C.; it is dissolved by water, probably as a chlorhydrin, as its solution in the smallest possible quantity of water deposits after some days crystals of the composition AsCl(OH) 2 .
Antimony trichloride is a crystalline compound which boils at 223° C.; its vapour density corresponds with the formula SbCl 3 . Antimony pentachloride, SbCl 5 , is a heavy, yellowish liquid, which cannot be distilled without decomposition. These chlorides are dissolved by hydrochloric acid or a small quantity of water, but a large quantity of water converts the trichloride into an insoluble oxychloride, SbOCl, and the pentachloride into a compound which corresponds in composition to pyrophosphoric acid, viz., H 4 Sb 2 O 7 . No oxychloride of arsenic or antimony corresponding to phosphorus oxychloride is known, but a. crystalline antimony sulphochloride, SbSCl 3 , exists.
Bismuth trichloride is a very fusible, volatile, deliquescent, colourless compound; its vapour density corresponds with the formula BiCl 3 ; water converts it into an insoluble oxychloride, BiOCl; when heated in hydrogen it furnishes an unstable dichloride, Bi 2 Cl 4 .
Niobium pentachloride, NbCl 5 , is a yellow, easily fusible,, crystalline substance, which boils at 240° C.; niobium oxychloride, NbOCl 3 , is a white, infusible, but volatile, crystalline compound. The formulae of these .bodies have: been established by the determination of their vapour densities; both are decomposed by water. The solution of niobium chloride in hydrochloric acid becomes blue when a. plate of zinc is immersed in it. Tantalum pentachloride is a yellow, fusible, crystalline substance, which boils at about 242° C.; its vapour density corresponds with the formula TaCl 5 ; it is decomposed by water.
Both oxides of arsenic are white; the trioxide sublimes at about 200° C. without fusing; the pentoxide, when suddenly heated to redness, fuses and is decomposed into the trioxide and oxygen. The vapour density of the tri oxide corresponds with the formula As 4 O 6 . Arsenic trioxide is sparingly soluble in cold water, more readily in hot water, and freely in hot hydrochloric acid; the aqueous solution reddens litmus. The pentoxide is deliquescent and easily soluble in water, forming a solution of arsenic acid, H 3 AsO 4 ; on evaporating the solution, however, even at ordinary temperatures, crystals of an acid corresponding in composition to pyrophosphoric acid, H 4 As 2 O 7 i , are obtained, which is converted into the pentoxide when heated at about 260° C. Arsenic acid furnishes a series of salts analogous in composition to and isomorphous with the corresponding phosphates. Arsenic trioxide is almost destitute of basic properties, and the salts formed by its union with basic oxides are not very stable, so that it differs greatly from the corresponding oxide of phosphorus; thermochemical investigation confirms this conclusion. Thus, on inspection of the following table, which shows the amounts of heat which are developed on adding sodium hydroxide solution to solutions of the acids of phosphorus and arsenic, it will be evident that, whereas the behaviour of arsenic acid is similar to that of phosphoric acid, the behaviour of a solution of arsenic trioxide is altogether different from that of a solution of phosphorus trioxide (phosphorous acid), the molecule of arsenic trioxide regarded as As 2 O 3 , like boron oxide, evidently existing in solution as a dibasic acid; it is not known whether the molecular composition of phosphorus trioxide is correctly represented by the formula P 2 O 3 , and whether therefore the difference between the phosphorus and arsenic compounds is to be referred to the existence of a difference similar to that exhibited by the formulae P 2 O 3 and As 4 O 6 . [table]
It will be evident also on comparing the following table with that on p. 518 that much less heat is developed in the formation of the arsenic compounds than in the formation of the corresponding phosphorus compounds :— [table]
The oxides of antimony are white or yellowish; the trioxide is fusible and volatile, and is partly converted into the tetroxide when heated in air; the tetroxide is infusible, and unalterable by heat; the pentoxide is converted into the tetroxide at a red heat. The vapour density of the trioxide has not been determined, but as it may be obtained in crystals isomorphous with those of arsenic trioxide its formula is probably Sb 4 O 6 . The oxides of antimony all exhibit feeble acid properties, dissolving in alkalies, but the salts mostly have little stability; the most stable salts apparently are those which correspond to the potassium salt of the formula K 2 Sb 4 O 11 . The· oxides of antimony are soluble in hydrochloric and concentrated sulphuric acids; the trioxide may thus be converted into the sulphate, Sb 2 (SO 4 ) 3 , which is decomposed by water into an acid and a basic salt.
The bismuth oxides are yellow or brown, and are devoid of acid properties, but they dissolve in acids; the pentoxide is decomposed and converted into the tetroxide when heated to about 220° C., and the latter is reduced to the trioxide at a somewhat higher temperature. Bismuth sulphate, Bi 2 (SO 4 ) 3 , like antimony sulphate, is decomposed by water into an acid and a basic salt.
On boiling a solution of acid potassium tartrate, HKC 4 H 4 O o , with antimony trioxide the so-called tartar emetic (SbO)KC 4 H 4 O 6 , is formed, the oxygenated radicle SbO displacing an atom of hydrogen in the potassium tartrate; corresponding arsenic and bismuth compounds are formed in a similar manner.
Niobium and tantalum pentoxides are white and infusible; they do not volatilize; after ignition they are insoluble in all acids, but dissolve on fusion with alkalies. They exhibit marked acid properties, combining with basic oxides in various proportions. The hydrated oxides obtained on decomposing the chlorides by water are acid to litmus, and are soluble in hydrochloric and hydrofluoric acids. The fluorides of niobium and tantalum, and niobium oxyfluoride, readily form double salts with other fluorides, which in many cases are isomorphous with certain titanium, zirconium, tin, and tungsten compounds which do not exactly correspond to them in composition, but in which apparently fluorine and oxygen displace one another isomorphously; thus, the following pairs of compounds are isomorphous :— [table]
Arsenic and antimony, like phosphorus and nitrogen, form gaseous trihydrides, AsH 3 and SbH 3 , and arsenic is said to furnish also a solid dihydride, As 2 H 4 . Arsine and stibine, AsH 3 and SbH 3 , are formed whenever hydrogen is evolved in presence of an arsenic or antimony compound; they are extremely unstable compounds, especially the latter, and are entirely devoid of basic properties. The corresponding compounds of arsenic and antimony with positive hydrocarbon radicles, however, like the analogous nitrogen and phosphorus compounds, are far more stable and combine readily with other elements. Thus, triethylamine, N(C 2 H 5 ) 3 , and triethylphosphine, P(C 2 H 5 ) 3 , combine with acids, forming salts corresponding to the ammonium salts such as the chloride PH(C 2 H 5 ) 3 C1; the triethylphosphonium salts, moreover, are stable in presence of water, which at once decomposes phosphonium iodide. Triethylstibine, Sb(C 2 H 5 ) 3 , decomposes hydrochloric acid with evolution of hydrogen, forming Sb(C 2 H 5 ) 3 Cl 2 ; trimethylarsine, however, does not react with acids, but with chlorine forms the compound As(CH 3 ) 3 Cl 2 ,the existence of which furnishes a proof that arsenic may function as a pentad element. The phosphorus, antimony, and arsenic compounds absorb oxygen with great avidity, forming powerfully basic oxides such as PO(C 2 H 5 ) 3 . They also furnish hydroxides such as N(C 2 H 5 ) 4 .OH, P(C 2 H 5 ) 4 .OH, &c. (p. 575), which exhibit the closest resemblance to the alkalies; Thomsen, indeed, has shown that the heat of neutralization of the hydroxide N(CH 3 ) 4 .OH is equal to that of the alkalies. The existence of these hydroxides, and their behaviour, afford the strongest reason for believing that a solution of ammonia at least partially consists of ammonium hydroxide, NH 4 .OH (p. 510).
Molybdenum — Tungsten — Uranium. [table]
Molybdenum and tungsten are closely related to chromium. They may be obtained by reducing theiι oxides at a high temperature in a stream of hydrogen. Molybdenum, after fusion, is white, and has a silvery lustre; the powder obtained by reducing tungsten trioxide has a grey colour, and strong lustre. Both are difficult of fusion, and insoluble in most acids except nitric acid; when heated in air or oxygen they are finally converted into the trioxides MoO 3 and WO 3 . They form the following chlorides and oxides :—∙ [table]
Molybdenum pentachloride is the highest chloride obtained by heating molybdenum in chlorine gas; it is a black, crystalline substance, which forms a colourless solution when dissolved in a considerable quantity of water; when heated in air it is converted into a colourless oxychloride, MoO 2 Cl 2 ; according to Debray, its vapour density corresponds with the formula MoC1 5 . The remaining chlorides are not distinctly crystalline; the tetrachloride is brown, the trichloride red, and the dichloride yellow; the molecular formulae of these chlorides are undetermined. In an atmosphere of carbon dioxide the dichloride bears a bright red heat without melting or volatilizing, the trichloride is resolved into di-and tetrachloride, and the latter when again heated, splits up into pentachloride which sublimes, and trichloride which remains; the di-and tri-chlorides are insoluble, but the tetrachloride is readily soluble in water.
Tungsten is converted by chlorine at a moderate heat into a hexachloride, WCl 6 , which forms black-violet crystals; it melts at 275° C., and boils at 341° C.; from the determination of its vapour density it appears that it undergoes decomposition when heated. When pure it is insoluble in water below 60° C., but if the slightest trace of oxychloride be present it is decomposed by water even in the cold. Tungsten pentachloride crystallizes in black, highly deliquescent needles; it melts at 248° C., and boils at 376° C., and its vapour density at 350 β C. corresponds with the formula WCI 5 . It has already been pointed out that the existence of pentachlorides of tungsten and molybdenum, and of a vanadium tetrachloride, is apparently anomalous; at present, these chlorides, chlorine dioxide, ClO 2 , nitrogen monoxide and dioxide, and, perhaps, uranium pentachloride are the only instances known to us in which apparently an odd number of affinities are disengaged (see p. 473).
Tungsten tetrachloride is a crystalline, greyish-brown substance; when strongly heated it decomposes into the volatile pentachloride and the non-volatile dichloride, which is a loose grey powder, destitute of crystalline structure. Tungsten also forms two oxychlorides, WOCl 4 and WO 2 Cl 2 ; the former is obtained in beautiful ruby-red, and the latter in yellow crystals. The formula of the monoxychloride, WOC1 4 , has been established by determining its vapour density.
Molybdenum trioxide has marked acid properties, but the remaining oxides are feebly basic; both oxides of tungsten exhibit acid functions. Solutions of molybdous salts corresponding to the oxide Mo 2 O 3 are almost black, and solutions of molybdic salts corresponding to the oxide Mo 2 O 4 are reddish-yellow; the molybdates are mostly white. Molybdenum trioxide, MoO 3 , melts at a. red heat to a dark yellow liquid, which on cooling forms a strawyellow mass, breaking up into crystalline scales; it is scarcely soluble in water, but dissolves in hydrochloric, nitric, and sulphuric acids. Like silica and other feebly acid oxides it combines with basic oxides in a large number of different proportions, forming molybdates.. Tungsten trioxide, WO 3 , is a canary-yellow substance, insoluble in water and in most acids; it melts at a high temperature, and, like molybdenum trioxide, may be volatilized by heat; it dissolves in alkaline solutions, forming tungstates, of which a great variety may be obtained by the union of tungsten trioxide with basic oxides in various proportions. The tungstates are mostly colourless. When a hot solution of an alkaline tungstate is mixed with an acid a yellow precipitate of tungstic acid, H 2 WO 4 , is obtained, which is insoluble in water, and forms insoluble salts with all metals, except the alkali metals; but by decomposing tungstates by the stronger acids at ordinary temperatures a soluble variety of the acid, called metatungstic acid, is produced, which forms soluble salts with nearly all metals.
Uranium. — The atomic weight of uranium has not yet been ascertained, and its position among the elements is therefore uncertain, especially as it does not exhibit marked analogy to any other element. The generally received atomic weight is 120, but Mendeljeff has proposed to double this; L. Meyer, however, considers that from the high specific gravity of the metal it is more probable that its atomic weight is 180, and that it occupies a position in the series near to tungsten. A redetermination of the equivalent of this element and an investigation of its compounds are much required.
Uranium is like iron in colour; it tarnishes when exposed to the air, and in the pulverulent state takes fire at a comparatively low temperature, and burns with great brilliancy, forming a green oxide, U 3 O 8 ; it is dissolved by hydrochloric acid, forming a green solution, which has powerful reducing properties. The highest oxide has the composition UO 3 . Uranium forms a volatile stable tetrachloride, UC1 4 , which crystallizes in dark-green octahedra, and an unstable chloride regarded as the pentachloride, UC1 5 ; both are dissolved and decomposed by water.
When uranium, or either of its oxides, is dissolved in nitric acid, a solution of uranic nitrate is formed from which the salt may be crystallized in beautiful yellow prisms of the composition (UO 2 )(NO 3 ) 2 , 6H 2 O; on the addition of potassium hydroxide to a solution of uranic nitrate a yellow insoluble precipitate of potassium uranate of the composition K 2 U 2 O 7 is produced. An oxychloride, UO 2 Cl 2 , corresponding to the nitrate, is obtained on passing chlorine over an ignited mixture of charcoal with an oxide of uranium. In all these formulae, U = 240.
Lanthanum — Cerium — Didymium — Yttrium — Erbium — Thorium. [table]
Excepting thorium, these elements are closely related in properties; their compounds are of extremely rare occurrence. Lanthanum, cerium, and didymium are always associated together, and on account of their occurrence in the mineral cerite are usually termed cerite metals; they appear to occupy a position intermediate between the alkaline earth metals and the metals of the aluminium group. The atomic weights of yttrium and erbium, the so-called gadolinite metals, have not been determined; their compounds much resemble those of the cerite metals. The atomic weight of thorium is also unknown; from the properties of its compounds, however, it is not improbable that it is a member of the aluminium group.
Lanthanum, cerium, and didymium have been obtained by the electrolysis of their chlorides. Cerium is an extremely ductile metal, resembling iron in colour and lustre; it melts more readily than silver, but much less readily than antimony. It retains its lustre for a considerable period in dry air, but takes fire much more readily than magnesium, and burns with great brilliancy. It slowly decomposes cold water, and rapidly dissolves in hydrochloric acid.
Lanthanum and didymium closely resemble cerium in their general chemical behaviour. Lanthanum is a slightly harder, less ductile metal than cerium, and appears to melt nearly at the same temperature. It readily oxidizes superficially even in dry air, although it takes fire at a much higher temperature than cerium. Didymium more closely resembles lanthanum than cerium, but has a distinct yellowish colour.
Only one oxide of lanthanum is known, namely, La 2 O 3 ; it is white, and unalterable by heat, and combines directly with water, forming the hydroxide La 2 (OH) 6 , which is a soft, white powder resembling calcium hydroxide. The lanthanum salts of colourless acids are all colourless.
Cerium forms the two oxides Ce 2 O 3 and Ce 2 O 4 . Cerons oxide, Ce 2 O 3 , has a(?) greyish-blue colour, and on exposure to the air is converted into yellow ceric oxide, Ce 2 O 4 . Cerous hydroxide, Ce 2 (OH) β , precipitated from the solution of a cerous salt by an alkali, is white; on exposure to the air it is converted into a mixture of cerous carbonate and ceric hydroxide; the latter is also obtained on passing chlorine into water in which cerous hydroxide is suspended. Ceric hydroxide dissolves in hydrochloric acid with evolution of chlorine, forming a solution of cerous chloride, Ce 2 Cl 6 , but it forms with sulphuric acid a brown-red solution of ceric sulphate, which is a powerful oxidizing agent. Cerous salts in solution are colourless, but a few possess a pale rose colour in the solid state.
Two oxides of didymium corresponding to the two oxides of cerium are known. The lower oxide, Di 2 O 3 , is white, and in contact with water is slowly converted into the hydroxide, Di 2 (OH) 6 ; didymium hydroxide resembles aluminium hydroxide in appearance, but has a pale rose colour, and the salts formed by dissolving it in acids have either a pure rose or violet colour. By gently heating the oxide Di 2 O 3 in air it is converted into the higher oxide Di 2 O 4 , which has a dark brown colour; it dissolves in hydrochloric acid with evolution of chlorine, and in oxyacids with evolution of oxygen, forming the same salts as are produced on dissolving the lower oxide.
Yttrium and erbium each form only one oxide, and like lanthanum, cerium, and didymium only one corresponding chloride. Yttrium oxide, Y 2 O 3 , is white, and erbium oxide, Er 2 O 3 , has a pale rose colour; the former is readily soluble in acids, even after ignition, but the latter dissolves with great difficulty. The yttrium salts are colourless; but the salts of erbium have a beautiful rose colour, and their solutions give an absorption spectrum, which is not the case with yttrium salts. When erbium oxide is strongly heated it glows with an intense green light, which in the spectroscope exhibits a continuous spectrum intersected by a number of bright bands corresponding in position to the dark bands in the absorption spectrum of solutions of erbium salts. Didymium oxide behaves similarly, and solutions of didymium salts give an absorption spectrum by which they are readily distinguished from erbium salts.
The chlorides, nitrates, and sulphates of the cerite and gadolinite metals are readily soluble in water, but their carbonates are insoluble; the solutions of their salts all possess a sweet astringent taste. The most conclusive evidence that lanthanum, cerium, didymium, and yttrium are closely allied to the alkaline earth metals is afforded by thermochemical investigation, the amounts of heat developed on neutralizing their hydroxides being for hydrochloric acid, according to Thomsen’s experiments, inferior only to that developed on neutralizing the hydroxides of the alkali and alkaline earth metals, as shown by the following examples :— [table]
Metallic thorium has been obtained as a dark-grey powder of the specific gravity 7*65 to 7 - 79, easily soluble in nitric acid, but difficultly soluble in hydrochloric acid. Its oxide, Th 2 O 3 (if Th = 178’5), is white, and after ignition is insoluble in all acids except concentrated sulphuric acid; it appears to be destitute of acid properties as it does not expel carbon dioxide when fused with alkaline carbonates. Thorium hydroxide is precipitated from solutions of thorium salts by alkalies as a gelatinous mass, soluble in most acids, but insoluble in alkalies. Thorium chloride, Th 2 Cl 6 , is a white crystalline substance, which dissolves in water with rise of temperature. Thorium sulphate, Th 2 (SO 4 ) 3 , is crystalline, and, like the sulphates of the cerite and gadolinite metals, is more soluble in cold than in hot water.
Periodic Relations of the Elements.
The foregoing description of the elements and of some of their more important compounds will be sufficient to show that, while each element manifests certain characters which distinguish it from all others, many of the elements are more or less closely related in properties, as indeed we have already frequently had occasion to point out.
Elements which exhibit similar properties often differ in atomic weight to the same or nearly the same extent ; for example, the difference between the atomic weights of potassium and rubidium is about 46, and of rubidium and caesium 47'5, since K = 39, Rb = 85 ’ 2, and Cs=132 ’ 7. As these three elements are closely related in properties, rubidium differing from potassium to about the same extent that caesium differs from rubidium, we are led to suspect a connection between the atomic weight of an element and its properties, especially as with very few exceptions the elements all possess different atomic weights. Many instances of relation between atomic weight and properties, similar in character to that which obtains in the case of potassium, rubidium, and caesium, have been commented on from time to time by various chemists, but the connection between the atomic weights of the elements generally and their properties was not recognized until Mendeljeff in 1869 pointed out that the latter are periodic functions of the former. In other words, if the elements are grouped in the order of their atomic weights, it will be found that nearly the same properties recur periodically throughout the entire series. Hence the whole of the elements may be arranged in a number of groups, each group consisting of members of the same natural family following each other in the same order. The elements are arranged in this manner in the following table, although, in order to retain elements which are undoubtedly members of the same natural family in the same vertical series, a few departures from the order of atomic weights are necessary, but probably they are necessary merely because the atomic weights are incorrectly determined; thus, tellurium is placed before iodine, and osmium, iridium, and platinum before gold. If the position assigned to uranium be correct, the number at present accepted as its atomic weight is much too low. [table]
The position of hydrogen at the head of a series of metals is in accordance with Graham’s conclusion that this element possesses the characters of a metal, a conclusion which many chemists are inclined to accept on account of the striking analogies to the metals which it exhibits in its relation to the halogens, oxygen, &c. As the properties of alloys are usually similar to those of their constituent metals, whereas the compounds of metals with non-metallic elements are in most cases widely different from those of the elements which enter into their composition, the fact that Troost and Hautefeuille’s sodium and potassium hydrides(lsodium-hydrogen and potassium-hydrogen alloys) (p. 524) retain the metallic character of sodium and potassium affords evidence confirmatory of the assumption that hydrogen has metallic properties.
We leave it to the reader to trace out the alteration in properties which accompanies alteration in atomic weight; but we may point out as an instance that the power to combine with other elements is subject to periodic variation; thus, the composition of the highest chlorides and oxides of the elements in the first and second horizontal series in the above table is as follows :— [table]
The chlorides and oxides of the elements of each of the following horizontal series exhibit similar differences in composition. With very few exceptions, only the elements which are included in the same vertical series as carbon, nitrogen, oxygen, and fluorine form compounds with hydrogen, the stability of which diminishes with increase of atomic weight; but the composition of these hydrides is subject to periodic variation, as is shown by the following list :— [table]
The periodic character of the relation between the properties of the elements and their atomic weights is especially evident in their physical properties, which are more readily compared than the so-called chemical properties, as they admit of measurement, although, on account of our imperfect knowledge, comparison is at present possible only to a very limited extent. One of the few physical properties which has been determined for most of the elements is the specific gravity in the solid state; now, on comparing the specific gravities of the various elements we find that they do not increase progressively with increase of atomic weight, but that they increase and diminish periodically. Hence, as Lothar Meyer has shown, on taking the atomic volumes, or quotients of the atomic weights of the elements divided by their specific gravities in the solid state, as abscissae, and the atomic weights as ordinates, a curve is obtained which exhibits a series of maxima and minima, viz., five maxima and five minima in the portion which includes barium and the elements of lower atomic weight. The most electro-positive elements, lithium, sodium, potassium, rubidium, and caesium form the five maxima; but with increase of atomic weight the height to which the curve rises also rapidly increases, namely, in the proportion
Li: Na: K: Rb: Cs = 12 :24 :46 :57 :79.
The remaining physical properties of the elements, so far as they are known to us, appear to be subject to similar periodic variation, but for a full discussion of the connection between the atomic weights of the elements and their properties the reader is referred to L. Meyer’s work, Die modernen Theorien der Chemie (Breslau).
The establishment of the periodic law may truly be said to mark an era in chemical science, and we may anticipate that its application and extension will be fraught with the most important consequences. It reminds us how important above all things is the correct determination of the fundamental constants of our science—the atomic weights of the elements, about which in many cases great uncertainty prevails; it is much to be desired that this may not long remain the case. It also affords the strongest encouragement to the chemist to persevere in the search for new elements. (∏, e . a .)
ORGANIC CHEMISTRY.
As has been already explained (p. 520) it was at one time thought that certain chemical compounds were producible only through the agency of living things, and the name “organic chemistry” was in consequence conferred upon this branch of the science. The progress of discovery has, however, served to break down the barrier that was supposed to exist between those and purely inorganic compounds, since it has been found possible to build up artificially a very large number of compounds formerly regarded as essentially organic. Nor must it be forgotten that many living things elaborate compounds which can in no sense be regarded as organic—such as the silica of grasses, the carbonate of calcium secreted by molluscs, &c., and the tricalcic phosphate forming the bones of the higher animals.
Although the number of elements entering into the composition of organic compounds is comparatively small, yet these compounds far exceed inorganic bodies both in number and complexity of composition. It is indeed the vast and ever-increasing numbers of the organic compounds that render it necessary to form a separate branch of the science for their study, and not any real chemical difference between the matter forming these and mineral compounds.
Definition and Character of Organic Compounds. — Carbon is an invariable constituent of organic bodies, so that this branch of the science is sometimes defined as “the chemistry of the carbon compounds.” The best chemical definition of organic compounds, however, is that proposed by Frankland, viz., “compounds the molecules of which consist of one or more atoms of carbon directly combined either with carbon, nitrogen, or hydrogen.”
After carbon, the elements of most frequent occurrence in organic compounds are hydrogen, oxygen, and nitrogen. Sulphur, phosphorus, the halogens, silicon, boron, and the metals are of comparatively rare occurrence. The number of atoms entering into the composition of organic molecules is often very great—far exceeding the atoms of the most complex molecule of a mineral substance. Thus, stearin contains 173 atoms, and albumin no less than 226. Chemical bodies that are possessed of great complexity of composition are generally distinguished by the facility with which they split up under the influence of the physical forces and chemical reagents; thus organic compounds are as a class characterized by their instability.
Ultimate Analysis of Organic Compounds.
Before the formula of an organic compound can be determined, it is necessary to ascertain the amounts of its several constituents—that is to say, to make a quantitative analysis of it. The elements carbon and hydrogen being of paramount importance are generally first determined. The method employed depends upon the fact that all organic compounds undergo combustion when heated with easily reducible oxidizing substances, their carbon being oxidized to carbon dioxide and their hydrogen to water. These two products of combustion being collected in suitable apparatus and weighed, the necessary data for calculating the amounts of carbon and hydrogen are obtained. The following is a brief description of the process.
Determination of Carbon and Hydrogen. — A tube of Bohemian hard glass of about 10 or 14 millimetres internal diameter, after being scrupulously cleaned and dried, is drawn off in the blo¼pipe flame so as to leave a tailed extremity of the form shown in figure 1. The open end of the tube is then held in the flame of the blowpipe till the sharp edge of the glass is rounded off. This tube, known as the combustion tube, is from 500 to 800 millimetres in length, according as the substance contains little or much carbon.
The substance generally employed to effect the combustion of carbon compounds is cupric oxide (CuO), and as this oxide is of a somewhat hygroscopic nature, it is necessary to dry it thoroughly before use, and then exclude it carefully from the air. Should these precautions not be observed, the oxide absorbs water from the air, and when heated in the combustion tube this water distils over and is weighed with that produced by the combustion of the substance. It is therefore found convenient in practice to have anothér glass tube sealed at one end and closely corked at the other for the reception of the cupric oxide. This tube is of sufficient capacity to contain enough cupric oxide to fill the combustion tube, and its diameter is such as to allow it to be introduced into the open end of the combustion tube.
The substance for analysis, after being finely powdered, is dried in a water oven at 100° C. till it ceases to lose weight, or, if decomposed by heat, in a vacuum over some waterabsorbing substance, such as strong sulphuric acid or calcium chloride. When dry it is transferred to a small stoppered bottle or corked tube, sealed at one end, and of about 40 or 50 millimetres in length.
The apparatus for absorbing the carbon dioxide is a series of light glass bulbs (figs. 2 and 3) containing a solution of caustic potash (KHO) of sp. gr. 1∙27, through which the gases evolved during the combustion are obliged to pass. It is customary to attach to one end of the potash bulbs a small glass tube drawn to a point at its free end, and containing a small piece of solid potash enclosed between plugs of cotton wool. This appendage is for the purpose of retaining any trace of carbon dioxide or moisture carried over from the bulbs by the escaping gases.
The water formed by the combustion is collected in a tube (fig. 4) filled with fragments l 'of dried calcium chloride or pumice stone wetted with strong sulphuric acid,
The combustion tube when charged is heated in a furnace which consists essentially of a row of gas burners constructed so as to burn a mixture of coal gas and air, and thus obtain a non-luminous but very hot flame. Two forms of furnace are in common use in laboratories. In Hofmann’s furnace a clay cylinder perforated by a large number of very fine holes is fitted over the end of each gas burner, thus dividing the escaping gas into a number of fine streams, and so ensuring its perfect combustion. In Griffin’s furnace the burners are constructed on Bunsen’s principle, that is, each burner is perforated by holes at its lower end so that a mixture of gas and air escapes from its upper extremity. The combustion tube is supported on a thin iron trough lined with asbestos, the whole being enclosed by unglazed tiles.
The cupric oxide intended for use is first heated to redness in a crucible, and transferred while still hot to its tube, wherein it is allowed to cool. When sufficiently cool the cork is withdrawn, and enough cupric oxide poured into the combustion tube to occupy about ½th of its length. The tube or bottle containing the dried pulverized substance having been weighed, about half a gramme is then introduced into the combustion tube, and a small quantity of cupric oxide is poured down after it. The bottle being again weighed gives the exact weight of the substance employed. The substance in the combustion tube is thoroughly mixed with cupric oxide by stirring it about with a long wire with corkscrew point.
More cupric oxide is now added till the tube is filled nearly to its open end, and a plug of freshly ignited asbestos is finally introduced. The tube is then gently tapped while in a horizontal position so as to cause a slight subsidence of the contents, thus allowing a free passage for gas throughout its whole length. The calcium chloride tube is fitted into the open end of the combustion tube by means of a tightly fitting cork or caoutchouc stopper, and the potash apparatus is attached to the calcium chloride tube by means of a short piece of caoutchouc tubing bound tightly round with thin copper wire. Both the potash bulbs and calcium chloride tube are carefully weighed before being attached. The whole apparatus as arranged for the combustion is shown in fig. 5 :—
The front part of the tube is first heated to redness, and the gas burners are then gradually turned on so as to cause the slow combustion of the substance. When no more bubbles pass into the potash solution the gas burners are turned out, the point of the tail of the combustion tube is broken off, and dry air passed through the whole system t© sweep out the last traces of carbon dioxide and water vapour lingering in the tube. In cases where great accuracy is required it is desirable to sweep out the combustion tube first with a current of pure oxygen and then with air. This ensures the complete oxidation of any trace of carbon that may have escaped the first combustion.
When the operation is completed the potash bulbs and calcium chloride tube are allowed to cool down to the atmospheric temperature and then again weighed. The increase in weight gives the respective amounts of carbon dioxide and water produced. Since carbon dioxide contains in 11 parts by weight 3 of carbon (CO 2 = 12 + 2×16 = 44), ⅛ of the weight of the CO 2 obtained is due to carbon. Similarly ½ of the weight of the water found is due to hydrogen (H 2 O = 2x1 + 16 = 18). It is customary to express the results in parts per cent. In practice the hydrogen usually comes out a little too high, and the carbon a little too low.
Cupric oxide is sometimes replaced by lead chromate in cases where the substance is difficultly combustible, and more particularly when the compound contains chlorine, bromine, or sulphur. In these latter cases were cupric oxide employed, cupric chloride or bromide would be formed, and might volatilize over into the calcium chloride tube, thus unduly increasing its weight. Sulphur in presence of cupric oxide is oxidized to SO 2 , and this gas if allowed to pass into the potash bulbs would be absorbed. These sources of error are avoided by the employment of lead chromate, because non-volatile lead chloride or bromide are formed, and sulphur is oxidized to SO 3 , which is fixed as PbSO 4 . When substances rich in nitrogen are burnt with cupric oxide, nitrogen dioxide is likely to be formed, and this on meeting the air and potash in the bulbs is absorbed, and increases the weight of the CO 2 apparatus. To obviate this source of error it is necessary when dealing with nitrogenous bodies to plug the end of the combustion tube with a roll of freshly reduced copper gauze. This being kept at a bright red heat during the entire operation decomposes the oxides of nitrogen, and retaining the oxygen, allows only nitrogen to escape. 1
When the substance to be analyzed is a liquid, a known weight is sealed up in a small glass bulb (fig. 6).
After sealing and weighing (the weight of the empty bulb having been previously determined) the neck of the bulb is broken off, and the broken portion, together with the bulb itself, introduced into the combustion tube, which is afterwards filled up with cupric oxide, <fcc., in the usual manner. In burning liquids the anterior portion of the combustion tube is, as with solids, first heated to redness; the portion containing the bulb is then gradually warmed so as to expel the liquid, which is thus made to distil slowly over the red-hot cupric oxide which effects its combustion.
The analysis of gaseous organic compounds is effected in eudiometers, or in special apparatus, of which several forms have been devised. (See Bunsen’s Gasometry, Sutton’s Volumetric Analysis, Thorpe’s Quantitative Analysis, &c.)
Determination of Nitrogen. — Two methods are in use for determining this element. Will and Varrentrapp’s method depends upon the fact that many nitrogenous bodies when heated with caustic alkalies yield their nitrogen in the form of ammonia. A known weight of the substance is intimately mixed with soda-lime, 2 and then heated in a combustion tube, the ammonia being absorbed by dilute hydrochloric acid contained in a glass apparatus attached to the end of the tube. (See fig. 7.)
In conducting an operation of this kind some soda-lime is first introduced into the tube; the substance is then mixed intimately with more sodalime in a mortar, and the mixture transferred to the tube; the mortar is next rinsed out with more soda-lime, and the rinsings are added to the contents of the tube, which is finally filled up with pure soda-lime. When charged the tube is tapped so as to secure a free passage for the escaping gas
(as in the combustion for the determination of C and H), and after it is laid in the trough of the gas furnace, the acid bulb is attached to it by means of a tightly fitting cork. The combustion is performed in precisely the same manner as in the process of burning with cupric oxide; when it is completed the tail of the tube is broken off, and air drawn through the apparatus so as to draw the last traces of ammonia into the acid. If the substance contains much nitrogen it is advisable to mix it with some compound which gives off a large quantity of gas when heated in the tube. Sugar or starch are convenient substances for this purpose. This dilution of the ammonia prevents its too rapid absorption by the acid, and thus diminishes the risk of acid being drawn back into the hot tube.
Having by the above process obtained the nitrogen in the form of ammonium chloride, its quantity is determined by precipitation as ammonio-platinic chloride, (NH 4 ) 2 PtCl 6 , the analysis being performed in the same manner as in the quantitative determination of ammonium by means of platinic chloride. In calculating the results the requisite data are furnished by the facts that 223 ’ 2 parts of the double salt contain 14 of nitrogen, or 197'5 parts of platinum correspond to 28 of nitrogen. Thus, if w = weight of substance taken, W the weight of double salt obtained, P the weight of platinum, and N the weight . . . W×14 P×28 . 100N , of nitrogen, we have N = oo ., or 1 ----; also —— for σ ’ 223'2 197'5 ’ w the percentage of nitrogen.
Instead of determining the ammonia gravimetrically it may be absorbed by a known quantity of dilute sulphuric acid of standard strength, and the amount of acid neutralized determined by adding litmus, and then a standard solution of soda till complete neutralization is effected as in ordinary acidimetrical determinations.
In cases where nitrogen-containing bodies do not yield the whole of that element as ammonia on ignition with soda-lime, Dumas’s process is employed. This process depends upon the fact that all nitrogenous substances yield their nitrogen in the free state when burnt with cupric oxide, and in presence of ignited metallic copper. The method is thus carried out.
A combustion tube of suitable length is sealed off at one end, and about 2 or 3 inches of hydrogen sodium carbonate (NaHCO 3 ) introduced, after which a small quantity of cupric oxide is added. A known quantity of the substance, intimately mixed with cupric oxide, is next introduced, and then a further quantity of pure cupric oxide, the remainder of the tube being finally filled up with 3 or 4 inches of freshly reduced copper gauze plugging. No special precautions need be taken to keep the cupric oxide dry. A gas passage having been secured through the tube in the usual manner, a delivery tube is tightly adapted to the open end of the combustion tube.
The tube is placed in a gas furnace, and the end of the delivery tube plunged beneath the surface of mercury contained in a mercurial trough. The first proceeding is to expel the air from the apparatus. This is effected by heating a portion of the hydrogen sodium carbonate in the posterior part of the combustion tube till a bubble of the gas received into caustic potash solution is wholly absorbed. When the apparatus is entirely filled with carbon dioxide, a graduated receiver containing about ½ of its volume of caustic potash solution, the remainder being filled with mercury, is inverted over the end of the delivery tube as in ordinary cases of gas-receiving over the mercurial trough. The combustion is then proceeded with in the usual manner,—the anterior portion of the tube containing the metallic copper being first heated to redness, and the heat gradually carried back till the combustion is completed, when more of the NaHCO 3 in the posterior part is heated 1 Although the principles upon which the determination of carbon and hydrogen in organic bodies depends remain unaltered, the process has received several modifications which it is unnecessary to describe in detail. Thus in some laboratories, instead of the potash bulbs, a tube filled with “soda-lime” is employed, and the substance, instead of being mixed with the CuO in the tube, is placed in a small boat of porcelain or platinum, and a slow stream of oxygen kept going during the whole process. Other oxidizers have been also proposed instead of cupric oxide, such as potassium perchlorate or a mixture of sulphuric acid and silver iodate, in which latter process the carbon is determined directly from the amount of CO 2 produced, and the hydrogen indirectly from the amount of oxygen consumed minus the quantity contained in the CO 2 , 2 Soda-lime is prepared by slaking quicklime with a strong solution of caustic soda, and then heating till thoroughly dry, so as to expel the last traces of nitrogen. The CO 2 produced by the combustion and from the NaHCO 3 being absorbed by the potash solution, the graduated receiver contains only the whole volume of pure nitrogen. The receiver and its contents are accordingly transferred to a vessel of water—the mercury and potash solution ’ allowed to be replaced by water, the receiver is raised till the pressure is equalized by the water being at the same level both inside and outside, and the volume of gas is read off, the temperature of the air of the room and the height of the barometer being at the same time noted.
If N = weight of nitrogen in grammes, t o the temperature of the air, b the height of the barometer, e the tension of aqueous vapour at the temperature i β , and V the volume of nitrogen in cubic centimetres—
xτ λt- 0∙0012562
∙ n - V ' (l + 0∙00367i o )760 ' °~ e
(0 0012562 being wt. of 1 c.c. of N at 0° C. and 760 mm. bar.).
Determination of ths Halogen Elements, Sulphur, and Phosphorus. — The halogens are sometimes determined as silver salts, by burning a known weight of the substance with pure quicklime in a combustion tube, dissolving in dilute nitric acid, and adding silver nitrate. Sulphur and phosphorus may be determined by fusing a known quantity of the substance with a mixture of potassium hydroxide and nitrate in a silver dish. The sulphur is by this means oxidized to sulphuric and the phoβp∩orus to phosphoric acid, and, on dissolving the fused mass and acidulating, these acids can be estimated.
By the method of Carius the halogens, sulphur, and phosphorus can be determined, if necessary, in one operation. A known weight of the substance is sealed up in a strong glass tube with about 20 times its weight of nitric acid (sp. gr. 1 -4), and the tube then heated for some hours in an oi]-bath to a temperature of 140° - 300° Q The substance is completely oxidized by this operation, the sulphur and phosphorus being converted into their respective acids, so that their determination then becomes an operation of inorganic analysis. If halogens are present, it is customary to add a few crystals of silver nitrate before sealing up the tube. After the operation the haloid silver salt is filtered off, the excess of silver is removed from the filtrate by HC1, and H 2 SO 4 or H 3 PO 4 is determined in the ordinary way.
Determination of Oxygen. — Although several processes have been devised for the direct estimation of oxygen (by Baumhauer, Maumené, and Mitscherlich), they are seldom employed in laboratories. This element is usually determined by “difference,” is., by adding the percentages of the other elements, and subtracting the result from 100.
Formulae and Constitution of Organic Compounds. Empirical Formulae. — Having by the above methods of analysis arrived at the percentage composition of a substance, the next step is to determine its formula. The empirical formula is obtained by dividing each percentage number by the atomic weight of its respective element. Thus, supposing an analysis of common alcohol gave the following percentage numbers :— [table]
Dividing these by the respective atomic weights— c =52T5 = 4∙3,h=^^=13∙0, O = ⅛^⅛2∙2.
12 1 16
These numbers show that the atoms of C, H, and O are present in numbers having the ratios 2:6:1, since by the atomic theory the atoms of each element must exist in a compound in integral numbers. The differences between the integral and fractional numbers are justly assignable to the unavoidable “experimental errors” of analysis. Thus the simplest formula that can be given to alcohol from the foregoing analysis is C 2 H 6 O, and it is usual to express the results in the following manner :— [table]
This imaginary example may serve to show that the determination of empirical formulae cannot be made according to any fixed set of rules. The errors of experiment are seldom so small as in the supposed illustration, and in cases where these are large, and where the substance contains a large number of atoms in its molecule, great difficulty is often experienced.
Molecular Formulae. — The formulae obtained by the method just described express simply the ratios existing between the numbers of atoms in the molecule of a substance, without regard to the actual number of atoms in such a molecule. Reasoning downwards from the law of Avogadro, which has been explained in the foregoing portions of this article, it will be seen that the volume of such molecules as do not undergo dissociation when heated is always equal to the volume of the molecule (2 parts by weight) of hydrogen. To ascertain molecular formula, therefore, all that is necessary is to determine the vapourdensity of the substance as referred to hydrogen. Thus the analysis of benzene, a hydrocarbon obtained from coal tar, leads to the formula CH, but there is no evidence to show whether its molecular formula is CH, C 2 H 2 , C 3 H 3 , C 4 H 4 , C 5 H 5 , or C M H„. By experiment its vapour density is found to be 39, so that its molecular weight is 78. Dividing this number in the ratio C: H, is., 12:1, we obtain 72: 6 as the actual ratio of the weights of C and H existing in the molecule. Seventy-two parts of C correspond to ½½ = 6 atoms, and 6 parts by weight of hydrogen correspond to ⅜ = 6 atoms, so that the molecular formula of benzene is C 6 H 6 .
Again, with respect to alcohol. The vapour-density (H = 1) is 23, so that its molecular weight is 46. This number, however, agrees with the molecular weight of a substance having the formula C 2 H 6 O, since 12 × 2 + 6 +16 = 46; hence this formula must be assigned to alcohol, and we have an illustration of a case in which the empirical and molecular formulae are identical.
Determination of Vapour-density.
In practice, the vapour-density is determined by the methods of Dumas or Hofmann.
Dumas's Method. — In Dumas’s process the weight of a known volume of vapour is ascertained in the following manner (see fig. 8).
A globular glass flask, as light as possible, with a neck fused into it, is first provided. The capacity of the globe may vaιy from ⅛ to ⅜ litre, according to the amount of substance to be operated upon. The neck is drawn out in the blowpipe flame to a capillary termination of about one millimetre diameter, and then bent up so as to project above the surface of the liquid of the bath in which the globe is to be immersed. The globe is first weighed full of air—the temperature and height of the barometer being noted.
By warming the globe, and plunging the point of the neck into some of the liquid of which the vapour-density is to be determined, a few grammes of the latter are introduced. The globe and its contents are then plunged into a bath of water, paraffin, or fusible metal kept at a constant temperature, at least 20° or 30° C above the boiling-point of the substance. As soon as the vapour ceases to rush out of the capillary orifice of the neck, the point is sealed hermetically by the blowpipe flame, the height of the barometer and the temperature of the bath being observed.
The globe, after being cleaned and allowed to cool, is again weighed, the temperature and height of the barometer being at the same time observed. The capacity of the globe is measured by breaking the point of the neck under mercury, when the metal rushes in to supply the place of the condensed vapour. As the expulsion of air by the escaping vapour is seldom complete, there usually remains a bubble of residual air, which must be allowed for by running in a known quantity of mercury from a burette. The total quantity of mercury is then poured out and its volume measured. The calculation is made from these data by the following approximative method :—
Let m = weight of globe + air at the temperature t of weighing and height of barometer b ; m'= weight of globe + vapour at the temperature f of sealing and height of barometer V ; V = capacity of globe in cubic centimetres;
λ = weight of V cubic centimetres of air at t and b.
Then m — λ = weight of vacuous globe; m'-(m- λ)=weight of substance.
Let H = wt. of V c.c. of hydrogen at t' and b' t then the vapourdensity (d) referred to hydrogen is . ∞'- (m— λ)
H
When residual air (r) is found in the globe, V — r=capacity of globe.
In very exact determinations corrections must be made for (1) the expansion of glass, (2) the difference of temperature and pressure between the first and second weighings of the globe, and (3) the difference in density between the drop of fluid remaining in the globe and the density oi mercury. For most chemical purposes, however, the above-given approximation formula is sufficiently accurate.
For high temperatures the globe is immersed in the vapours of boiling mercury, cadmium, or zinc, and the apparatus is modified accordingly.
Gay-Lussac and Hofmann's Methods. —These methods have for their object the measurement of the volume of a known weight of vapour. Gay-Lussac’s method, being available only for substances boiling below 100° C., has been gradually replaced by Hofmann’s modification (fig. 9).
_ A glass tube about 1 metre in length and 20 mm. diameter, closed at one end, is graduated and calibrated. The tube being filled with mercury, and inverted in a vessel of the same liquid, is practically a barometer with an exaggerated Torricellian vacuum. Surrounding this tube is a wider tube, through which the vapour of any liquid boiling at a constant temperature can be passed, and thus the barometer tube and its contents kept at that temperature.
The substance of which the vapourdensity is to be determined is weighed (about ⅛ gram.) in a minute stoppered bottle, and passed up into the Torricellian vacuum. According to the boiling-point of the substance (which is, of course, much lowered by the reduced pressure), the vapour of alcohol, water, aniline, or amylic alcohol is passed through the space between the two tubes till the temperature and volume of vapour remain constant. The height of the mercury column, the temperature to which the vapour is heated, and the height of the barometer in the room being observed, all the necessary data are obtained.
Let m = weight of substance in grammes;
V== volume (in c.c.) occupied by vapour at temperature t ; h'= height of mercury in tube above mercury in reservoir; b = height of barometer in room.
Then b - h = pressure upon vapour.
Let H = weight of V c.c. of hydrogen at a pressure b- h and temperature t.
Then the vapour-density (d) referred to hydrogen is . _ m
H'
For exact determinations at high temperatures the tension of mer∙ cury vapour (e) at the temperature (t) must be allowed for, and the pressure upon the vapour then becomes b—h— e.
In some cases the substance of which the molecular formula is to be determined does not admit of vaporization, being decomposed by heat. With such substances, some method other than the determination of the vapour-density must consequently be resorted to. In the case of acid or basic compounds, the problem admits of easy solution. Thus, supposing we desired to determine the molecular weight of acetic acid without having recourse to a vapourdensity determination. Having ascertained that the acid contains one atom of hydrogen replaceable by metals, or, in other words, that it is monobasic, the silver salt is prepared, and the amount of silver determined. All the necessary data are then obtained. Thus, supposing the analysis to give 64√67 per cent, of metal, the molecular weight of the salt, t.e., the weight containing one atom of silver, will be given by the proportion—
6467: 100 :: 108 .x.
Whence , x = 167 .
The weight of the “acid-radicle” is therefore 167-108 = 59.
And as one atom of H is replaced by the Ag, the molecular weight of the acid is 60.
The empirical formula deduced from the ultimate analysis would be CH 2 O = 30, so that the molecular formula is 2(CH 2 O) = C 2 H 4 O 2 ∙
With poly basic acids the problem is somewhat more complex, but the solution is effected in a similar manner, í.e., by estimating the metal in a normal salt. Silver salts are employed when obtainable, as they are generally anhydrous and easily purified by crystallization.
As a further illustration we now give an example of the determination of the molecular weight of a basic substance. Supposing an analysis of the base triethylamine to have given the following results :— [table]
The base is monacid, forming a hydrochloride containing one molecule of HC1, and this hydrochloride forms a double platinum salt containing two molecules of the hydrochloride to one molecule of platinum; 100 parts of the platinum salt left, on ignition, 32'14 parts of platinum, so that, to find out the amount of salt containing one atom of platinum, we have—
32∙14: 197'5:: 100: x .∙. x=614∙5.
Putting X for the unknown molecular weight of the base, the molecular weight of the salt is—
2HC1= 73∙0
Pt=197∙5
Cl 4 =142'0
2x + 412 ’ 5 It was shown by analysis that 197’5 parts of platinum represented 614'5 parts of the double salt, so that—
2z+412∙5=614∙5 .∙. x=101.
Making use of the percentage numbers given by ultimate analysis, we have the following obvious proportions for finding the weights of the respective elements contained in this molecular weight :— [table]
Whence C=72 0, H = 14 - 9, N = 13 - 9, and the numbers of the atoms are :—
C=¾θ= 6 ,H=ψ=14∙9,N=⅛ 9 = 0∙99.
By the atomic theory these numbers must be integral, so that the numbers of atoms are 6, 15, and 1; and the molecular formula is c 6 h 15 n.
When a base does not readily form double platinum salts, the molecular formula is deduced from the analysis of an anhydrous normal salt. In the case of compounds which are neither acid nor basic, and which do not admit of vaporization, the molecular formula can only be indirectly arrived at by considering the chemical transformations of the compound and its relationship to known substances. A molecular formula obtained by this means implies that, could the substance be vaporized, its molecular volume would correspond to the molecular volume of hydrogen.
Rational, Constitutional, or Structural Formulae. — The molecular formula of an organic compound simply expresses the fact that the molecule of such compound contains so many atoms of each of its constituent elements, and in the earlier stages of the science chemists were contented with such representation of their analytical results. As the science developed, however, it soon became evident that substances might have the same percentage composition, or even the same molecular formula, and yet exhibit under the influence of the same reagents totally distinct characters. These facts, which will be more fully discussed in a subsequent part of this article, led to the necessity of devising some method by which organic formulae could be made to represent the behaviour of the respective compounds under the influence of decomposing agents—in other words, the manner in which the compound was capable of splitting up or of being resolved, and, as a necessary result, the converse fact of representing the manner in which the elements of a compound were grouped together. These rational, constitutional, or structural formulae must be regarded solely from a chemical point of view; they are symbolic representations of chemical facts, and in no way represent the physical grouping of the atoms in space. They may be most conveniently defined as artificial epitomes of the reactions of compounds, indicating that when decomposed the compounds separate into such and such groups, and that when it is possible to bring these groups or radicles together, the compound can in most cases be built up or synthesized.
Let us now, by way of illustration, proceed to consider the method of arriving at the constitutional formula of some typical compound.
The molecular formula of acetic acid, as previously shown by its ultimate analysis and the determination of silver in its silver salt, is—
C 2 H 4 O 2 .
Being a monobasic acid, one of its hydrogen atoms is replaceable by metals. This fact is expressed, as in the case of inorganic acids, by the formula—
H.C 2 H 3 O 2 .
But this formula does not express the whole of the decompositions possible to the acid; the residue C 2 H 3 O 2 being capable of further subdivision, the formula may be further developed. Thus, acetic acid may be formed by the action of acetyl chloride upon water, according to the reaction—
C 2 H 3 O.C1 + OH 2 = C 2 H 3 O 2 ∙H + HC1.
Thus the radicle acetyl C 2 H 3 O is shown to enter into the composition of acetic acid, and the formula therefore becomes—
H.O.C 2 H 3 O.
In confirmation of this formula several reactions might be mentioned in which the acetyl group is left unchanged, while the hydroxyl, HO, is withdrawn and replaced by other elements or radicles. For example—
H,O.C 2 H 3 O + PC1 5 = C 2 H 3 O.C1 + HC1 + POOL
P Ä C Aee.yl.bMd.. ∙W≤Me »gj, 5(H∙O.C 2 H s O) + P 2 S 5 = 5(H,S.C 2 H s O) + P 2 O 5 A«..!...«., »g- m ∣∞ te . a4 . ij-jgj.
Na,O.C 2 H 3 O + C 2 H s O.C1 = C 2 H a O.C 2 H 3 O.O + NaCl.
Sodium acetate. Acetyl chloride. Acetic oxide. chloride
Next with respect to acetyl itself. When acetic acid is electrolyzed, hydrogen is evolved at the positive pole and carbon dioxide and ethane (C 2 H c ) at the negative. Now, ethane can be shown to be identical with di-methyl (CH 3 ) 2 , so that the radicle methyl is thus shown to exist in acetic acid—a fact which receives confirmation from several reactions, two of which may be now considered.
When potassium cyanide acts upon methyl iodide, a substance known as acetonitrile (CH 3 .CN) is produced—
CH 3 I + KCN = CH 3 .CN + KI.
Methyl iodide. Potassium cyanide. Acetonitrile. Potassium iodide.
By heating acetonitrile with water or caustic potash solution, acetic acid and ammonia are formed, thus—
CH 3 .GN + 2OH 2 = CH 3 .CO.O. H + NH 3 .
Acetonitrile. Water. Acetic acid. Ammonia.
When barium acetate is submitted to dry distillation, it decomposes iu the manner shown by the following equation :—
Ba.O2.(CH 3 .CO) 2 = CO.(CH 3 ) 2 + BaCO 3 .
Barium acetate. Acetone. c X i "‰
Thus the most developed formula of acetic acid is CH 3 .CO.O.H, or, as it is more conveniently expressed—
↑ COOH.
The bracket signifies that the two carbon atoms are directly united.
Graphic Formulae. —Graphic formulae having already been explained (see p. 473), it is here only necessary to illustrate their application to organic compounds. The following are typical examples :—
Name of Compound. Rational Formula. Graphic Formula.
Acetic acid. {(χχ)Η f
H—G—H
0=0— ι Ο—H
( c 2 ¾ - H
Triethylamine. N √ C 2 H 5
( C 2 H 5 H—C—H
H—C—H
Η H I Η H
i i I ii
H—C—C N C—C—H
II II
Η Η Η H Name of Compound. Rational Formula. Graphic Formula.
IT ™ í nh 2 Η Ο H
Urea. C0 ( NH 2 ∣ ∣∣ ∣
N—C—N
I I
II H
H ττ H ττ
CH I
Trimethyl-methane. C ∙ qjj 3 jj ø q ¶¶
Μ IM
Μ J il H-C-H
I
H
Isomerism. — It has been mentioned that organic substances may have the same percentage composition and molecular formula, and yet exhibit totally distinct chemical and physical characters. To this phenomenon the general term isomerism is applied, and the compounds are termed isomerides, or are said to be isomeric. The differences between isomeric bodies are well explained on the view that such bodies possess different constitutions, or that their elements are grouped in different ways. Constitutional formulae are thus absolutely indispensable for the representation of isomerides. As will be seen subsequently, there can be several kinds of isomerism, but we shall here restrict the term to two classes of cases.
(1.) Isomerism proper or Physical Isomerism. — In these cases the substances are identical in their composition, vapour-density, and chemical behaviour, but exhibit different physical properties. Thus there are several hydrocarbons known as terpenes, having the formula C 10 H 16 , which exist in the oils of turpentine, lemon, bergamot, orange, &c., and which exhibit the same behaviour under the influence of chemical reagents, differing only in their odour and action upon polarized light. Again, the empirical formula C 4 H 6 O 6 includes several organic acids (tartaric acid being one of the number) which are physical isomerides differing from one another only in their crystalline form and action upon polarized light.
(2.) Chemical Isomerism. — In these cases the compounds are identical in composition and in their molecular formulae, but differ in physical properties and chemical behaviour in certain reactions.
For instance, by taking one atom of hydrogen from the hydrocarbon propane (C 3 H 8 ) we obtain the radicle propyl (C 3 H 7 ), and if we suppose one atom of hydrogen in marsh gas or methane (CH 4 ) to be replaced by propyl, we get the hydrocarbon tetrane or diethyl—
∕H (C 3 H 7 c⅛ c -⅛
(h (h
Methane. Tetranβ.
Now propyl can be written in two different ways, according as the elements are grouped differently; thus—
(CH 3 CH 2 CH 2 ) ’, or
(CHCH 3 CH 3 )'.
The differences are still more strikingly shown by the use of graphic formulae—
H H H H H H
III III
H—C—C—C— H—C-C—C—H
∏ Å Å u
Propyl. Pseudopropyl.
Distinguishing these radicles from each other by the names propyl and pseudopropyl, it is clear that we can have a tetrane containing propyl and another containing pseudopropyl—
∕CH 2 CH 2 CH 3 (CHCH 3 CH 3 rJ H pjJ H
l )H υ )H
( H ( H
Tetrane or propyl-m ethane. Pseudopropyl methane.
These formulae can, of course, be written in a more condensed form; thus—
CH 3 .C 3 H 7 CH 3 .CH(CH 3 ) 2 , propyï-methane. Methyl-pseudopropyL
or
(C 2 H 5 ) 2 CH(CH s ) s .
Tetrane or diethyl. Trimethyl-methane.
The graphic formulae help to show still more clearly that the elements can be grouped only in two different ways—
H
I H H H
H-C-H I I I
H—C—C—C—H
H-C—H I I
H H
H—C—H
H—C—H
H—C—H I
I H
H
Tetrane or d ethyt Trimethjì-methane.
The fact thus shown possible by formulation is borne out experimentally. Two tetranes actually exist,—one, which from its mode of formation can be shown to be propyi-methane or diethyl, having a boiling point of Γ C., the other, which can be shown to be pseudopropyl-methane, boiling at - 15° C.
The experimental confirmation does not, however, end here. The same isomerism can be shown to exist among all the derivatives of these two propyls. Thus we have— [table]
This list might be considerably extended, but sufficient examples have been given to illustrate the phenomena under consideration.
Not only do these isomerides differ in physical properties, but they exhibit different and characteristic transformations under the action of the same reagent. Thus, normal propyl alcohol when oxidized yields propionic acid :—
H H H H H 0
H—C—C—C—O—H + 0.=H—0—C—C—0—H + 0H 2
H H H ∙ H H
Normal propyl alcohol. Propionic acid. Water.
Under the same circumstances pseudopropyl alcohol yields the substance known as acetone or dimethyl ketone :—
H H
I I
H—C—H ' H—C—H
I I
H—C—O-H + O= C=O + OH 2
I I
H—C—H H—C—H
I I
H H
Pseudopropyl alcohol. Acetone. Water.
The distinguishing character of this class of isomers is that the isomerides can be shown to belong to the same series of compounds, or, according to Schorlemmer, “they contain the same number of carbon atoms linked together.”
Reactions that give rise to the formation of a compound capable of existing in two isomeric modifications frequently result in the production of both isomers, but the precise conditions which regulate the relative quantities of the two compounds are not yet known.
Certain relationships have been shown to exist between the physical and chemical properties of isomeric bodies, although numerous exceptions render the exact expression of these relationships an impossibility in the present state of knowledge. Thus, as a rule, the boiling points of the compounds of an isomeric series are lower than those of the normal series (see preceding table)—or, more generally, the boiling point is higher the more simple the constitution of the substance.
Since different amounts of heat are concerned in the production of isomeric bodies, it must be admitted that such bodies are stored with different amounts of potential energy. Adopting this view, some chemists have recently sought an explanation of isomerism in the different amounts of potential energy thus contained in isomerides, and have thrown distrust upon the “constitutional” theory. In taking this view, however, the true position seems reversed —the fact that isomerides contain different stores of potential energy by no means does away with the hypothesis that they possess different constitutions. It seems, on the contrary, that the difference of energy is accounted for on the view that the bodies possess a difference of constitution, since the contained energy results from the relative positions of the atoms or radicles with regard to the intra-molecular chemical forces.
Metamerism. — Compounds having the same molecular formula may result from the combination of totally distinct radicles, and exhibit in consequence not only a marked difference of physical properties, but in almost all cases different chemical transformations under the influence of the same reagent. Such substances are said to be metameric. The subjoined examples illustrate this class of cases. It is to be observed that, as with true isomeric compounds, two metameric bodies are frequently produced in the same reaction. For instance, most of the methods given for obtaining ethyl cyanide yield a mixture of this compound with the isocyanide.
Empirical Formula, C 3 H 9 N.
¶H 7 CH 3 CH 3
N N N
H H C 2 Ilf X H Cnf^C¾
Propylamine. Methyl-ethylamine. Trimethylamine.
Empirical Formula, C 3 H 6 O.
H H 0
H 5 C 3 -0—H H—C—C—C—Ο—II H 3 C-C—CH a
I I H H
Allyl alcohol. Propionic aldehyde. Acetone.
Empirical Formula, C 3 H 5 N.
H 5 C 2 -C≡N H 5 C 2 -N≡C
Ethyl cyanide or Ethyl isocyanide or propionitrile. ethyl carbamine.
Polymerism. — Compounds having the same percentage composition but different vapour-densities are said to be polymeric; thus— [table]
The above list exhibits the rise in boiling-point with increase of vapour-density.
Polymeric bodies may be isomeric or metameric; the following, for example, are metameric polymerides :—
í TT ì
Formic aldehyde j COH ( =CH 2 0
Acetic acid θθβ∏ ∣ =C 2 H 4 O 2 =2(CH 2 O)
Lactic acid ∣ COOH H j = θ3 Η βθ3 = 3(£Η 2 Ο)
Classification of Organic Compounds. Homologous Series. — Carbon being a tetrad element is only saturated by four atoms of a monad element, or by any number of atoms the joint atomicities of which are equal to four. Thus the following are saturated compounds :—
CH 4 ', CHC1 3 ', CO"C1 2 ', CO 2 ", CN'"C1', CI 4 , .
Among inorganic compounds many instances occur in which two or more atoms of the same element unite, such, for example, as in the ferric and manganic salts. In these cases, however, the number of atoms thus entering into combination is seldom great. In the carbon atom we meet with the greatest tendency to unite with similar atoms, and this special property of carbon accounts for the great multiplicity of organic compounds.
Making use of graphic notation, the tetrad carbon atom is thus represented—C—. If two atoms of carbon unite by one bond of each, the resultant atomicity of the group is 6; if 3 atoms unite the resultant atomicity is 8, and so on :—
—C— _ c— -c- — C—
’ X X X
l J μ ~
1 -Ο
ι
Thus every additional atom of carbon brings two active units of atomicity into the molecule, and if we suppose these carbon atoms to be saturated by hydrogen, we shall have an ascending series of hydrocarbons, each member of which differs from the one beneath it by + CH 2 . Such series are known as homologous series. The following are examples :— [table]
Isologous Series. — An inspection of the foregoing formulae will show that hydrocarbons of the general formula C m H 2 , +2 are the only saturated compounds, i.e., the only compounds in which the atom-fixing powers of the carbon atoms are completely satisfied. If an odd number of hydrogen atoms be withdrawn, a compound having a certain odd number of unsatisfied units of atomicity (i.e., a “perissad” radicle) is the result, and such compounds are incapable of existing in the free state. From this and previous considerations it follows, that “all hydrocarbons contain an even number of atoms of hydrogen,” and further, “that the sum of the atoms of monad and triad elements contained in the molecule of a carbon compound must also always be an even number” (Schorlemmer).
Compounds containing an even number of unsatisfied units of atomicity, although non-saturated, are capable of existing in the free state. Thus, commencing with the saturated hydrocarbons of the C n II 2n+2 series, hydrogen atoms can be withdrawn by pairs, giving rise to a descending series of hydrocarbons, each member of which differs from the one below it by + H 2 . Such series are termed isologous series. The following table shows at a glance the relationship between homologous and isologous series, and at the same time the system of nomenclature used :— [table]
Organic Radicles. — The meaning of the term radide has already been explained (p. 474), so that it is here only necessary to point out the part played by such unsaturated groups of atoms in the formation of organic compounds.
It has already been mentioned that an odd number of hydrogen atoms withdrawn from a saturated hydrocarbon of the C w H 2n+2 series leaves an unsaturated group having an odd number 1 of unsatisfied units of atomicity. For example :—
CH 4 -H = (CH 3 )' (c 2 h 4 )'- H = (C 2 H 3 )'"
Methane. Methyl. Ethene. Ethenyl.
c ⅛-H = (¾H 5 >' (¾H β )∙- H = (C s H 5 )'"
Ethane. Ethyl. Propene. Propenyl.
c⅛÷ 2 -h =(C‰ 1 )' (C⅛)'-H=(C.w
Thus the C.H 3h+2 hydrocarbons may be conveniently regarded for some purposes as hydrides of C b H 2m+1 radicles, a series the members of which enter largely into the composition of organic compounds of all classes.
Perissad radicles are incapable of existing in the free state because on isolation two semi-molecules unite; for instance—
i H î ì H i i H H í ∙. H H! i I H I Μ I I H I I:
: H—C-+-C—H:: H—C—C — C—C—H:
i I L≡ I i i I I H I I i
H:: H s .: H H :: H H
Dimethyl = ethane. Diethyl = tetrano.
Generally 2(C n H 2w+1 ) = C 2λ H 4 . +2 .
It is by no means necessary that a group of atoms should be capable of isolation in order to constitute such group an organic radicle. Any unsaturated group which through several reactions remains unchanged may be so regarded. 2 (See, for example, the previously quoted reactions of acetic acid.)
The following are additional examples of organic radicles :—
Cyanogen N≡C—; in the free state, N≡C—C≡N.
z ch 3 .
Acetyl (C 2 H 3 O)' = CO×
COOH
Carboxyl (COOH) ’ j in the free state, oxalic acid, ∣
COOH
While hydrocarbon radicles are positive (see p. 476), cyanogen and the oxygen-containing or acid radicles are negative.
Fatty and Aromatic Groups. — In order to assist in bringing the vast numbers of organic compounds within the scope of some system of classification, chemists frequently adopt the convenient division of them into fatty and aromatic groups.
It will be seen subsequently that most organic compounds may be regarded as derived by substitution from hydrocarbons. 3 Starting, then, with the saturated hydrocarbons C.H 2n+2 , the isologous series, down to C„H 2 ._ 6 , and some of the members of the C λ H 2b-4 series, with their derivatives, constitute the fatty group, so called because many of its members exist in fatty bodies. The hydrocarbons of the C.H 2b „ 6 , &c., series, with their derivatives, are termed the aromatic group, because many of the compounds are obtained from balsams, essential oils, gumresins, and other aromatic substances. The chief characteristic of the aromatic group is the comparative stability of its compounds, for, whereas the artiad (t.e., even) radicles of the fatty group act as unsaturated groups entering freely into direct combination with other elements, the aromatic radicles act more like saturated groups entering into direct combination only with difficulty, and forming substitution compounds with comparative ease.
The division here made between fatty and aromatic substances must not be regarded as one having a sharply defined boundary line. In point of fact, the two series merge into one another, and compounds belonging to one group can be transformed into compounds of the other. Thus benzene, the typical hydrocarbon of the aromatic group, can be formed directly from ethine or acetylene, a hydrocarbon of the fatty group, by the polymerization effected by heat :-—
3C 2 H 2 = C 6 H 6 . Ethine. Benzene.
Also acetone, a substance directly obtainable from the 1 The names of perissad radicles are made to end in yl. 2 So numerous, indeed, are these radicles in Organic Chemistry that this branch of the science has been named the “Chemistry of compound radicles.”
3 “We may therefore define that part of our science which is generally known as Organic Chemistry as the Chemistry of the Hydro· carbons and their Derivatives.”—Schorlemmer.
salts of the fatty body acetic acid, when it is heated with sulphuric acid, loses water and is transformed into mesitylene or trimethylbenzene :—
3CO(CH 3 ) 2 - 3OH 2 = C 6 H 3 (CH 3 ) 3 .
Acetone. Water. Tiimethylbenzene.
Families of Organic Compounds. — Most of the organic compounds at present known can be referred to certain famüies which we shall now proceed to enumerate.
I. Hydrocarbons. — These are the parent bodies from which the remaining families are derived. Their arrangement into homologous and isologous series has already been discussed.
II. Alcohols. — Formed from hydrocarbons by the substitution of hydroxyl for hydrogen :—
CH 4 , CH 3 OH; C 2 H 6 , C 2 H 5 OH , C 2 H 4 "(OH) 2 ;
Methane. Methyl alcohol. Ethane. Ethyl alcohol. Ethene alcohol·
C 3 H 8 , C 3 H 7 OH , C 2 H 6 "(OH) 2 , C 3 H 5 "'(OH) 3 .
Propane. Propyl alcohol. Propene alcohol. Propenyl alcohol.
III. Haloid Ethers. — Formed from hydrocarbons by the substitution of halogen elements for hydrogen :—
C 3 H 8 , C 3 H 7 C1 , C 3 H 6 "Br 2 , C 3 H 5 '"L.
Ο O ' O i ' Ο O Δ ' öD.J
Propane. Propyl chloride. Propene dibromide. Propenyl triiodide.
IV. Ethers. — Derived from alcohols by the substitution of oxygen for hydroxyl; thus, from
2(C 2 H 5 OH) , C 2 H 4 "(OH) 2 , C 3 H 5 "'(OH) 3
Ethyl alcohol· Ethene alcohol· Propenyl alcohol·
we have
(C 2 H 5 ) 2 O , C 2 H∕O , (C 3 H 5 ) 2 "O 3 .
Ethyl oxide or ether. Ethene oxide. Propenyl oxide.
V. Sulphur, Selenium, and Tellurium Alcohols and Ethers. — These compounds are the analogues of the alcohols and ethers, oxygen being replaced by sulphur, &c. The thio-alcohols are known as mercaptans :— C 2 H 5 OH , C 2 H 5 SH , C 2 H 5 SeH , (C 2 H 5 ) 2 S , (C 2 H 5 ) 2 Te;
C 2 H 1 '(OH) 1 , C 2 H 4 '(SH) 2 ; C 3 H 5 "'(OH) 3 , C 3 H 5 '"(SH) 3 . Ethene alcohol. Ethene íulphydrate. Propenyl alcohol. Propenyl »ulphydrate.
VI. Aldehydes. — Derived from hydrocarbons by the replacement of hydrogen by the radicle (COH)':—
CH 4 , CH 3 .COH; C 2 H β , C 2 H 5 .COH .
Methane. Acetic aldehyde. Ethane. Propionic aldehyde.
VΠ. Ketones. — Derived from aldehydes by the replacement of hydrogen in the COH group by monad hydrocarbon radicles :—
CH 3 ,COH , CH 3 ,CO.CH 3 ; C 2 H 5 .COH , C 2 H 5 .CO.CH 3 . Acetic aldehyde. Dimethyl-ketone. Propionic aldehyde. Ethyl-methyl ketone.
VIII. Organic Acids. — These compounds may be regarded as hydrocarbons in which hydrogen is replaced by carboxyl :—
CH 4 , CH 3 .COOH; C 2 H 6 , C 2 H∕.(COOH) 2 .
Methane. Acetic acid. Ethane. Succinic acid.
The organic acids may likewise be regarded as derived from alcohols by the replacement of H 2 by 0 :—
CH 3 .CH 2 .OH , CH 3 ,CO.OH;
Ethyl alcohol· Acetic acid.
⅞HΛpH) 2 , Ç H 2 O'(OH) 2 , C 2 O 2 '(OH) 2 . Ethene alcohol· Glycollic acid. Oxalic acid.
Corresponding thio-acids are known thus
C 2 H 3 O.OH C 2 H 3 O.SH
Acetic acid. Thiacetic acid.
IX. Anhydrides. — Derived from acids by the substitution of oxygen for hydroxyl, and thus bearing the same relationship to the acids that the ethers bear to the alcohols; thus, from
2(CH 3 .CO.OH) , C 2 H∕C 2 O 2 (OH) 2 , 2(C 6 H 5 .CO.OH) Acetic acid. Succinic acid. Benzoic acid.
we have
(CH 3 ,CO) 2 O , C 2 TI∕.C 2 O 2 .O , (C 6 H 5 .CO) 2 O.
Acetic anhydride. Succinic anhydride. Benzoic anhydride Corresponding thio-compounds are capable of existing :— (CH 3 .CO) 2 S = thiacetic anhydride.
X. Add Halides. — Derived from acids by the substitution of halogen elements for hydroxyl. These compounds are thus most conveniently formulated as compounds of an oxygenated or acid radicle with the halogens :—
C 2 H 3 O.HO , C 2 H 3 O.C1 , (C 2 H 3 O) 3 Γ';
Acetic acid. Acetyl chloride. Acetyl iodide.
C 4 H 4 O 2 ' ’.(OH) 2 , C 4 H 4 O 2 ".C1 2 .
Succinic acid. Succinyl chloride.
XI. Ethereal Salts or Compound Ethers. — Derived from acids, organic or inorganic, by the substitution of a hydrocarbon radicle for the hydrogen of the hydroxyl :—
NO 2 OH , NO 2 .O(C 2 H 5 );
Nitric add. Ethyl nitrate.
SO 2 ,(OH) 2 , SO 2 ,OH.(OC 2 H δ ) , SO 2 -(OC 2 H 5 ) 2 ;
Sulphuric add. Ethyl-hydrogen sulphate. Diethyl sulphate.
PO.(OH) 3 , PO.(OH) 2 (OC 2 H 5 ) , PO.(OC 2 H 5 ) 3 ;
Phosphoric acid. Monethyl phosphate. Triethyl phosphate.
C 2 H 3 O.OH , C 2 H 3 O.(OC 2 H 5 ) .
Acetic add. Ethyl acetate
The thio-acids also form ethereal salts :—
C 2 H 3 O.SH , C 2 H 3 O.(SC 2 H δ ) .
Thiacetic acid. Ethyl thiacetate
XII. Organo-metallix Bodies. — Compounds of hydrocarbon radicles with metals :—
NaC 2 H 5 ; Zn(C 2 H 5 ) 2 Sn(C 2 H 5 ) 4 - SnI 2 (C 2 H 5 ) 2 . Sodium ethide. Zinc ethide. Stannic ethide. Stannic <Uiodo-
diethide.
XIII. Amiιιes or Compound Ammonias. — These compounds are most conveniently regarded as derivatives of ammonia and its hydrate, and of ammonium haloid salts, hydrogen being replaced by hydrocarbon radicles :—
f(¾H∕
(H (C 2 H 5 í C 2 H 5 (C g H 5 H
N√H N√H N√C 2 H 5 N C 2 H 5 N 2 H
(h (h (c’h‘ (h ∣h
H
Ammonia. Ethylamine. Triethylamine. Eth ^^ ny1 ' Ethene diamine.
' H (θ 2 ¾ [ c 2 h 5 fC 2 H 5 H H c 2≡5 (r ττ C 2 H 3
N H n]h N C 2 H 5 N∫g¾ n c 2 h h h c 2 h 5 I c ⅛
I Cl [Cl [l [OH
Ammonium Ethylammonium Tetrethyl-ammo-Methyl carb-Tetrethylchloride. chloride. nium iodide. amine. ammonium hydrate.
These compounds may likewise be formulated as hydrocarbon derivatives :—
C≡ i , CH (NH 2 ); C 2 H 6 , C 2 H 5 (NH 2 ) , C 2 H 4 '(NH 2 ) 2 j Methylamine τ ,.. . , - i , ,. .
Methane. or amido-Ethane. Ethylamine or Ethene diamine or
»ethane. amido-ethane. diamido-ethane.
C β H β , C,H 5 (NH 2 ) , C c H∕(NH 2 ) 2 , C 6 H 3 "(NH⅛. Benzene. Amidobenzene. Diamidobenzene. Triamidobenzene.
The phosphorus, arsenic, antimony, and bismuth analogues of ammonia yield derivatives corresponding to the amines :—
(C 2 H δ (C 2 H δ (CH 3
P H P< C 2 H δ As√CH 3
I H IC 2 H 5 I H
Ethylphosphîne. Triethylphosphine. Dimetlιylarsine.
c 2 H 5 c 2 h 5
( C 2 H δ C 2 H δ C 2 H 5
Bi√C 2 H δ Sb-{C 2 H δ P-j C 2 H δ
( C 2 H δ C 2 H δ C 2 H δ
Tetrethylstihonium Tetrethylphosphonium Triethylbismuthine. chloride. iodide.
XIV. Amides. — These compounds are analogous to the amines, being derived from ammonia by the substitution of acid radicles for hydrogen :—
(H {C 2 H 3 O (C 2 H 3 O (C 2 H 3 O
N √H N(H N ∙<C 2 H 3 O N√ C 2 H 3 O (H IH (H (C 2 H 3 O
Ammonia. Acetamide. Diacetamide. Triacetamide.
ÍÍCAO/
H ((C 4 H 4 O 2 )"
NJH N 5 J(C 4 H 4 O 2 )"
H ((¾W
. Succinamide. Trisuceinamide.
The amides may also be written as compounds of acid radicles with amidogen (NH 2 )', imidogen (NH)", and nitrogen N*':—
C 2 H 3 O.NH 2 , (C 2 H 3 O) 2 .(NH) λ ',(C 2 H 3 O) 3 N λ "j(C 4 H 4 O 2 )' , .NH Acetamide. Diacetamide. Triacetamide. Succinimide.
The fourteen families now enumerated, although comprising most of the known organic substances, still leave outstanding a large number of compounds, of which the constitution has not yet been determined, and which consequently cannot be referred to any of the above groups. We propose to consider briefly the individual families in succession, describing the mode of preparation and properties of the most important members of each family, and then proceeding to the consideration of the unclassed organic compounds.
Cyanogen and its Compounds.
The compound which in its chemical behaviour most closely resembles inorganic substances, and which forms as it wère a connecting link between these and organic bodies, is the radicle cyanogen. Before proceeding, therefore, to the systematic consideration of the great organic families, cyanogen and its compounds may be conveniently treated of.
Cyanogen, CN or Cy, in the free state Cy 2 , is generally prepared by heating the cyanide of some heavy metal (usually mercury): HgCy 2 = Hg + Cy 2 . A brown substance, most probably a polymeride of cyanogen, known as paroeyanogen, is always formed in this reaction.
Cyanogen is a colourless gas, having a pungent odour resembling that of bitter almond oil. It burns in air with a purple flame, and is extremely poisonous. The gas is condensable into a liquid under a pressure of about 4 atmospheres. The liquid boils at - 21° C., and solidifies at - 34° C. Water absorbs about 4 volumes of the gas in the cold. The aqueous solution decomposes on standing, ammonium oxalate being the chief product of the reaction: C 2 N 2 + 4OH 2 = (NH 4 ) 2 C 2 O 4 . At the same time small quantities of urea, ammonium carbonate, and cyanide are formed. The addition of a mineral acid to the solution greatly retards the decomposition, oxamide being then produced: C 2 N 2 + 2OH 2 = C 2 O 2 (NH 2 ) 2 . Conversely, when oxamide or ammonium oxalate is heated cyanogen is produced: (NH 4 ) 2 C 2 O 4 - 4OH 2 = C 2 N 2 j C 2 O 2 (NH 2 ) 2 - 2OH 2 = C 2 N 2 .
In its chemical relationships cyanogen is the exact analogue of the halogen elements. Its compounds with metals or positive radicles are called cyanides.
Compounds of cyanogen with Cl, Br, I, S, OH, and NH 2 are known, and are remarkable for their polymeric modifications. The following is a list of the more important compounds :— [table]
Hydrogen Cyanide,or Hydrocyanic or Prussic Acid, HCN or HCy.—This compound is formed synthetically by passing electric sparks through a mixture of nitrogen and ethine gases: C 2 H 2 + N 2 = 2HCN^. In practice benzene vapour may be used instead of pure ethine, as it is partially resolved into the latter substance by the action of the spark. The anhydrous acid is also obtained by passing dry hydrogen sulphide over mercuric cyanide. The aqueous solution of the acid is prepared by the action of acids upon metallic cyanides: HC1 + KCN = KC1 + HCN; by the action of ammonia on chloroform: NH 3 + CHCl 3 = HCN + 3HC1; and also (most conveniently) by heating a mixture of 5 parts of potassium ferrocyanide with 3 parts of sulphuric acid and 4 parts of water.
The pure acid is a colourless liquid, having an odour of bitter almonds; it is a most violent poison. Its boiling-point is 26 0, 5 C. and its point of solidification - 15° C. The pure acid and its strong aqueous solution are both inflam ’ mable, burning in air with a violet flame. Both thé anhy drous and aqueous acids are very unstable, the formel decomposing into ammonia and a brown substance, and the latter undergoing the same decomposition with the additional formation of ammonium formate. Small quantities of formic or of a mineral acid prevent this décomposition, but on mixture with strong acids, a complete decomposition into formic acid ensues :—
HCN + 2OH 2 + HC1 = H.COOH + NH 4 C1 Hydrocyanic w f Hydrochloric Formic Ammonium acid. water. a cid. add chloride.
Alkalies induce a similar change, alkaline formate and free ammonia being produced. When ammonium formate is heated, the inverse reaction takes place :—
HCO(ONH 4 ) - 2OH 2 = HCN Ammonium Hydrocyanic formate. water ∙ acid.
A polymeride, H 3 C 3 N 3 , is known.
Metallic Cyanides. — Cyanogen being a monad radicl∙ forms, like Cl, Br, and I, a series of salts typified by the formulae M'Cy, M' ’ Cy 2 , M"' 2 Cy 6 , M"Cy n . Of these the most important is potassium cyanide, which can be formed by passing nitrogen over a mixture of red hot carbon and potassium carbonate :—
K 2 CO 3 + 4C + N 2 = 2KCN + 3CO.
This salt is also produced by heating potassium ferrocyanide either alone or mixed with potassium carbonate. Most of the metallic cyanides can be prepared by the action of hydrocyanic acid upon the oxides and hydroxides of the metals, or by double decomposition.
The cyanides have a remarkable tendency to form double salts, such, for example, as the double cyanide of potassium and silver, KCy,AgCy.
Of these double cyanides a certain number are decomposed by the action of a mineral acid yielding free HCy and salts of the acid :—
KCN, AgCN + 2HNO 3 = KNO 3 + AgNO 3 + 2HCN^.
Other double cyanides do not decompose in this manner by the action of mineral acids; the heavy metal cannot be detected by its ordinary reagents, and no HCy is evolved. Thus the double cyanide of potassium and iron, 4KCy,FeCy 2 , when acted on by acids exchanges its potassium for hydrogen, but the iron is not removed :—
K 4 FeCy 6 + 4HC1 = H 4 FeCy 6 + 4KC1.
Potassium ferrocyanide. Hydroferrocyanic acid.
Thus the group FeCy 6 is regarded as an acid radicle lferrocyanogen), and a large number of its salts (ferrο-cuanides} are known. In the same manner we have— [table]
Of these compounds the most important are the ferrocyanides and ferricyanides of potassium.
Potassium ferrocyanide (known commonly as yellow prussiate of potash), is prepared on the large scale by fusing refuse animal matter, such as horn parings, leather scraps, &c., with crude potassium carbonate and iron filings. This salt is also formed when a ferrous salt is added to a solution of potassium cyanide—
6KCy + FeSO 4 = K 4 FeCy 6 + K 2 SO 4 .
Potassium ferrocyanide crystallizes in large pale yellow crystals belonging to the quadratic system, and having the composition K 4 FeCy 6 , 3OH 2 , Solutions of ferric salts produce in an aqueous solution of potassium ferrocyanide a deep blue precipitate of ferric ferrocyanide or Prussian blue :—
3K 4 FeCy β + 2Fe 2 Cl 6 = 2Fe 2 Cy p .3FeCy 2 + 12KC1.
Prussian blue.
Soluble ferrocyanides are thus a delicate test for the presence of ferric salts; and conversely, ferric salts are used to detect ferrocyanides, and also cyanides by the simultaneous addition of a ferrous salt (Scheele’s test for prussic acid).
Potassium ferricyanide (or red prussiate of potash) is prepared by passing chlorine into a solution of the ferrocyanide :—
2K 4 FeCy 6 + Cl 2 = K 6 Fe 2 Cy 12 +2KC1
Ferricyanide.
The ferricyanide forms large prismatic crystals of a dark red colour soluble in water. Ferric salts give a brown coloration with ferricyanides, while ferrous salts give a blue precipitate of Turnbull's blue (Fe 5 Cy 12 + Aq).
Nitroprussides are salts of the general formula M 2 'Fe"Cy 5 NO obtained by the action of nitric acid upon ferro-and ferri-cyanides.
Cyanides of Hydrocarbon Radicles. — The compounds formed by the union of cyanogen with hydrocarbon radicles are of the utmost interest and importance on account of their metameric modifications. Thus we have—
C 2 H 6 , C 2 H 5 .CN , C 2 H∕.(CN) 2 .
Ethane · Eth a yi e cy a aX Ethene
On developing the graphic formulae of one of these cyanides containing a monad radicle, taking for brevity the simplest case, viz., CH 3 .CN, methyl cyanide, it will be seen that two metamerides are possible (see also p. Õ51)
H H
H—C— N≡C H—C—C≡N
i ⅛
In the one compound the carbon of the radicle is in combination with the nitrogen of the cyanogen. In the other compound the same carbon atom is in combination with the carbon of the cyanogen. Two such metameric series are actually known, one (cyanides or nitriles) being formulated as hydrocarbon derivatives, and the other (isocyanides or carbamines) as ammonia derivatives or amines (see p. 553); thus (representing the monad radicle by R')—
( h íïï l∏ ( IST'" I ∙ α ∙ ( Γ«’
c H > C {r· n ∣≡ N ’ R-
( h < h u ' fl "≡- c ≡ 0r , ≡⅛ r
In accordance with what has been previously said concerning isomeric bodies (p. 551), these two series exhibit different behaviour under the action of the same reagent. For instance, the following equations show the ultimate action of water on them :—
Cyanides. Iβocyanides.
f √N"' (R' n , <C" ÍH , π b I r , +20h 2 =nh 3 +c j o" λ 1r , +2oh 2 =n h + j g 0011
Acid Amine. Formic acid.
In accordance with these reactions cyanides or nitrile» can be produced by the action of dehydrating agents on the ammonium salts of the corresponding acids—
H.CO(ONH 4 ) - 2OH 2 = H.CN = (CH)"'N
Ammonium formate. Hydrogen cyanide. Formonitrile.
CH 3 .CO(ONH 4 ) - 2OH 2 = CH 3 .CN = (C 2 H 3 )" ’ N
Ammonium acetate. Methyl cyanide. Acetonitrile.
C 2 H 5 .CO(ONH 4 ) - 2OH 2 = C 2 H 5 .CN = (C 3 H 5 )'"N
Ammonium propionate. Propyl cyanide. Propionitrile.
C 6 H 5 .CO(ONH 4 ) - 2OH 2 = C 6 H 5 .CN = (C 7 H 5 )" , N
Ammonium benzoate. Phenyl cyanide. Benzonitrile.
The action of water upon nitriles thus gives rise to the formation of an acid containing the same number of carbon atoms as the nitrile; the radicle remaining unchanged, we may consider that in these reactions cyanogen (CN) is converted into carboxyl (COOH)—
R'.CN + 2OH 2 = NH 3 + R'.COOH.
It has been further stated that both the nitriles and organic acids can be formulated as hydrocarbon derivatives —in the former the H of the radicle being replaced by (CN)', and in the latter by (COOH)'. Starting then with the cyanide of a hydrocarbon radicle, we get by the action of water an acid containing one atom of carbon more than the hydrocarbon from which the acid is derived, and in this manner acids can be built up from their parent hydrocarbons, and the number of their contained semimolecules of carboxyl increased; thus—
H.CN H.COOH
Hydrogen cyanide is converted by the action of water into Formic acid.
CH 3 .CN CH 3 ∙COOH
Methyl cyanide. t ,, Acetic acid.
C 2 H 4 ".(CN) 2 C 2 H 4 ".(COOH) 2
Ethene cyanide. s , i , Succinic acid.
R*.(CN) n R n .(COOH)n CH 2 .CN.COOH CH 2 ∙(COOH) 2
Cyanacetic acid. t . v Malonic acid.
R' - H n . (CN) n . (COOH) R' - H n . (COOH) n+ .
The cyanides of the hydrocarbon radicles will be again referred to when treating of the haloid ethers, of which bodies they may be considered the analogues. The isocyanides will be treated of as amines.
I. Hydrocarbons.
Starting with the saturated compounds of the C m H 2m+2 ceries, the isologous series C h H 2m , C n H 2 ,. 2 , C.H 2w - 32 are known.
First Series, C m H 2m+2 . — Marsh Gas or Paraffin Series.
The first member of the series is CH 4 (marsh gas or methane), and the homologues down to C 16 H 34 (hexdecane) have been obtained. The names and formulae will be given later on.
General Properties. — The members of the series exhibit a regular gradation in physical properties with each successive addition of CH 2 . Thus, the first four members are gaseous at ordinary temperatures, while the succeeding terms are liquids of increasing specific gravity and viscidity as the series is ascended; those containing 20 or more carbon atoms are solid crystalline bodies. The boiling-point increases as the series is ascended, the difference between the boiling-points of the successive terms of the normal series decreasing regularly by about 4° C. as far as C 12 H 26 , after which there appears to be a constant difference of I9° C.
The paraffins are chemically distinguished for their indifference, hence the name applied to the series (parum affinis). Being saturated molecules, they are incapable of uniting directly with any other element or radicle. The halogen elements produce by their action on paraffins substitution derivatives (haloid ethers), in which the hydrogen of the hydrocarbon is replaced by the halogen.
CH 4 , CH 3 C1 , CH 2 C1 2 , CHC1 3 , CC1 4 ;
C 2 H β , C 2 H 5 C1 , C 2 H 4 C1 2 , C 2 H 3 C1 3 , C 2 H 2 C1 4 , C 2 HC1 5 , C 2 C1 6 , &c., &c.
These compounds will be again referred to under the family of the haloid ethers. Sulphuric acid is without action on paraffins, and at ordinary temperatures they are scarcely attacked by oxidizing agents; but on heating with these agents they are either entirely oxidized to CO 2 and OH 2 , or, in addition, to acids of the acetic (C - H 2 , +1 COOH) and succinic [C m H 2m .(COOH) 2 ] Series. Nitric acid exerts no action in the cold, but the concentrated acid attacks the higher members of the series when heated with them, forming nitro-substitution compounds in which hydrogen is replaced by nitryl (NO 2 ), and, in some instances, acids of the acetic and succinic series and nitriles. The lower members of the series are not acted on by nitric acid, but their nitro-derivatives can be obtained by indirect methods, the general method of preparation being the action of silver nitrite on the iodo-paraffin :—
• C⅛ l+1 I + AgNO 2 = C n II 2 , +1 (NO 2 ) + Agi.
These nitro-paraffins are interesting as being metameric with the corresponding ethereal salts of nitrous acid; thus—
ÌH ∙ CH
Nq , is metameric with methyl nitrite, 0 ^θ 3
Nitroethane, C , i , „ ethyl nitrite, 0
Generally—
C n H 2n+1 C n H 2n+1 are metameric with ∣
O=N=O O—N=O
Nitro-paraffins. Ethereal nitrites.
In accordance with the above formulation the two classes of compounds are differently acted on by nascent hydrogen.
C m H 2m+1 C k H 2 , +1
I +3H 2 = 2OH 2 + I
NO 2 nh 2
Nitro-paraffin. Amine, or amido-paraffin.
C b H 2k+1 C m H 2m+1
I +3H 2 = OH 2 + NH 3 + I O—NO OH
Ethereal nitrite. Ammonia. Alcohol.
The higher members of the series are resolved by prolonged heating in sealed tubes into members lower in the series and members of the isologous series, C n H 2n (olefines), in accordance with the general equation—
C w+i H 2 (n+p ) +2 = C n H 2n+2 + C j ,H 21 , . Occurrence of Paraffins. — Methane, or marsh gas, is an invariable product of the decomposition of organic matter excluded from the air, and is thus produced at the bottom of stagnant ponds, marshes, &c. This gas is the “firedamp” of coal mines; it is found in volcanic gases, and the gas issuing from a mud volcano in the Crimea is almost pure methane. The normal paraffins are found in coal formations and other bituminous strata. Thus, American petroleum contains a mixture of all the paraffins from marsh gas to pentadecane (C 15 H 32 ), or even higher members. In boring for rock oil large quantities of methane escape, and from the freshly drawn petroleum ethane and propane are given off as gases at the ordinary temperature of the air. On distilling off the liquid portion (known commercially as “paraffin oil ”) the higher members of the series remain as white crystalline solids. Many paraffins occur also in the oil obtained by the destructive distillation of Boghead and cannel coal. Solid paraffins are likewise formed in the destructive distillation of wood, coal, bituminous shale, &c., and are found native, as fossil wax, ozokerite, hatchetin, &c.
General Methods of Formation of Paraffins.
1. By the action of nascent hydrogen upon haloid ethers of C n H jn+1 radicles—
CnH 2n 4- 1 AΓα + H a = C⅛ +l + HHa ; (Ha being used to represent a halogen).
2. By the action of zinc and water upon the iodides of C n H an + 1 radicles—
2C n H 2n+1 I + 2Zn + 2OH s = Zn(HO) 2 + Znl 2 + 2C n H a n-∣- a .
3. By the action of water upon the organo-metallic bodies of the general formula Zn(C n H sn + 1 ) a —
Zn(C n H 2n+1 ) a + 2OH a =Zn(H0) a +2C n H an+a .
4. By heating iodides of C n H an+1 radicles with zinc—
Zn + 2C n H an+1 I = Znl 2 + C 2n H 4n ψ a .
In this reaction the paraffin is, at the same time, resolved into a lower member of the series containing half the number of carbon atoms, and the corresponding isologue of the C n H 2 n series—
C an H 4 n+ a = CnH an -∣- a + C n H an .
5. By the action of sodium on the iodides of C n H an + 1 radicles— 2C n H an+1 I + Na a = C 2n H 4n+2 + 2NaI.
6. By the electrolysis of the fatty acids of the series C n H an O a (see ascetic series)—
2C n H 2n O a = 2CO a + C an —»H 4 n— 2 + H 2
7. By heating the dry sodium salts of C n H a n+ 1 . COOH and C n H an .(COOH) a acids with sodium hydroxide—
C n H 2 n +1 . COONa + NaHO=Na a CO 3 + C n H an+a
C n H an (C∞Na) a + 2NaH0 = 2Na a CO 3 +C n H an+a .
8. By the dry distillation of acetates and butyrates several paraffins are formed.
9. Almost all organic compounds yield paraffins when heated in sealed tubes to a very high temperature with excess of strong hydriodic acid.
10. Many paraffins have been obtained by acting on a mixture of the corresponding iodides and ethyl alcohol with copper-coated zinc foil (Gladstone and Tribe’s “copper-zinc couple ”)—
C, i H an+1 I + C 2 H 5 .OH + Zn=ZnI(C 2 H 5 O) + C n H an+a .
Formulation and Classification of Paraffins. — All the members of the paraffin series can be regarded as methane, in which one atom of hydrogen is replaced by C n H 2n+1 radicles; thus—
(H ( CH 3 ( C 2 H 8 ( CH 2 .CH 3 ( C n H 2 n +1 p)H ∩ ) H ο ) H nr (’ ) H η ) H
C )H C )H C )H orb )H c )H
(H (h (H (h (H
Ethane, or p,™»™ nr
Methane. methyl-cthyl-methane.
methane. j
They may likewise be formulated as hydrides of C m H 2m+1 radicles, in accordance with their formation from the haloid ethers of these radicles by the action of nascent hydrogen—
CH 4 = CH 3 .H; C 2 H 6 = C 2 H 5 .H; C 3 H 8 = C 3 H 7 .H .
Methane, hydride Ethane. Ethyl hydride. Propane. Propyl hydride.
More generally, the paraffins may be regarded as formed by the coalescence of any hydrocarbon radicles, furnishing by their addition the necessary number of carbon and hydrogen atoms; thus—
CH 4 = CH 3 'H = CH 2 ".H 2 = CH'".H 3
Methane. Methyl hydride. Methene hydride. Methenyl hydride.
C 2 H 6 = CH 3 .CH 3 = C 2 H 4 ".H 2 = C 2 H 3 ' , '.H 3 A Ο Ο O A ⅛ A A Ο O
Ethane. Μ ®„ dtaethy " β, Ethene hydride. Ethenyl hydride.
C 3 H 8 = CH 3 .C 2 H 5 = ð O O A u
Prcmane Ethyl ∏>etnane, rropane. 0r methyl ethane .
∕CH 3 CH 2 7CH 3 ) 2 or C √ ι c ι ∏ 3 or II 2 C∫
(h 3
Methene dimethide, or dimethyl methane c 4 h 10 = c,⅝c 1 n s = CH,.C s H τ = C 2 H 1 "(CH s ) 2 = Butane, or Ethyl ethane, Propyl methane, or Ethene dlmethide or tetrane. or diethyl. methyl propane. dimethyl ethane.
CH -(CH 3 ) 3 .
Trimethyl methane, or methenyl trimethide (isotetrane).
The number of possible methods of representing a paraffin thus greatly increases with the complexity of the molecule, but it must not be inferred from these formulae that the radicles represented as composing a paraffin molecule have a separate existence in the compound. Such formulation expresses simply the possible modes of formation by which the compound can be produced. For instance—
a. The ethyl hydride obtained by the action of nascent hydrogen upon ethyl iodide (C 2 H b I + H 2 = C 2 H 6 + HI) is identical with the ethene hydride produced by the action of nascent hydrogen on an ethene haloid ether (C 2 H 4 "I 2 + 2H 2 = C 2 H 4 "H 2 + 2HI), and with the dimethyl formed by heating CH 3 I with asmetal [2CH 3 I + Zn= (CH 3 ) 2 + Znl 2 ].
β. The propane (propyl hydride) obtained by the action of nascent hydrogen on propyl iodide, &c. (C 3 H 7 I + H 2 = C 3 H 7 .H + HI), is identical with the ethyl-methyl produced by the action of a metal on a mixture of the iodides of methyl and ethyl (CH s I + C 2 H 5 I + Na 2 = CH 3 .C 2 H 5 + 2NaI), or by the action of zinc-ethyl on methyl iodide [Zn(C 2 H 5 ) 2 + 2CH 3 I=2CH 3 .C 2 H 6 +ZnI 2 ].
γ. Methane obtained by the action of nascent hydrogen on methyl iodide is identical with the methenyl hydride formed by the action of nascent hydrogen on chloroform (CH"'Cb + 3H 0 = CH"'. H, + 3HC1).
Thus it must not be supposed that, because ethane may be written as dimethyl, ethane contains methyl. On treating ethane with chlorine, for example, we do not obtain methyl chloride (CH 3 C1), but substitution products of ethane, C 2 H 5 C1, and similar relations obtain throughout the series.
Turning to the graphic formulae made use of in illustrating the formation of homologous series by the continuous coalescence of carbon atoms with the consequent increase of atomicity (p. 552), it will be seen that the homologous series of C n H 2n+1 radicles can be regarded as derived from the first member, methyl, CH 3 , by the continuous addition of methene, CH 2 "; and as the paraffins can be regarded as derived from the first member, methane, by the substitution of C.H 2b+1 radicles for hydrogen, we have the following constitutional formulae for the four first members :— [table]
Thus, with the increase of the number of atoms in the molecule we have an increased number of hydrocarbon radicles coalescing to form the paraffin; in other words, we have increased complexity of structure, and thus the possible modes of arrangement, or the possible number of isomerides (see p. 550), becomes greater as the number of atoms becomes greater. The three first members, as will be seen from the above formulae, can only be written in the manner shown, and no isomerides exist. The fourth member, butane or tetrane, as already shown, when treating of isomerism (p. 550), can be written in two ways, and two isomerides are known. Similarly there can be three pentanes, four hexanes, six heptanes, &c.
It has been found by Schorlemmer that all the paraffins of which the constitution is known can be classified under four series, viz. :—
1. Normat paraffins, in which no carbon atom is combined with more than two other carbon atoms. (See formulae above. )
2. Isoparaffins, in which one carbon atom is combined with three others. Typical formula :—
í C m H 2 ∙+ 1 HC j C⅛ +1 ( C 7 >H 2i >+ 1
3. Neoparaffins (Odling), in which one carbon atom is combined with four others. Typical formula :—
( c"⅛ n+1
C ( ~*W≡"+ι ∕ p⅛ p ÷ 1 ( υ r H 2 , +1
4. Mesoparaffins (Odling), containing the group [HC(CH 3 ) 2 ]'twice. Typical formula :—
HC(CH 3 ) 2
C⅛
HC(CH 3 ) 2
With regard to the general properties of a paraffin, as compared with those of its isomerides, it has been observed that the boiling-points and specific gravities of the normal compounds are higher than those of the isomers. With respect to chemical stability, the normal paraffins are more difficultly decomposable than their isomers.
The following list contains the names, formulae, boiling-points, and specific gravities of the most important paraffins known at the present time :— [table]
[table] [table] [table]
We now proceed to give a few methods of preparing some of the more important paraffins, which methods do not come under the previously described general processes of formation.
Methane. — This hydrocarbon is found as a constituent of coal-gas where it is produced by the destructive distillation of coal. It is formed synthetically by passing a mixture of carbon disulphide vapour and steam (or sulphuretted hydrogen) over copper heated to redness in a tube :—
CS 2 + 2SH 2 + 4Cu = CH 4 + 4CuS.
Tetrane may be prepared by exposing ethyl iodide over mercury to the action of sunlight :—
2C 2 H 5 I + Hg = Hgl 2 + C 4 Hj 0 . Isopentane (ethyl-dimethyl-methane) is formed by the dehydration of amyl alcohol by means of zinc chloride.
Neopentane (tetramethyl-methane) is prepared by the action of zinc-methide on trimethyl-iodomethane(katabutyl-iodide) :—
2C(CH 3 )I + Zn(CH 3 ) 2 = Znl 2 + 2C(CH 3 ) 4 .
Hexane is produced by the action of hydriodic acid on benzene at 280° :—
C 6 H 6 + 8HI = C 6 H 14 + 4I 2 .
Meso-hexane (tetramethyl-ethane) is formed by the action of hydriodic acid on pinacone a substance having the formula
HOC(CH 3 ) 2
HOC(CH 3 ) 2 produced by the action of sodium amalgam on dimethylketone in presence of water.
Isolaeptane (tetryl-dimethyl-methane) is produced by the action of sodium on a mixture of the bromides of ethyl and amyl.
Solid Paraffin is, when purified, a white wax-like substance, melting between 40° and 60°, and boiling about 370°. It is contained in the tar produced by the distillation of Boghead and cannel coals, and is probably a mixture of several of the higher members of the C.H 2b+2 series.
Second Series, C n H 2 ,.— Olefine Series. General Properties. — These hydrocarbons furnish a good illustration of polymerism, as they form a series which is both homologous and polymeric. Being unsaturated compounds, they act as dyad radicles, uniting with Cl 2 , Br 2 , 0, HC1, HBr, HI, C1HO, &c. The haloid compounds of the olefines treated with alcoholic potash give up one halogen atom and furnish mono-haloid derivatives :—
C n H 2 √∕α 2 + KHO = C b H 2 ^ 1 5λ + KAea + OH 2 .
These mono-haloid olefines can again combine with IIa 2 and again lose Ha by the action of KHO :—
C n II 2n -√∕α√Az 2 + KHO = C n H 2n -⅛ + K77α + OH 2 .
Di-haloid olefine.
In this way 3,4 n atoms of hydrogen can be replaced by Ha 3 , Ha i , Ha n atoms of halogen, and the final product can again unite with Ha 2 , forming a per-haloid paraffin derivative. For instance :—
C 2 C1 4 + Cl 2 = C 2 C1 6 .
Tetraclιlorethene. Hexachlorethane.
The compounds of olefines with hypochlorous acid yield the corresponding alcohols by the action of nascent hydrogen :—
C n H 2n ClHO + H 2 = C n H 2n+1 .OH + HC1.
Many olefines combine with sulphuric acid, forming ethereal salts, which are converted by the action of water into sulphuric acid and the corresponding alcohol :—
C⅛ + H 2 SO 4 = C n H 2n+1 HSO 4 .
C n H 2n+1 HSO 4 + H(HO) = C n H 2n+1 .OH + H 2 SO 4 .
General Methods of Formation. — 1. By heating the C n H 2n+1 OH alcohols with dehydrating agents (H 2 SO 4 , ZnCl 2 , <fcc ) :—
C⅛ +1 OH-OH 2 = C⅛.
2. By electrolyzing C n H 2n (COOΞ) 2 acids :— C n H 2n (COOH) 2 = C n H 2n + 2CO 2 + H 2 .
3. By heating N(C n H 2n+1 ) 4 HO amines :— N(C n H 2 , 1+1 ) 4 HO = N(C n H 2n+1 ) 3 + C n H 2n + OH 2 .
4. By the action of copper-coated zinc foil (Gladstone and Tribe’s “copper-zinc couple”) on the corresponding dibromides :—
C n H 2n Br 2 + Zn = ZnBr 2 + C n H 2n .
Isomerism. — The isomerides of the olefine series have been less completely investigated than those of the paraffins. The formula of the first member (ethene) might be written—
CH 2 — ch 2 ch 3 î| or I or ∣ CH 2 —CH 2 =CH
The first is probably the radicle in the free state, and the second the radicle in combination with Cl, Br, &c. The third formula represents a compound unknown in the free state, but known in combination. While the normal series are known as ethene compounds, the isomeric series are known as ethidene compounds.
Similarly there may be 4 propenes, 9 tetrenes, &c.
The following is a list of the C n H 2n hydrocarbons known at the present time. An idea of the general physical properties will be obtained from the table. [table]
The numerous isomerides of the higher members of the series cannot be described within the limits of the present article.
The first member, ethene, is a colourless inodorous gas, which can be condensed by cold and pressure. It burns with a brilliant flame, and is the constituent to which coal-gas chiefly owes its luminosity. Several olefines can be synthesized from other hydrocarbons; propene, for instance, by passing methane and carbon monoxide through a red-hot tube :—
2CH 4 + CO = OH 2 + C 3 H 6 .
Third Series, C n H 2n - 2 .
General Properties. — These hydrocarbons are capable of acting as tetrad radicles uniting with 4 atoms of Cl, Br, ⅛c., and producing tetra-haloid paraffin derivatives; or the combination may be limited to 2 atoms, thus forming di-haloid olefines. Similarly they can combine directly with one or two molecules of haloid acids, producing C n H 2n ^ 1 Hα, or C n H 2n Z7α 2 . These combinations show their relations to the olefines and paraffins. Their relationship to the latter parent hydrocarbons is further shown by the ultimate action of chlorine: —
CJÏ ^ + C^CJI^¾· C n H 2n-2 Cl 2 + Cl 2 = C,H 2n - 2 01 4 .
Perchlorinated paraffin.
General Methods of Formation. — 1. By the action of an alcoholic solution of potassium hydroxide on mono-haloid olefine derivatives :—
(‰⅛ + KHO = C n H 2n-2 + KAea + OH 2 .
2. By the electrolysis of C n H 2n-2 (COOH) 2 acids;·— C⅛ 2 (COOH) 2 = C n Ξ 2n , 2 + 2CO 2 + H 2 .
The following is a list of the hydrocarbons of this series :— [table]
But few of the isomerides of this series are known. Diallyl, C 6 H 10 (b. p. 51 ’ C.), is isomeric with hexine, and rutylene, C 10 H 18 (b. p. lbθ°), with decine.
Ethine or acetylene may be formed directly from its elements by causing the electric arc to pass between poles of dense carbon in an atmosphere of hydrogen. It is formed also by passing a mixture of methane and carbon monoxide through a red hot tube(CH 4 + CO = C, 2 H 2 + OH 2 ), bypassing chloroform vapour over red hot copper (2CHCl 3 + 6Cu = 3Cu 2 C1 2 + C 2 H 2 ), and by the imperfect combustion of most organic compounds. Ethine is a colourless, incondensable gas, having a characteristic odour. It is absorbed by an ammoniacal solution of cuprous chloride, forming a red precipitate of cuprous acetylide, (C 2 HCu 2 ) 2 O, which when heated with HC1 furnishes pure ethine, so that this is a convenient method for purifying the hydrocarbon. Ethine can be made to combine with nascent hydrogen with the formation of ethene: C 2 H 2 + H 2 = C 2 H 4 .
Fourth Series, C n H 2n-i .
The hydrocarbons of this series may be conveniently divided into two groups. The first, consisting only of one member, pentone or valylene (C 5 H 8 ), is related to the preceding series of hydrocarbons, since, by the addition of bromine the hexabrom-paraffin C 5 H 6 Br 6 is formed.
The other group is composed of decone (C 40 H 16 ) and a series of hydrocarbons known as terpenes, having the general formula C 10 H 16 , and found in the essential oils of various plants, chiefly of the orders Cowiferoe and Aurantioeeoe. Terpenes. — These hydrocarbons are found in the wood, leaves, &c., of Coniferous trees, and in the essential oils of lemon, lavender, bergamot, juniper, &c., &c. Many of these oils contain oxidized bodies besides terpenes. The terpenes possess very similar chemical properties, differing chiefly in their boiling-points (which range from 160° to 176° C.), specific gravities, and action upon a ray of polarized light, thus furnishing a good illustration of physical isomerism (see p. 550). They are generally clear, highly refractive liquids, possessing characteristic odours, and easily polymerized by the action of heat, H 2 SO 4 , BF 3 , &c. Several polymerides exist ready formed in oils of copaiba, cubebs, &c. The terpenes are characterized by the extreme readiness with which they undergo isomeric changes under the influence of reagents. They act as unsaturated compounds combining with at most four monad atoms. Thus, they combine with HC1, forming such compounds as C 10 H 17 Cl, C 10 H 18 Cl 2 , and some terpenes can be made to combine with water, forming such bodies as C 10 H 20 O 2 , C 10 H 18 O, and C 20 H 34 O, which, in certain respects, resemble alcohols. The best known terpene is turpentine oil, which is obtained from the resinous exudation of certain French and American species of Pinus and Abies. Camphor (C 10 H 16 O) is an oxidized compound closely allied to the terpenes contained in the leaves of Laurus Camphora.
The relationship of the terpenes to the group of aromatic hydrocarbons is shown by the production of cymene (C 10 H 14 ), a member of the C n H 2n-6 or benzene series, from turpentine oil and camphor, by the action of certain reagents.
Fifth Series, C„H 2n _ 6 .— Benzene or Aromatic Series.
The general characters of these hydrocarbons, as compared with the preceding series, have been previously considered (p. 552).
Occurrence. — Small quantities of these hydrocarbons are found in petroleum, but the chief source is the tar obtained in the destructive distillation of coal for the manufacture of coal-gas.
General Method of Synthetic Formation. — By heating a mixture of a mono-iodated paraffin and a brominated benzene hydrocarbon with sodium :—
(C n H 2n+1 I) m + C n H 2n _ (6+m) Br m + Na 2m =
C - H 2 ^ β+w >.(C n H 2n ' +1 ) w + (NaI) m + (NaBr) m .
Formulation and Isomerism of Benzene and its Derivatives. — Like the paraflins, the C n H 2n-6 hydrocarbons form a parent series giving rise to a vast number of derivatives, and just as the higher paraffins can be regarded as formed from methane by the substitution of C n H 2n+1 radicles for H, so the homologues of benzene can be considered as derived from this latter hydrocarbon by similar substitutions. Thus—
C 6 H 6 C 6 H 5 .CH 3 = C 7 H 8 C 6 H 5 .C 2 H 5 = C 8 H io Benzene. Methyl benzene. Ethyl benzene.
C 6 H5.C 3 H 7 = c 9 h 12 Propyl benzene.
The synthesis of benzene from ethine has already been alluded to (p. 552). In accordance with this mode of formation, the structural formula of benzene may be written—
H H
^C- C
/’ ∙x
H—C> <C—H
∖ δ 6 ∕ c=c ∏ z ∖
This formula, as will be seen on inspection, is in accordance with the fact that benzene can act as a hexatomic radicle, combining directly (though with difficulty) with 6 halogen atoms. Owing to the complex structure of the benzene molecule, an immense number of isomers are capable of existing. A few of the most important cases may be here elucidated.
Let us, in the first instance, suppose one atom of H to be replaced by another element or radicle. Supposing (as there is at present no reason for doubting) that all the H and C atoms in the molecule have the same property, it is obvious that there can be no isomer of a mono-substitution derivative, and thus we know only one methyl benzene, chlorobenzene, nitrobenzene, &c.
Passing on to the poly-substitution derivatives, we have several possible causes of isomerism. Thus, the H atom may be replaced by one radicle, or more than one H atom may be replaced by an equivalent number of radicles, furnishing the same total number of C and H atoms. For instance—
H H H H
Μ U
∕ ∖ ∕ X
H—C C—H H—C C—H
\=ο Ζ ∖ =c
II II
CH 3 CH 3 h C 2 H 5
Dimethyl-benzone Is metameric with Ethyl-benzene.
Similarly diethyl-benzene is isomeric with tetramethylbenzene, &c., &c. Another cause of isomerism is to be found in the C n H 2n+1 or other radicles replacing the H in the benzene molecule, since these radicals themselves admit of isomeric modifications, as has already been seen when the paraffin series were discussed. Thus, we may have—
C 6 H 5 .C 3 H 7 and C 6 H 5 .CH(CH 3 ) 2 . Propyl-benzene Pseudo-propyl-benzene.
With regard to substitution derivatives of the homologues of benzene, the H atoms may be replaced either in the benzene molecule or in the radicle; thus—
C 6 H 4 Cl.CH 3 C 6 H 5 .CH 2 C1
Methyl-chlorobcnzene or ) . τnptnmpric wi . h (Chloromethyl-benzene monochlorotoluene f ls metameric with j or benzyl chloride.
Another cause of isomerism depends on the relative positions of the replaced H atoms of the benzene molecule. Thus, there can be only one of each mono-substitution derivative, because it is immaterial which of the six H atoms is replaced. Di-derivatives containing two atoms oi the same element or radicle are capable of three isomeric modifications, according as the replaced atoms are those numbered in the graphic formula 1 2 or 1 3 or 1 4. All other arrangements will be found on trial to be reducible to these (X representing the substituted element or radicle)—
Benzene. Ortho-di-derlvative. Meta-di-derivative. Para-di-derivative. C 6 HHHHHH C 6 XXHHHH C 6 XHXHHH c 6 xhhxhh
1 2 3 4 6 6 1 3456 123456 23456
In like manner there can be three tri-or tetra-derivatives, and only one penta-or hexa-derivative, when the H atoms are replaced by the same element or radicle.
When the H atoms are replaced by three or more different elements or radicles the possible number of isomerides is greater than when replaced by the same number of similar elements or radicles, because the arrangements of the latter with regard to one another are then capable of variation. Thus, there can be only three di-derivatives containing two different elements or radicles XY, because the reversal of their order does not affect the relative positions of X to Y. But when three H atoms are replaced by XXY elements or radicles we may have, for every isomeride depending on the relative position of XXY to one another and to the remaining H atoms, corresponding isomerides depending on the order XYX. The same reasoning applies to the higher derivatives.
The following is a list of the hydrocarbons of this series :—
[table]
Other hydrocarbons compounded of benzene residues and C n H 2n-1 , <fcc., radicles are known. Dipropargyl, C 6 H 6 (b.p. 85'C.), is a substance isomericwith benzene, having the formula HO ≡ C—(CH 2 ) 2 —C ≡ CH. The hydrocarbons of the benzene series are all colourless liquids (durene is solid), having peculiar and characteristic odours. By the action of oxidizing agents aromatic hydrocarbons of the general formula C 6 H 6-m (C n H 2n+1 ) m are ultimately converted into acids of the general form C 6 H G _ m (COOH) M . Nitric acid gives rise to the formation of nitro-derivatives C n H 2 „_ (6+w) (NO 2 ) M . Sulphuric acid (concentrated) forms sulphonix acids C n H 2rt _ (6+w) (HSO 3 )„. The halogens produce substitution derivatives C n H 2n _ , the parti cular metameric modification formed being often determined by the temperature at which the halogen is made to act. Benzene, toluene, &c., are formed by heating their corresponding acids with lime: C n H 2n - 7 (COOH) - CO 2 ≈ C n H 2n - 6 .
The remaining series of hydrocarbons are as follows :— [table] [table]
The hydrocarbons benzene (C 6 H 6 ), naphthalene (C 10 H 8 ), anthracene (C 14 H 10 ), chrysene (C 18 H 12 ), and idrialeue (C 22 H 14 ) are noteworthy as forming a homologous series, the successive terms of which differ by +C 4 H 2 . These compounds present certain analogies in their chemical behaviour, and are particularly characterized by their giving rise to quinones when oxidized. In the formation of a quinone H 2 is replaced by O 2 ; thus—
C 10 H 8 + 30 = C 10 H 6 O 2 " + 0H 2 .
Naphthalene. Naphthoquinone.
Pyrene (C 16 H 10 ), although not a member of the series, furnishes a quinone on oxidation. Benzene does not yield a quinone by direct oxidation.
The relationship of the second and third members of the series to benzene will be seen from their graphic formulae :—
H
I
c
Z ∖
H H H H H—C Ο—II
I II I I !l
C CO C O C—H
√'∖ ∕ ∖ ∕ ¼ ∕ ∖ % ∕
H—C C—H H—C C C—H H—C C C
H—C C—H H—C C C—H H—C C C—H
¼ ∕ ¼ ∕ ∖ Z % ∕ ∖ ∕∕
c c c c c
I II II
H H H H H
Benzene Naphthalene. Anthracene.
The hydrocarbons, being the parent compounds of the remaining organic bodies, have been treated in greater detail than is necessary in the case of the other families.
II. Alcohols. Formulation and Classification. —It has already been pointed out that alcohols can be considered as derivatives of hydrocarbons by the substitution of hydroxyl for hydrogen, and, consequently, as compounds of hydroxyl with hydrocarbon radicles (p. 553). Each series of hydrocarbons can thus be supposed to give rise to a corresponding series of alcohols :— [table]
Alcohols containing 1, 2, n semi-molecules of hydroxyl are said to be monohydric, dihydric, n-hydrù. The dihydric alcohols of the ethene series are termed glycols. Alcohols derived from unsaturated hydrocarbons act as unsaturated compounds.
Series C n H 2n+1 . HO.—Some of these alcohols (derived from paraffins) have been previously quoted in illustration of homologous series (p. 552). The first member, methyl alcohol (CH 3 HO), is termed carbinol, and the remaining terms can be derived from this by the replacement of II by C n H 2n+1 radicles. Thus,— [table]
Since the paraffins can be formulated as derivatives of methane (see p. 556), and the present series of alcohols as derivatives of paraffins, it follows that these alcohols can be likewise regarded as methane derivatives :— [table]
It will be convenient to distinguish the carbon atom in combination with hydroxyl as the “typical carbon atom.”
The higher members of the series are susceptible of several kinds of isomerism, which may now be considered. Passing over the first two members, which are not capable of being formulated in more than one way, the third member (propyl alcohol) may be written as—
C. C 2 H 5 . H 2 . HO or C(CH 3 ) 2 H. HO Ethyl carbinol. Dimethyl carbinol.
Thus, the first kind of isomerism consists in the replacement of one C n H 2n+1 radicle by two others, each containing a smaller number of carbon atoms. The fourth member of the series (tetryl or butyl alcohol) may be written :—
C.C 3 H 7 .H 2 .HO C(CH 3 )(C 2 H 5 )H,HO C(CH 3 ) 3 .HO
Propyl carbinol. Methyl-ethyl carbinol. Trimethyl carbinol.
The next kind of isomerism, as exemplified by the third of the above formulae, consists in the substitution of three C n H 2n+1 radicles for the three hydrogen atoms of carbinol.
Now, as carbinol contains but three replaceable hydrogen atoms (not regarding the atom contained in the hydroxyl, the replacement of which gives rise to a new class of compounds), trimethyl carbinol is the type of a series of alcohols, in which the number of C n H 2n+1 radicles combined with the typical carbon atom is a maximum. The isomeric alcohols of the present series can thus be conveniently classified under the three following groups :—
1. Primary alcohols, in which the typical carbon atom is combined with but one other carbon atom—
C(C n H 2n+1 )H 2 ,HO.
2. Secondary alcohols, in which the typical carbon atom is combined with two other carbon atoms—
C(C n H 2n+1 ) 2 H.HO.
3. Tertiary alcohols, in which the typical carbon atom is combined with three other carbon atoms—
C(C n H 2n+1 ) 3 ∙HO.
Another cause of isomerism among alcohols is to be found in the constitution of the C n H 2n+1 radicles (see also p. 550). Thus we may have—
C.C 3 H 7 .H 2 .HO isomeric with C.CH(CH 3 )2.H 2 .HO Propyl carbinot Pseudopropyl carbinol.
C(CH 3 )(C 4 H 9 )H.HO „ C.CH 3 [CH 2 .CH(CH 3 ) 2 ]H.HO Methyl butyl carbinol. Methyl isobutyl carbinol.
Each of the three above-named groups is thus capable of a further division into two sub-groups, such as normal primary and iso-primary, normal secondary and iso-secondary, <kc., &c. In the higher members a further number of isomeric modifications become possible. Many of these isomeric alcohols are known, but a large number have yet to be discovered to complete the series. The following is a list of the normal primary alcohols :— [table]
These alcohols (up to nonyl alcohol) are limpid liquids, the viscidity of which increases with the molecular weight. The three other members are white crystalline solids. Methyl, ethyl, and propyl alcohols are readily miscible with water, the remaining members are more or less of the nature of oils. Most of these alcohols possess peculiar and characteristic odours. The boiling-points of the isomers are generally lower than those of the corresponding normal alcohols.
The normal and iso-primary alcohols can be obtained from the corresponding haloid paraffin derivatives by the action of potassium hydroxide :—
Normal primary.. C,,H 2n+1 Zfo + KHO = C n H 2n+1 .HO + KZfα.
Iso-primary (CH 3 ) 2 CH(CH 2 ) n .ZΓα + KHO
= (CH 3 ) 2 CΠ(CH 2 ) n . HO + KHa.
Likewise by the action of nascent hydrogen on the corresponding normal primary and iso-primary aldehydes :—
(C n H 2n+1 )COH + H 2 = C n H 2n+1 .CH 2 . HO.
The synthesis of alcohols from the corresponding olefines has already been alluded to (p. 559). Since ethine can be formed directly from its elements (p. 559), and ethene from ethine (p. 559), it is obvious that ethyl alcohol can by these means be formed synthetically.
Secondary alcohols are produced by the action of nascent hydrogen on ketones :—
CO (C n H 2n+1 ) 2 + H 2 = C(C n H 2n+1 ) 2 H.HO.
Tertiary alcohols are prepared by the action of organo-zinc compounds on the chlorides of acid radicles, C n H 2n+1 .COCl, and treatment of the product with water.
The metamerism of the three groups of alcohols is strikingly displayed by their behaviour on oxidation. Thus, primary or iso-primary alcohols are oxidized first to aldehydes—
(C.H 2 ,, 41 )CH 2 .HO + 0 ≈ (C⅛ +1 )COH + OH 2 ;
Alcohol· Aldehyde.
then to acids containing the same number of carbon atoms—
(C n H 2n+1 )COH + 0 = C n H 2n+1 .COOH.
Aldehyde. Acid.
Secondary alcohols are oxidized first to ketones—
C(C n H 2n+1 ) 2 H.HO + 0 = CO(C n H 2n+1 ) 2 + OH 2 ;
Alcohol· Ketone.
then to one or more acids containing a smaller number of carbon atoms than the alcohol—
CO (C n H 2n+1 ) 2 + 3O = C n H 2n+1 .COOH + C n H 2n O 2 . Ketone. Acid. Acid.
Tertiary alcohols are probably first oxidized to ketones, and ultimately to a mixture of two or more acids containing a smaller number of carbon atoms than the alcohol.
By particular reactions an alcohol of one group can be converted into an isomer of another group, such, for instance, as normal primary butyl alcohol into the isoprimary alcohol, and the latter into the tertiary alcohol.
In many of their reactions alcohols are the analogues of water; for instance— [table]
In other reactions they resemble metallic hydroxides [table]
Of this series the best known members are methyl and ethyl alcohols. The former (or “wood-spirit ”) is chiefly obtained from the crude “wood vinegar” produced by the destructive distillation of wood. Ethyl alcohol is common “spirits of wine,” and is always procured by the fermentation of sugar, which, in contact with yeast, undergoes the following change—C 6 H 12 O 6 = 2CO 2 + 2C 2 H 5 .HO.
Many of the higher members are formed by the fermentation of grain, sugar-beet molasses, potato, <fcc. Fusel oil consists chiefly of isomeric modifications of amyl alcohol [table]
Series C n H 2n-7 .HO, Phenols. — The alcohols of this series bear the same relation to the benzene hydrocarbons that the ethyl series of alcohols bear to the paraffins. Thus these alcohols at once divide themselves into two metameric series according as the hydroxyl replaces hydrogen in the benzene or paraffin radicle; for instance, methylbenzene or toluene (C 6 H 5 .CH 3 ) gives rise to—
C 6 H 4 (HO).CH 3 and C 6 H 5 .CH 2 (HO). cresol. Benzyl alcohol.
The alcohols of the benzylic series ^C 6 H 6 - w ∣
are true alcohols in their chemical behaviour, being formed from the haloid derivatives of their corresponding hydrocarbons by the action of potassium hydroxide, and being oxidized to acids containing the same number of carbon atoms. The alcohols represented by cresol, C 6 H 5 - w HO.(C n H 2n+1 ) m , or phenols, on the other hand, act more like saturated compounds, and in some respects resemble acids. They cannot be directly obtained from their corresponding haloid hydrocarbon derivatives, and are converted by oxidation into quinones. The phenols yield substitution-derivatives by the action of halogens, nitric acid, &c.; such, for example, as— [table]
The phenols yield sulphonix acids when acted on by strong sulphuric acid, C 6 H 4 (HO)(HSO 3 ), C 6 H 8 (HO)(HSO 8 ) 2 , &c.; the metameric alcohols do not undergo an analogous reaction.
The homologues of phenol being di-derivatives of benzene are susceptible of the isomeric modifications already pointed out (p. 560). Thus there are known— [table]
Isomerism analogous to that exhibited by the derivatives of benzene and its homologues likewise extends to aU the derivatives of phenol and its homologues.
The following is a brief list of the compounds of the present series :— [table]
The boiling-points of the alcohols are generally higher than those of the metameric phenols.
Many of the compounds of the present series are transparent oily liquids possessed of powerful odours; others are white crystalline solids. Of the phenols the best known is phenol or carbolic acid, obtained from the acid portions of coal-tar.
Series C n H 2n _ 9 .HO comprises cinnamic alcohol, C 9 H 9 .HO, and cholesterin, C 26 H 43 ,HO.
Series C n H 2 „_ ]3 .HO.—Only one member is known, viz., naphthol, C 10 H 7 .HO.
Series (C n H 2n )"(HO) 2 , Glycols. — Of these dihydric alcohols the terms corresponding to the first six members of the olefine series are known, and are named ethene glycol, propene glycol, &c. The glycols are colourless, viscid liquids, miscible with water. They are generally prepared by acting on di-haloid olefine compounds with silver acetate: C n H 2n Br 2 + 2AgC 2 H 3 O 2 = C n H 2n (C 2 H 3 O 2 ) 2 + 2AgBr The resulting acetate is then heated with potassium hydroxide: C n H 2n (C 2 H 3 O 2 ) 2 + 2KHO = C n H 2n (HO) 2 + 2KC 2 H 3 O 2 . The glycols are capable of existing in isomeric forms similar to those of analogous olefine compounds (see p. 559), but these isomers are at present but imperfectly known. A series of compounds known as p>olyethxnic glycols are formed from ethene glycol by condensation, with elimination of water. Their general formula is re{C 2 H 4 (HO) 2 } - (n- 1)(OH 2 ), where the highest value of n is 6.
Series C n H 2n-8 (HO) 2 , Dihydric Phenols. — The following are known :— [table]
Many of these compounds can be formed by suitable reactions from benzene hydrocarbons, phenols, &c.; others are found ready formed in lichens, or can be prepared from substances of vegetable origin.
Series (C n H 2n-1 )'"(HO) 3 comprises propenyl alcohol or glycerin, C 3 H 5 .(HO) 3 , and amyl-glycerin, C 5 H 9 .(HO) 3 . Glycerin can be formed synthetically, but is generally obtained from glycerides, a group of ethereal salts composed of glycerin and acids of the acetic and oleic series (see p. 572), which exist ready formed in most animal and vegetable oils and fats. Thus stearin (propenyl tristearatc), a constituent of many natural fats, when acted upon by super-heated steam, decomposes as shown in the equation— C 3 H s (C i8 H 35 O 2 ) 3 + 3OH 2 - C 3 H 3 (HO) 3 + 3H.C 1s ¾O 2 .
Stearin. Glycerin. Stearic acid.
Large quantities of glycerin are obtained by this means. Glycerin is a colourless, transparent, viscid liquid of a sweetish taste, readily soluble in water, and boiling (under 50 mm. pressure) at 210° C. Pólyglycerins analogous to the polyethenic glycols are known.
Series (C n H 2n - 9 )"'(HO) 3 .— Trihydric phenols comprising the isomerides pyrogallic acid (or pyrogαllol) and phloro-glucin, C 6 H 3 (HO) 3 .
Series (C n H 2n-2 ) lτ (HO) 4 .— Erythrite, erythromannite or phycite, C 4 H 6 (HO) 4 , and prοpyl-phydte, C 3 H 4 (HO) 4 .
Series (C n H 2n - 5 ) τ (HO) 5 .— Pinite and quercite, C 6 H 7 (HO) 5 .
Series (C„H 2 „_ 4 ) t, (HO) 6 . — Mannite and dulcite,
C 6 H 8 (HO) 6 , two natural sugars.
Sugar, starch, gum, &c., are members of a group of compounds termed carbohydrates , which are closely allied to the present series of alcohols. The names, classification, and relationship of these bodies are shown in the following table :— [table]
The glucoses exhibit the characters both of aldehydes and of alcohols; the saccharoses bear the same relationship to the glucoses that the polyethenic glycols bear to glycol; the amyloses are formed from the saccharoses by the elimination of one molecule of water.
Many of the sugars are widely diffused throughout the vegetable kingdom—the sugar in ordinary use being obtained from the sugar-cane and sugar-beet. Dextrose and levulose are found in honey, ripe fruits, &c., lactose in milk, inosite in flesh. Gum-arabic is the dried exudation of certain species of Acacia. All the amyloses, with the exception of glycogen, which is found in the liver of animals, are of vegetable origin. Cellulose or lignin is the chief constituent of the cells of plants; starch is found in the cells of most plants. Gun-cotton or pyroxylin is obtained from cotton-wool, which is almost pure cellulose, by the action of strong nitric acid; it has the composition of trinitro-cellulose, C 6 H 7 (NO 2 ) 3 O 5 . Many of the varieties of the sugars are physical isomerides differing from each other only in their rotatory action on polarized light.
III. Haloid Ethers.
These compounds can be formed from saturated hydrocarbons, ⅛and from hydrocarbons which behave as such by the direct action of the halogen: —
C n H 2n+2 + Cl 2 = C⅛ +1 C1 + HC1.
The substitution can in many cases be continued atom by atom, till the whole of the hydrogen is replaced; thus :— [table]
The reactions are often accelerated and in some cases remarkably modified by light or heat. The presence of iodine, antimony, <kc., frequently assists the action of chlorine. Chlorine is more energetic than bromine, and the action of iodine less energetic than that of bromine,—iodine, in fact, does not act directly on the paraffins. By the direct action of halogens on paraffins a mixture containing more than one derivative is generally obtained. Thus, when chlorine and methane are mixed and exposed to diffused light (explosion takes place in direct sunlight), the whole series of derivatives CH 3 C1, CH 2 C1 2 , CHC1 3 , and CC1 4 are formed.
Unsaturated hydrocarbons first of all combine directly with a certain number of halogen atoms. In some cases the number of halogen atoms thus taken up restores the compound to the type of a saturated body:—
C 2 H 4 + Cl 2 = C 2 H 4 Cl 2 = ethene dichloride = dichldrethane. Ethene.
In other cases the restoration is not to the type of a saturated hydrocarbon :—
C 10 H 8 + 2Cl 2 =C 10 H 8 Cl 4 , tetrachlorinated naphthalene, corresponding ¾ 3 tha - [to type C n H jn - 8 .
The mono-haloid derivatives of paraffins and some other hydrocarbons are conveniently prepared by the action of haloid acids on the corresponding alcohols :— [table]
Polyhydric alcohols may by these reactions have their "hydroxyl only partially replaced by halogens :—thus, by the action of hydrochloric acid on glycerin there are produced C 3 H 5 C1(HO) 2 (phlorhydrin) and C 3 H 5 C1 2 (HO) (dixhlor-hydrin).
The haloid phosphorus compounds act upon alcohols in a manner similar to that of the halogen acids :— [table]
The haloid ethers are liable to the same isomeric modifications as their parent hydrocarbons and derivatives. The conditions which determine the production of any particular modification cannot in all oases be laid down with precision. For instance, the di-haloid paraffin derivatives obtained from the paraffins are in some cases identical and in other cases isomeric with the corresponding compounds produced by the direct union of halogens with olefines. Again, the higher members of the paraffin series when acted on by chlorine yield simultaneously two isomeric (primary and secondary) mono-chlorinated derivatives. As a final example, when chlorine acts upon toluene in the cold chloro-toluenes (C β H 4 Cl.CH 3 , C 6 H 3 C1 2 .CH 3 , &c.) are formed, but when the action takes place with boiling toluene the substitution takes place in the methyl radicle, thus forming benzyl compounds (C 6 H 5 .CH 2 C1, &c.)
Alcohols are reproduced from their haloid ethers by the action of KHO (see p. 563), or by first converting the haloid ether into an acetate of the radicle by treatment with silver acetate, and then acting on the acetate with KHO. Tho alcohols thus formed are similar in isomeric constitution to the haloid ethers from which they are obtained. For instance, (CH 3 ) 2 CH.I (pseudopropyl iodide) when made to undergo the reactions specified is converted into (CH 3 ) 2 CH.HO (pseudopropyl alcohol) (see also p. 563). Isomeric haloid ethers are thus named in accordance with the radicle composing the alcohols from which they are obtained, or to which they give rise on conversion into acetates and treatment with KHO; thus, we have the following monoiodobutanes :— [table]
The same relations exist between the haloid ethers and their isomerides as between the paraffins and their isomerides. Thus, the boiling points and specific gravities of the normal ethers are higher, and their chemical stability is greater than that of the isomerides.
By the action of nascent hydrogen many haloid ethers are restored to the hydrocarbons from which they are derived by an inverse substitution. Thus tetrachlormethane (CC1 4 ) when treated with sodium amalgam and water is ultimately converted into methane (CH 4 ), monochlorbenzene (C 6 H 5 C1) into benzene, &c. The action of caustic alkalies upon mono-haloid ethers will be best illustrated by the following examples :— [table]
The mono-haloid olefines (typified by monochlorethene in the last equation) are obtained by the action of a caustic alkali on di-haloid paraffins (see p. 558). These mono-haloid olefines, hke their parent olefines, are dyad radicles combining with Ha i and H7Zα. Their compounds with haloid acids are under some conditions identical and under other conditions isomeric with the corresponding compounds produced by the direct combination of the olefine with two halogen atoms.
Of the haloid ethers derived from paraffins the best known is trichloromethane or chloroform (CHC1 3 ), a colourless, transparent, sweet-tasting, and somewhat fragrant liquid boiling at 62°, largely used as an anaesthetic. It can be prepared by the direct action of chlorine on methane, but it is usually obtained "⅜y the action of bleaching-powder on common alcohol. The corresponding bromine and iodine compounds, CHBr 3 and CHI 3 lfromo∙ form and iodoform), and CH(NO 2 ) 3 (nitroform), are known.
Cyano-derivatives of Hydrocarbons. — The metamerism exhibited by these compounds has been already referred to (p. 555), andjsince they resemble in many respects the haloid ethers, their modes of preparation may be now touched upon.
The most general method of obtaining the cyanides (nitriles) in a state of purity is the dehydration of the ammonium salts of acids containing the corresponding hydrocarbon radicles (see also p. 555).
C 6 H δ .CO(ONH 4 ) - 2OH 2 = C 6 H 5 .CN Ammonium benzoate. Phenyl cyanide.
Most of the other processes yield a mixture of cyanides and isocyanides :— [table]
In the first of these methods the cyanide predominates, and in the second the isocyanide. By distilling a mixture of potassium cyanide with an ethereo-metallic salt of sulphuric or sulphurous acid, a mixture of cyanide and isocyanide of the radicle contained in the ethereal salt is obtained, the cyanide or nitrile being in excess :— [table]
The cyanides or nitriles do not exhibit strongly marked basic properties, nor are they oxidized by Ag 2 O or HgO. They are not poisonous, and do not possess unpleasant odours. Their boiling-points are generally higher than those of the metameric isocyanides. The nitriles are susceptible of isomeric modifications depending on the isomerism of their contained hydrocarbon radicles.
IV. Ethers.
These compounds bear to the alcohols the same relations that the metallic oxides bear to the hydroxides :— [table]
In accordance with this view, many ethers are formed by dehydrating their corresponding alcohols by means of strong sulphuric acid, zinc chloride, boron trioxide, &c. :— [table]
Conversely, many ethers, by prolonged heating with water, are retransformed into their corresponding alcohols.
Oxygen being a diatomic element combines with two monatomic radicles. When the two radicles are similar the compound is termed a simple ether, when dissimilar a compound ether. Thus, the ethers formulated above (methyl and benzyl) are simple ethers; while CH 3 .O.C 2 H δ , C 5 H 11 .O.CH 3 , C 6 H δ .O.C 3 H δ
Methyl-ethyl ether Methyl-amyl ether. Phenyl-allyl ether, are examples of compound ethers. Compound ethers containing a C n H 2n+1 and a C n Ξ 2 ^ 7 radicle are termed anisols.
Ethers are liable to isomeric modifications dependent on the isomerism of their contained radicles, and likewise to metamerism depending on the presence of different radicles. For example, C 2 H 5 .O.C 4 H a (ethyl-butyl ether) would be metameric with (C 3 H 7 ) 2 O (propyl ether).
A general method of preparing ethers (simple and compound) containing monatomic radicles is to act upon the sodium derivative of the alcohol containing the one radicle with the haloid ether containing the other radicle. When the alcohol and haloid ether contain the same radicles, the result is a simple ether; when they contain different radicles a compound ether is produced :— [table]
The ethers derived from the phenols are prepared by special methods. Ethers containing diatomic radicles are obtained from the mono-haloid derivatives of the dihydric alcohols (glycols) by the action of KHO :—
C 2 H 4 C1(HO) + KHO = C 2 H 4 O + KC1 + OH 2 . Monochlorhydrin. Ethene ether.
Compound ethers can be obtained from the sodium derivatives of the glycols by the action of mono-haloid ethers :— [table]
The ethene series of ethers are more active in their chemical behaviour than those containing C n H 2n+1 radicles.
Ethers corresponding to the first seven normal primary alcohols of the series C n H 2n+1 . HO are known. Of these “sulphuric ether” is most frequently met with. It is the oxide of ethyl (C 2 H 5 ) 2 O, and is prepared by allowing a stream of ethyl alcohol to flow into a mixture of strong sulphuric acid and alcohol kept at a constant temperature of about 140° C. The following reactions occur:— [table]
As will be seen from these equations, a given quantity of H 2 SO 4 is theoretically capable of etherifying an unlimited quantity of alcohol; practically, however, a limit is reached. If a mixture of alcohols is employed at starting, a compound ether is produced.
Ethyl ether is a colourless, mobile liquid, but very slightly soluble in water, and possessing an agreeable odour. It boils at 35° ’ 5 C. By the action of chlorine it yields substitution derivatives, C 2 H 4 C1∙O.C 2 H 6 ,
(C 2 C1 5 ) 2 O.
V. Sulphur, Selenium, and Tellurium Alcohols and Ethers.
These compounds are the analogues of the alcohols and ethers (see p. 553), and are in many cases prepared by analogous methods. The following are typical reactions :—
Preparation of Sulphur Alcohols. [table]
Preparation of Sulphur, dec., Ethers. [table]
Diethyl teUuride is prepared by an analogous reaction, K 2 Te being substituted for K 2 Se. Polysulphides, (C 2 H 5 ) 2 S 2 , (C 2 H 5 ) 2 S 3 , can be obtained by employing alkaline polysulphides in these reactions. Certain thio-alcohols can be prepared from the corresponding alcohols by the actioa
The cyanide thus formed unites with AgCN, forming the compound CH 3 . CN , AgCN, which, on treatment with KCN in excess, yields the methyl isocyanide with the formation of KAg(CN) 2 .
of phosphorus pentasulphide (5C 2 H 5 .HO + P 2 S 5 = 5C 2 H 5 .HS + P 2 O 5 ). The sulphur analogues of the phenols and dihydric phenols are prepared by a special general method. The sulphonic acid (see p. 561) of the corresponding hydrocarbon is converted into a sulphonic chloride by the action of PCl 5 , and the sulphonic chloride is then acted on by nascent hydrogen :— [table]
The sodium derivatives of these thio-phenols yield compound thio-ethers by the action of haloid ethers [table]
Compounds intermediate between alcohols and thioalcohols are known :—
C 3 H 5 (HO) 2 C1 + KHS ≈ C 3 H 5 (HO) 2 (HS) + KC1.
The compounds of this family are mobile or oily liquids or crystalline solids. Most of them possess characteristic and offensive odours. They are susceptible of the same isomerism as their oxygen analogues. The thio-alcohols combine energetically with alkali metals, and with certain metallic oxides and salts, to form derivatives analogous to the metallic alcohol derivatives,—hence the name mercaptans (corpus mercurio aptum) sometimes given to these bodies. Certain lead mercaptides, when heated in the dry state, yield the corresponding thio-ethers.
VL Aldehydes.
The relations between these compounds and the primary alcohols are clearly brought out by the mode of formulation adopted for these alcohols (p. 562). Thus :— [table]
It will be seen from these formulae that aldehydes are derived from alcohols by the elimination of H 2 from the group CH 2 OH, i.e., 2 hydrogen atoms are withdrawn from the typical carbon atom, leaving the group (COH)' (compare with definition previously given, p. 553); hence the generic name (alcohol dehydrogenatum). It is thus possible for any alcohol to furnish an aldehyde, although great numbers of these compounds have yet to be discovered to complete the various series. The following formulae will illustrate the derivation of aldehydes from polyhydric alcohols :— [table]
The aldehydes of monohydric alcohols are metameric with the oxides (ethers) of dyad radicles. Thus, ethyl aldehyde, CH 3 .COH, is metameric with ethene oxide (C 2 H 4 )"O; benzoic aldehyde would be metameric with methyl-phenylene oxide [C 6 H 3 (CH 3 )]"O. The aldehydes are liable also to the isomerism of their contained hydrocarbon radicles, while aromatic aldehydes of dihydric alcohols are susceptible of the isomerism incident to the relative positions of the COH groups.
By the action of oxidizing agents aldehydes are converted into acids containing the same number of carbon atoms; thus, R' being the monatomic radicle—
R' R'
I +O=∣
COH COOH
Aldehyde. Acid.
Starting, therefore, from the parent hydrocarbons, the primary alcohols are the first results of the introduction of hydroxyl into the methyl group. The next step in oxidation removes the two remaining hydrogen atoms from this group with the formation of aldehydes, and the final result of the oxidation is to convert COH into carboxyl COOH, with the formation of acids. Aldehydes thus occupy a position intermediate between alcohols and acids :— [table]
Aldehydes take up nascent hydrogen, reproducing alcohols—R'.COH + H 2 = R'.CH 2 OH. The acid sulphites of the alkali metals combine directly with aldehydes, forming crystalline compounds, which, on treatment with a mineral acid, yield the aldehyde unaltered.
Aldehydes corresponding to eleven of the primary alcohols of the C n H 2n+1 .HO series are known, and are generally prepared by the oxidation of these alcohols: R'.CH 2 OH - H 2 = R'.COH. The contained C n H 2n+1 radicles are in some cases normal primary, and in others iso-primary. The first member, methyl or formic aldehyde, H.COH, is gaseous; the succeeding terms are liquid, and hexdecyl aldehyde is a crystalline solid. The liquid aldehydes are colourless, transparent, and possessed of pungent ethereal odours; their mobility decreases, and their boiling-points rise as the series is ascended. Their solubility in water decreases in the same manner. By the action of PC1 5 aldehydes lose their oxygen, and Cl 2 is substituted. These chloraldehydes are isomeric with the dichlorides of the olefines, and (with the exception of the ethyl compound), with the chlorides of monochlorinated C n H 2n+1 radicles. Thus :— [table]
(See also p. 559.) The aldehydes of the present series combine directly with ammonia, forming aldehyde-ammonias of the general formula R'. C(OH)(NH. 2 )H, which, by losing the elements of water, condense into basic substances, termed aldines and oxaldines. When acted on by chlorine in large excess these aldehydes yield the corresponding chlorides of acid radicles (R'CO)'CL Trichloraldehyde, CC1 3 ∙COH, or chloral, can be obtained by the action of chlorine on ethyl alcohol· The acetals are compounds formed by the combination of aldehydes and alcohols of the C n H 2n+1 HO series, with elimination of water. By the action of zinc chloride, &c., acetic aldehyde yields crotonic aldehyde : 2C 2 H 4 O - OH 2 = C 4 H 6 O. The aldehydes oxidize with the greatest readiness to the corresponding acids; thus, silver oxide is reduced to the metallic state when heated with an aldehyde and water. A drop of acetic aldehyde let fall on blue litmus paper shows an acid reaction on mere exposure to the air, owing to the formation of acetic acid. Aldehydes are characterized by their extreme readiness to undergo polymeric modification. Thus the presence of a trace of certain reagents converts acetic aldehyde under some conditions into paraldehyde, C 6 H 12 O 3 = 3C 2 H 4 O, and under other conditions into metaldehyde, n(C 2 H 4 O).
The aldehydes corresponding to the C n H 2n ~ 1 .HO alcohols are acrylic aldehyde or acrolein C 2 H 3 .COH, and cratonic aldehyde, previously mentioned. These aldehydes do not yield compounds analogous to aldehyde-ammonia.
Of the aldehydes of the aromatic series mention may be made of benzoic aldehyde, which can be obtained from bitter-almond oil; cumic aldehyde, existing in oil of cumin; salicylic aldehyde, which exists in the flowers of “meadowsweet;” and cinnamic aldehyde, the chief constituent of the oils of cassia and cinnamon. The aromatic aldehydes, by the action of ammonia, are converted into hydramides of the general formula (C n H 2n .-CH) 3 "N 2
VII. Ketonhs.
These compounds are derived from secondary alcohols, in the same manner that aldehydes are derived from primary alcohols :—
(0⅛ +1 ) 2 ch.oh - h 2 ≡ (C⅛ +1 ) 2 CO. Secondary alcohol. Ketone.
The ketones may thus be regarded as compounds of C n H 2n+1 , <fcc., radicles with CO, or as aldehydes in which the H of the COH group is replaced by such radicles. The compounds of this family bear considerable resemblance to the aldehydes, but are distinguished by their behaviour on oxidation, for whereas the aldehydes are readily converted into acids containing the same number of carbon atoms, the ketones are converted (with some difficulty) into a mixture of two acids of the C n H 2n+1 .COOH series, each containing a smaller number of carbon atoms. The law of the oxidation of ketones appears to be that the less complex of the two hydrocarbon radicles remains attached to the CO; thus— [table]
Ketones are converted by the action of nascent hydιogen into secondary alcohols (see p. 563).
The ketones are liable to isomeric modification depending on the isomerism of their contained hydrocarbon radicles.
{CH
r , t? (methyl propyl ketone) is iso- ½ n 7
í cττ
meric with CO ∣ qjj¢qjj ) (methyl-pseudopropyl ketone). Furthermore, there can be metamerism among ketones owing to the presence of different radicles in the molecule, and every ketone is metameric with an aldehyde of the same series. Thus, the following are metameric :—
∞{c"⅛ c0 {¾¾ C s H 11 ,COH.
Metħyl-butyl Et, ’ yl X° pyl Hexyl aldehyde ketone. ketone. j j
Ketones corresponding to the general formulae CO(C λ H 2h+1 ) 2 , ∞{c⅛*√ and CO(C n H 2n . r ) 2 are known, and are formed by various methods, of which the following are the most important :—
1. By the oxidation of secondary alcohols (see beginning of section).
2. By the action of sodium organo-metallic bodies on CO: 2Na(C n H 2n+1 ) + CO = CO(C n H 2n+1 ) 2 + Na 2 .
3. By the action of zinc organo-metallic compounds on acid chlorides (chlorides of acid radicles) :—
2CO(C n H 2n+1 )Cl + Zn(C n H 2n+1 ) 2 = 2CO(C n H 2n+1 ) 2 + ZnCl 2 .
4. By the dry distillation of the Ca and Ba salts of monobasic acids of the series C n H 2n+1 .CO 2 H, C n H 2n - 7 .CO 2 H, &c.,— [table]
By employing a mixture of the salts of acids belonging to two different series, or of two different acids belonging to the same series, ketones containing two different radicles are obtained.
The ketones of the series CO(C n H 2rt+1 ) 2 and CO ∣ ^ n g 2n+1 are, with few exceptions, mobile or oily colourless liquids, possessed of most characteristic and penetrating odours; those of the series CO(C n H 2π-7 ) 2 are crystalline solids. Ketones containing methyl form with acid sulphites white crystalline compounds, from which the ketone is obtained unaltered on distillation with an alkali. The best known ketone of the series CO(C n H 2n + 1 ) 2 is dimethylketone or acetone, CO(CH 3 ) 2 , a limpid, inflammable liquid, boiling at 55 0, 5 C., and readily miscible with water. In addition to the general modes of formation previously given, this ketone can be obtained by the destructive distillation of citric acid, and also by distilling certain carbohydrates with quicklime. Heated with ammonia, acetone forms acetσnvne, a basic substance, of the formula (C 3 H 6 )" 3 N 2 . Of the aromatic ketones methyl-phenyl ketone or acetophenone is interesting as furnishing, by the action of fuming nitric acid, a nitro-derivative,
í CH
CO 1 ∏ ττ ∕∙∖τn ∖ > which, when heated with soda-lime and ( ...
zinc dust, yields indigotin or indigo-blue — [table]
VIII. Organic Acids.
The relationship of the organic acids to the hydrocarbons and to the alcohols has been previously pointed out (pp. 553 and 567). A further development of this relationship, as bearing on the formation of acids from polyhydric alcohols, is shown in the following examples :— [table]
In the formation of acids from alcohols, therefore (disregarding the intermediate formation of aldehydes), one atom of oxygen is substituted for H 2 in the group CH 2 OH, thus converting this group into carboxyl, COOH. It has been before mentioned that CH 2 OH is derived from methyl by the substitution of HO for H, so that carboxyl may be regarded as a methyl derivative, and a similar view may be extended to cyanogen CN, where N" z may be regarded as replacing H 3 . This connection between CN and COOH is shown by the various reactions in which the one radicle is converted into the other, but more particularly, so far as the compounds now under consideration are concerned, by the synthesis of organic acids from the corresponding nitriles (p. 555). A few comparative formulae will serve to illustrate still further this important relationship :— [table]
Classification. — By inspecting the formulae given at the commencement of the present section, it will be seen that the number of times COOH occurs in an acid is dependent on the number of times the group CH 2 .OH occurs in the parent alcohol. Now, organic like mineral acids when brought into contact with metallic oxides or hydroxides, give rise to the formation of salts, in which H, H 2 , H 3 , &c., are replaced by M', M' 2 or Μ", M' 3 or M'", &c.; but the hydrogen thus replaceable is found to be not that which enters simply into the composition of HO, as with mineral acids, but that which exists in the group COOH. The basicity of an organic acid depends therefore on the number of carboxyl groups it contains—acids containing 1, 2, 3, n carboxyl groups being mono-, di-, tri-, n-basic.
Monobasic Acids. Series C m II 2 , +1 (CO 2 H), Acetic or Fatty Series. — These acids are prepared—1. By the oxidation of the corresponding primary alcohols and aldehydes (pp. 563 and 567); 2. By the action of mineral acids or alkalies on the cyanoderivatives of the paraffins obtained by the methods previously given (p. 566; see also p. 555); 3. By the action of organo-sodium compounds on carbon dioxide :—
C w H 2 , +1 Na + CO 2 = C⅛ tl .COONa Sodium salt.
Isomerism. — Just as the alcohols of the methyl series are most conveniently formulated as derivatives of the first term carbinol, so the present acids may be regarded as derived from formic or acetic acids, the two first members of the series :—
H CH 3 CH 2 (CH 3 ) C‰ CH 2 ((‰)
COOH COOH COOH COOH COOH
Acetic or Propionic or r . nrτ 1 r -- α l Formic acid. methyl-formic methacetie f ξ r ⅛ c ⅛ d . ¾ d , acid. acid.
The first kind of isomerism which we have to consider is that depending on the replacement of one C.H 2 , +1 radicle by two or three others, each containing a smaller nuπft>er of carbon atoms. (Compare with isomerism of alcohols, p. 562.) Thus we may have the following isomerides :— ch 2 (c 3 h,) ch(ch 3 )(c 2 h 5 ) c(ch 3 ) COOH COOH COOH
Propacetic or Methyl-ethacetic Trimethacetic valeric acid. acid. acid.
Acids are thus classified in the same manner as the alcohols, into primary, secondary, and tertiary, according as the carbon atom in combination with carboxyl is combined with one, two, or three other carbon atoms.
Another kind of isomerism is dependent on the isomerism of the radicles replacing the H of the CH 3 of the acetic acid; thus—
C(C 3 H 7 )H 2 . COOH C[CH(CH 3 ) 2 ]H 2 . COOH
Propacetic acid. is isomeric with Pseudopropacetic acid.
Each of the three groups of acids is accordingly subdivided into normal and iso-acids, although the series are far from being complete.
The following is a list of the normal primary acids :— [table]
These acids exhibit the general physical properties of all homologous series. Up to pelargonic acid they are limpid liquids, possessing penetrating odours, and of increasing viscidity; the remaining members are solid fats. The acids of this series are found in nature in the free state (formic acid in ants, cerotic acid in beeswax, <fcc,), or as glyceric, &c., ethereal salts, in natural fats (mutton fat, glyceric stearate; spermaceti, cetylic palmitate, &c.); hence the term “fatty series.” Formic acid can be prepared by certain special methods—1. By passing carbon monoxide over moist potash (CO + KHO = H.COOK); 2. By heating dry oxalic acid with glycerine (H 2 C 2 O 4 = CO 2 + H.COOH). Acetic acid also can be synthesized from ethine and oxygen in presence of potash (C 2 H 2 + O-t-KHO = C 2 H 3 KO 2 ). This acid is obtained on the large scale by the dry distillation of wood, and is likewise contained in vinegar. Butyric acid is formed by the fermentation of sugar in presence of chalk and putrefying cheese or milk. The metallic salts formed by the fatty acids furnish the organic acids unchanged on distillation with a mineral acid. The potassium salts, when distilled with potassium formate, yield the corresponding aldehydes—
C⅛ +1 .COOK + HCOOK = C⅛ +1 .COH + CO(OK) 2 Potassium salt Formate Aldehyde.
The aldehydes can be converted into the alcohols by the action of nascent hydrogen, and from the alcohols the haloid ethers can be obtained, which, by methods now known to the reader, can be converted into nitriles, and these again into acids containing one more atom of carbon than the alcohol (p. 555). It will be obvious that by these reactions the homologous series of fatty acids can be ascended term by term. By the action of chlorine, &c., on the fatty acids substitution-compounds are obtained. Thus acetic acid is converted into CH 2 CLCO 2 H, CHC1 2 .CO 2 H, and CC1 3 .CO 2 H. Other derivatives can be obtained from these haloid derivatives by double decomposition; for example, (by the action of KCN), CH 2 (CN).CO 2 H, cyan-axetic arid; (by the action of ammonia), CH 2 (NH 2 ).CO 2 H, amidacetic acid or glycorine. Series C n H 2n (HO).CO 2 H, Lactic Series. — These acids are the hydroxyl derivatives of the fatty acids, from which they can be obtained by treating the mono-haloid substitution derivatives with Ag 2 O and water—
2C⅛ffα. CO 2 H + Ag 2 O + OH 2 =2C n H 2n (HO). CO s H + 2AgZZ<z.
The isomeric modifications of these acids admit of being grouped under four divisions :—
1. Primary Acids, C(C n H 5n+1 )H.HO.CO 2 H.—These are formed by the above reaction from primary fatty acids, and also by the slow oxidation of glycols of the form CH(C n H 2 n+ 1 )OH.CH 2 .OH.
Aldehydes of the acetic series form compounds with HCN, which, on heating with HC1 and water, yield acids of this group :—
CH 3 . COH + HCN + 2OH 2 + HC1=C(CH 3 )H. HO. CO 2 H + NH 4 C1. Acetic aldehyde. Ethidene-lactic acid.
2. Secondary Acids, C(C n H 2 n-L 1 ) 2 .HO.CO 2 H.—Produced by the action of Ag 2 O and water on mono-haloid derivatives of secondary fatty acids. Also by the action of organo-zinc compounds on ethyl oxalate, C 2 O 2 (OC 2 H s ) 2 , and treatment of the resulting compound with water. Ketones containing C n H 2 n+ 1 radicles, when heated with HCN, HC1, and water, produce acids of this group :—
CO (CH 3 ) 2 + HCN + 2OH 9 + HC1=C(CH 3 ) 2 . HO.CO 2 H + NH 4 C1. Acetone. Oxybutyric acid.
3. Primary Olefine Acids, C n H an j βθ s θ∩^.—Produced by the ÍCH OH
CH 2 OH ,
4. Secondary Olefine Adds, C n H sn ∣ CO OH, " +1 ^ θ^∙—^ r0 ^ uce ^ by the action of nascent hydrogen on certain compounds formed by the action of sodium on ethyl acetate, CH 3 .CO(C 2 H s O). (See p. 573.)
The following is a list of the acids of the lactic series :— [table]
Ethylidene lactic acid is the first product of the butyric fermentation of sugar; ethylene and paralactic acids are contained in the juice of flesh. Lactide and dilactix acid (C 3 H 4 O 2 and C 6 H 10 O 5 ) are produced by the elimination of water from one and two molecules of lactic acid by heat.
Senes C n H 2n - 1 O.CO 2 H, Pyruvic Series. — These acids may be regarded as derived from those of the lactic series by the abstraction of H 9 . The following are known :— [table]
Series C n H 2n+1 O 2 .CO 2 H comprises glyoxylic and glyceric acids, CH 3 O 2 .CO 2 H and C 2 H 5 O 2 .CO 2 H.
‰+C n H 2 n- 1 ∙CO 2 H, Acrylic Series. — The acids known in this series are the following :— [table]
These acids are produced by the oxidation of the corresponding aldehydes, and also by the dehydration of secondary lactic acids. They may be arranged in two isomeric series:— primary acids of the form CHR".CO 2 H,and secondary acids, CR'R". CO 2 H. These acids combine directly with ZZα 2 , ~RHa, and are raised to the C n H 2n+1 .CO 2 H type. On fusion with KHO they yield potassium salts of two fatty acids.
Series C n H 2n . 3 .CO 2 H contains tetroleic (C 3 H 3 .CO 2 H), sorbic (C 5 H 7 .CO 2 H), and stearolic (C 17 H 31 .CO 2 H) acids. Series C n H 2n - 7 .CO 2 H, Benzoic or Aromatic Series. — The following are known :— [table]
These acids are prepared by the general methods of oxidizing the corresponding alcohols and aldehydes, and by the action of KHO cn the cyanides of the C n H 2n-7 radicles. Benzoic acid exists in gum-benzoin, from which it sublimes, on heating, in white feathery crystals, having a fragrant odour. By the action of Cl, HNO 3 , &c., these acids yield substitution derivatives, such as C 6 H 4 C1.CO 2 H and C 6 H 4 (NO 2 ).CO 2 H, chlorobenzoic and nitrobenzoic acids. The latter, by the action of nascent hydrogen, is converted into C 6 H 4 (NH 2 ).CO 2 H, amidobenzoic acid. Hippuric acid, the potassium salt of which occurs in the urine of many herbivorous animals, is benzamido-acetic acid — CH 2 (NH.COC 6 H 5 ).CO 2 H.
The aromatic acids can be arranged in two metameric series corresponding to the benzylic alcohols and the phenols.
Series C n H 2n-s (HO).CO 2 H, Oxybenzoic Series. — This bears to the preceding series the same relationship as the lactic to the fatty series :— [table]
Series C,,H 2n-9 (HO) 2 .CO 2 H, Dioxybenzoic Series, contains oxysalicylic, hypogallic, protocol echwic, and carbohydro-quinonic acids — C 6 H 3 (HO) 2 .CO 2 H.
Allied to this series are piperic (C 12 H 10 O 4 ) and eugetic (C 77 H 72 O 4 ) acids. Series C n H 2n-10 (HO) 3 .CO 2 H, Gallic Series. — Gallic acid, C 6 H 2 (HO) 3 .CO 2 H, exists ready formed in many plants, such as sumach, hellebore, &c. Tannins are the astringent vegetable principles made use of in tanning. They all contain some form of tannic acid, C 27 H 22 O 17 , a compound resolved by the action of acids into gallic acid and glucose (see p. 572). Gallic acid is generally prepared from the gallotannic acid extracted from gall-nuts.
Series C n H 2n - 9 .CO 2 H :— Cinnamic, atropic, and isatropic acids, C 8 H 7 .CO 2 H.
Series C n H 9n . 10 (HO).CO 2 H :— Coumaric and paracourna∕ ric acids, C 8 H β (HO).CO 2 H.
Series C n H 2n-11 .CO 2 H :— Phenyl-propiolic acid,
C 8 H 5 .CO 2 H.
Series C n H 2n - 13 ∙CO 2 H, comprising two naphthoic acids, C 10 H 7 .CO 2 H, Series C n H 2n - 14 (HO).CO 2 H, containing oxynaphthοic adds, C 10 H 6 (HO).CO 2 H.
Series C n H 2n - 16 (HO).CO 2 H, containing benzilic acid, C 13 H 10 (HO).CO 2 H.
Series C n H 2n . 19 ,CO 2 H :— Anthracenecarbonic acid, C 14 H 9 ,CO 2 H.
Dibasic Adds. Series C n H 2n (CO 2 H) 2 ,Sucdnic Series. — The following are known :— [table]
These acids may be prepared by the general methods of oxidizing the corresponding alcohols (glycols), and from the cyano-olefine compounds, C n H 2n (CN) 2 , by the usual reactions. Oxalic acid (sodium salt) may be synthesized by heating sodium in dry CO 2 , and is manufactured on the large scale by fusing sawdust with caustic alkalies. This acid, free or combined, is found in many plants,—hence the name (from Oxalis, wood-sorrel). Succinic acid exists in amber. Many of these acids are produced by oxidizing various organic substances with nitric acid. The acids of the present series are crystalline solids, forming, like all
This acid (the hydroxyl derivative of formic acid) is unknown in the free state, and is dibasic.
dibasic acids, two series of salts, normal and acid, of the . 1 ∩ ∏ í COOH an Λ∏ π J COOM' τ formulae C n H 2n j COOM z anc ^ k>>⅛ ∣ OOOM ’ ’ Isθmθr^ ides of many of the acids are known. Oxalates may be formed from cyanogen compounds (see p. 554).
Series C n H 9n . 1 (HO)(CO 2 H) 9 , Malic Series:— [table]
Malic acid is found free or combined in the juice of most fruits.
Series C, 1 H 2n - 2 (HO) 2 (CO 2 H) 2 , Tartaric Series. — The following are known :— [table]
Tartaric, like malic and oxalic acids, is of frequent occurrence in the vegetable kingdom, being found free or combined in the juice of many fruits. It is generally obtained from argol or tartar (crude acid potassium tartrate), which is deposited from fermenting grape juice. It can be formed by the action of Ag 2 O and water on dibromsuccinic acid, C 2 H 2 Br 2 (CO 2 H) 2 . Five modifications of tartaric acid, differing chiefly in their optical properties, are known, viz., dextrotartaric, levotartaric, racemic, mesotartaric, and metatartaric acids. Normal and acid tartrates are known. Tartar emetic is potassio-antimonious tartrate; the acid potassium tartrate is known as cream of tartar. Series C n H 2n - 2 (CO 2 H) 2 , Fumaric Series. — This consists of fumaric and maleic acids, C 2 H 2 (CO 2 H) 2 ; citraconic, itaconic, and mesaconic acids, C 3 H 4 (CO 2 H) 2 .
Series C n H 2n √CO 2 H) 2 , Phthalic Series, comprises :— [table]
The phthalic acids are obtained by the oxidation of many aromatic hydrocarbons and their derivatives. Derived from these acids are hydro-phthalic and hydrotere-phihalic acids (C 8 H 8 O 4 ).
Tribasic Acids. Meconic acid, C 4 HO(CO 2 H) 3 , is obtained from opium; trixarballylic add, C 3 H 5 (CO 2 H) 3 , from tricyanopropane (see p. 568).
Citric acid, C 3 H 4 (HO) (CO 2 H) 3 , exists in many fruits, and is generally obtained from lemon juice. It forms colourless crystals readily soluble in water. Being a tribasic acid, it forms with metals three classes of salts, typified by M'C 6 H 7 O 7 , M 2 'C 6 H 6 O 7 , M 3 'C 6 H 5 O 7 , <fcc.
Aconitic acid, C 3 H 3 (CO 2 H) 3 , is derived from citric acid, and is found also in monkshood (Aconitum NapeUus). Trimellitic (para-), trimesic (meta-), and hemimellitic (ortho-) acids, C 6 H 3 (CO 2 H) 3 , are tricarboxyl derivatives of benzene.
Tetrabasic and Hexabasic Acids.
The following tetracarboxyl benzene derivatives are known :— pyromelUtic (para-), prehnitic (meta-), and mello-ρhanic (ortho-) acids, C 6 H 2 (CO 2 H) 4 .
The hexabasic acids known are mellitic acid, C 6 (CO 2 H) 6 , and its derivative hydromellitic acid, C 6 H 6 (CO 2 H) 6 .
It is to be remembered that the foregoing list is necessarily brief. Many acids of vegetable and animal origin, and others derived from them by artificial methods, are known, but their constitution is in many cases still undecided.
IX. Anhydrides.
These compounds are the ethers of acid radicles (see p. 553), and may be prepared in many cases by analogous reactions :— [table]
Many anhydrides are obtained by heating the corresponding acids :— [table]
Anhydrides when acted on by water yield the corresponding acids :— [table]
Compound anhydrides when distilled are resolved into two simple anhydrides—
2C 2 H 3 O.O.C 7 H 5 O = (C 2 H 3 O) 2 O + (C 7 H 5 O) 2 O.
Aceto-benzoic anhydride. Acetic anhydride. Benzoic anhydride.
Acetic anhydride is an oily liquid, boiling at 137°, and having a pungent odour. By the action of P 2 S 5 it yields thiacetic anhydride, (C 2 H 3 O) 2 S, and by BaO 2 acetic peroxide, (C 2 H 3 O) 2 O 2 . It combines directly with aldehydes.
X. Acid Halides.
These may be regarded as the haloid ethers of acid radicles. They can be prepared by the action of haloid phosphorus compounds on the acids containing the corresponding radicles— [table]
A similar reaction can be effected with the haloid substitution compounds of the acids. Thus, C 2 H 2 C1O.C1, C 2 HC1 2 O.C1, and C 2 Cl 3 O.Cl (mono-, di-, and tri-chlor-acetyl chlorides) can be obtained by the action of PCl 3 on mono-, di-, and tri-chloracetic acids.
The compounds of this family must be distinguished from the haloid salts of the acids which contain the halogen atom in the place of the carboxyl hydrogen; for instance—
0 0
II II
h 3 cc—ci h 3 c- C—OC1
Acetyl chloride. chlorine acetate.
The latter compound is a yellow liquid produced by the action of acetic anhydride upon hypochlorous anhydride: (C 2 H 3 O) 2 O + C1 2 O = 2C 2 H 3 O. OC1. It is very unstable, being decomposed by heat, Br, I, and most metals. The corresponding iodine compound produced by the action of iodine on the chlorine acetate is interesting as showing the triad nature of the iodine atom—[(C 2 H 3 O)O] 3 Γ".
The acid halides are as a rule colourless liquids more or less oily, and possessing characteristic odours. They are decomposed by water :— [table]
In some cases the whole of the halogen is not removed by the action of cold water :—
(C 3 H 4 O)"C1 2 + OH 2 = (C 3 H 4 C1O)HO + HC1.
Lactyl dichloride. Chloropropionic acid.
Acetyl chloride, when made to act upon potassium hydrosulphide, forms thiacetix arid :—
C 2 H 3 O.C1 + KHS = (C 2 H 3 O)HS + KC1.
When it is mixed with acetic acid, and acted on by sodium amalgam, the corresponding alcohol is produced :— [table]
Some acid halides can be synthesized from hydrocarbons by the action of carbonyl dichloride (phosgene gas)— [table]
XI. Ethereal Salts.
These compounds are derived from acids by the substitution of hydrocarbon radicles for hydrogen, but it is only the hydrogen entering into the composition of hydroxyl that can be thus replaced. Thus, the number of ethereal salts that any acid can form depends upon the number of times it contains HO; in other words, an íi-hydric acid can form n ethereal salts. It has been previously mentioned that the basicity of an organic acid depends upon the number of times it contains COOH, hence it follows that monohydric acids must be monobasic, but n-hydric acids may be 1,2, 3, η-basic. For example, tartaric acid is tetrahydric, but as only two of its hydrogen atoms are replaceable by metals it is dibasic, and consequently contains 2(COOH). But this acid being tetrahydric contains 4(HO), and has thus four hydrogen atoms replaceable by radicles, giving rise to four ethereal salts :— [table]
Ethereal salts are normal or acid according as the carboxyl (basic) hydrogen is entirely or partially replaced by radicles.
Every acid, inorganic and organic, can give rise to the formation of ethereal salts, these being in many cases produced by the direct action of the acid on the alcohol containing the necessary radicle :— [table]
Other methods of preparing ethereal salts are exemplified by the following reactions:— [table]
[In practice this reaction is effected by passing HC1 gas into a mixture of the acid and alcohol.l [table]
Most ethereal salts are decomposed on heating with water into an acid and an alcohol. The same reaction is more speedily induced by caustic alkalies:— [table]
The ethereal salts of nitrous acid are metameric with nitro-derivatives of hydrocarbons; thus— [table]
Ethyl nitrite is produced by the action of ethyl sulphuric acid on potassium nitrite. Nitroethane is formed when ethyl iodide acts on silver nitrite. By the action of nascent hydrogen the difference in constitution between these two compounds is well brought out :— [table]
Many ethereal salts exist ready formed in various animal and vegetable substances. Thus, oil of Wintergreen (Gaultksría procumbens) contains methyl salicylate, C 6 H 4 (HO).CO(OCH 3 ). The ethereal salts of glycerin (glycerides) have received special names:— [table]
Stearins.
C 3 H s (HO) 2 (C 18 H 35 O 2 ), Monostearin. C 3 H s (HO)(C 18 H 35 O 2 ) 2 , Distearin. C3H 8 (C j8 H3 0 O 2 )3, Tristearin.
(Stearins occur in suet, tallow, and other fats.)
Glucosides are compounds existing in vegetables which by the action of reagents or natural ferments are resolved into glucose and some other compound. Thus amygdalin, a crystalline substance from bitter almonds, <fcc., when decomposed by the action of synaptase or emnlsin, a ferment existing in the plant, is converted into glucose, benzoic aldehyde, and hydrocyanic acid: C 20 H 27 NO 11 + 2OH 2 = C 7 H β O + HCN + 2C 6 H 12 O 6 . Among the more important glucosides are :—
Gallotannis acid (C 27 H 22 O 17 ), from oak-galls, sumach, &c.
Salicin (C 13 Hι 8 O 7 ), from bark and leaves of poplar and willow.
Aesculin (C 21 H 24 O 13 ), from bark of horse-chestnut, &c.
Glycyrrhizin (C 24 H s6 O 9 ), from liquorice root.
Quercitrin (C 33 H 30 O 17 ), from bark of quercitron (Quercus infectoria). Phlorizin (C 21 H 24 O 10 .2OH s ), from root bark of apple, pear, plum, and cherry.
Many of the vegetable colouring matters employed in dyeing are glucosides. Thus indican, from woad (Isatis tinctoria), is a colourless substance decomposed by acids into indigo-blue and a glucose-like body (indiglucin) :—
2C 26 H 31 NO 17 + 4OH 2 = Cj 6 H 10 N 2 O 2 + 6C 6 H 10 O 6 .
indican. Indigo blue. Indiglucin.
The most valuable tinctorial constituent of madder-root (Rubia tinctoτda) is alizarin (C 14 H 8 O 4 ), which exists in the root as a glucoside (rubianic acid) (C 26 H 28 O 14 ). When heated with zinc dust alizarin yields anthracene (C 14 H 10 ), and conversely, anthracene can be converted into alizarin by first oxidizing it to anthraquinone (C 14 H 8 O 2 "), then by treatment with Br or H 2 SO 4 transforming the quinone into dibromanthraquinone or anthraquinone-disulphonic acid, and finally fusing either of these products with caustic alkalies: C 14 H 6 Br 2 O 2 " + 2KHO ≈ C 14 H 8 O 4 + 2KBr. Large quantities of artificial alizarin are thus manufactured.
Vanillin (C 8 H 8 O 3 ), the odorous principle of vanilla, has also been obtained artificially from coniferin (C 16 H 22 O 8 ), a glucoside obtained from the cambium of Coniferous trees. This glucoside when heated with water and emulsin is resolved into glucose and a crystalline substance (C 10 H 12 O 3 ) which on oxidation is converted into vanillin.
By the action of sodium on ethyl acetate there is produced among other compounds the sodium salt of an acid, C 6 H 10 O 3 , which is of extreme interest on account of its enabling the ascent of the series of ketones and fatty acids to be effected. Thus :— [table]
The ascent of the fatty acids is more simply explained by considering the first products of the action of Na on ethyl acetate to result from the replacement of the methyl hydrogen by Na. Thus:— [table]
The liberation of the acid from the ethereal salt is effected by the action of KHO, as previously explained.
XII. OκGAN0-METALLIC BODIES.
The constitution of these compounds has been already considered (p. 553); they have to be distinguished from organic compounds containing metals, such as metallic salts of organic acids. In organo-metallic bodies the metal is directly combined with the hydrocarbon radicle, while in organic compounds containing metals the metallic atom is connected with the radicle through the medium of some other element:—
H 2 H 2 0 0
II II II II
II 3 CC—Zn—CH 3 H 3 C-C—O—Zn—O—C—CH 3
Zinc ethide. Zinc acetate.
Organo-zinc compounds are prepared by the action of zinc upon the iodides of C n H 2n+1 radicles :—
2Zn + 2C n H 2n+1 I = Zn(C n H 2n+1 ) 2 ÷ Znl 2 .
The operation is most rapidly conducted by digesting the iodide with zinc foil coated with copper by previous immersion in a weak solution of cupric sulphate (Gladstone and Tribe). Zinc ethide, the first of these compounds obtained by Frankland, is a limpid, mobile liquid boiling at 118° C.,and spontaneously inflammable.
Organo-compounds of tin, lead, mercury, bismuth, arsenic, and antimony have been prepared by the action of the iodides and bromides of the corresponding radicles on the potassium or sodium alloys of the metals :—
HgNa 2 + C 2 H 5 I = Hg(C 2 H 5 ) 2 + Nal
Sodium amalgam. Ethyl iodide. Mercuric ethide.
HgNa 2 + 2C 6 H 5 Br = Hg(C 6 H 5 ). 2 + 2NaBr. Sodium amalgam. Phenyl bromide. Mercuric phenide.
Many organo-metallic bodies are derived from others by simple replacement or by double decomposition :—
SnCl 4 + 2Zn(C 2 H 5 ) 2 = Sn(C 2 H 5 ) 4 + 2ZnCl 2 stannic chloride. Zinc ethide. Stannic ethide.
3Hg(C 2 H 5 ) 2 + Al 2 = A1 2 (C 2 H 5 ) 6 + 3Hg Mercuric ethide. Aluminic ethide.
3Zn(C 2 H 5 ) 2 + Na 2 = 2[Na(C 2 H 5 ), Zn(C 2 H 5 ) 2 ] + Zn. Zinc ethide. Sodic zinc ethide.
In the case of organo-compounds of monad metals prepared by the direct action of the metal on some other body, as in the last reaction, the original compound occurs inseparably in the resulting compound.
The following is a list of the organo-metallic bodies :— [table]
Not many isomerides of the above compounds have as yet been obtained; zinc pseudopropide and mercuric benzylide (isomeric with the tolylide) are known.
The extreme readiness with which the organo-metallic bodies exchange their hydrocarbon radicles for more negative elements or radicles obviously renders them of particular value in organic synthesis.
The following are further important reactions :—
1. Synthesis of fatty acids by the direct absorption of CO 2 by organo-sodium compounds :—
CO 2 + Na(C n H 2n+1 ) = I cøONa ,
2. Displacement of halogens, oxygen, (HO), (OC n H 2n+1 ), <fcc., by C n H 2n+1 radicles (see the preparation of Sn(C 2 H 5 ) 4 from SnCl 4 given above) :—
í CO(OC 2 H 5 ) 7 í c 2 h 5 20h _ J C(C 2 H 5 ) 2 .OH
I CO(OC 2 ⅛ +zπ 1 C 2 H 5 + jυu 2- I CO(OC 2 H 5 ) +
Ethyl oxalate. Zinc ethide. Ethyl diethoxalate.
zn {θH + c Λ∙≡ 0 ∙ Zinc hydroxide. Alcohol.
Also, formation of paraffins by action of water on organo-zinc bodies :—
Zn(C n H 2n+1 ) i + 2OH 2 = Zn(HO) 2 + 2C n H 2n+2 .
Compounds intermediate between organo-metallic bodies and haloid salts of the metals have been obtained, e.g. —
Hg(CH 3 )I , Sn(CH 3 ) 2 I 2 , Sn 2 (C 2 H 5 ) 4 T,,
Mercuric methiodide. Stannic dimetho-diiodide. Distannic tetretho-diiodide.
Sn(C 6 H 5 )(C 2 H 5 )Cl 2 .
Stannic etho-phenyl-dichloride.
The action of slow oxidation upon the compounds of this family is shown in the following examples :—
Zn(C 2 H 5 ) 2 + 0 = Zn(C 2 H 5 )(OC 2 H 5 )
Zinc ethide. Zinc etho-ethylate.
Zn(C 2 H 5 ) 2 + 0 2 = Zn(0C 2 H 5 ) 2 .
Zinc ethide. Zinc ethylate.
Organo-boron compounds containing boron directly combined with hydrocarbon radicles have been obtained by the aid of zinc methide and ethide :—
2B(OC 2 H 5 ) 3 + 3 Zn(CH 3 ) 2 = 2B(CH 3 ) 3 Triethyl borate. Zinc methide. Boric met hide.
+ 3Zn(OC 2 H 5 ) 2 .
Zinc ethylate.
Boric ethide has been prepared by an analogous reaction. Boric methide is a gas, and boric ethide a limpid liquid; both are spontaneously inflammable, burning in air with a green-tinted flame. They combine with ammonia forming compounds of the formula NH 3 , BR' 3 .
Organo- silicon compounds are bodies in which carbon is replaced partially or entirely by silicon. Many of these are formed by means of zinc methide and its homologues :—
SiCl 4 + 2Zn(C 2 H 5 ) 2 = Si(C 2 H 5 ) 4 + 2ZnCl 2 .
Silicic chloride. Zinc ethide. Silicic ethide.
A large number of these compounds are now known. Their analogy to the carbon compounds will be seen from the following examples :—
Si(CH 3 ) 4 C(CH 3 ) 4 SiHCl 3 CHC1 3 .
Silicic methide. Tetramethyl-methane. Silicon chloroform. Chloroform.
H.SiO(OH) CH 3 .SiO(OH) C 2 H 5 .SiO(OH)
Silicoformic acid Siliconacetic acid. Siliconpropionic add.
H.CO(OH) CH 3 ,CO(OH) C 2 H 5 .CO(OH).
Formic add. Acetic acid. Propionic acid.
XIII. Amines.
It has been previously stated that amines are derivatives ©f ammonia or its hydrate and haloid salts (p. 553), hydrogen being partially or entirely replaced by hydrocarbon radicles. Amines built on the type of NH 3 may arise from the replacement of hydrogen in NH 3 , N 2 H 6 , or N 3 H 9 , thus forming monamines, diamines, or triamines :— [table]
Similarly triamines are formed by replacement of H 3 by R'".
Each of these groups of amines is further divisible into primary, secondary, and tertiary, according as one-third, two-thirds, or all the hydrogen of ammonia is replaced by hydrocarbon radicles. For example :— [table]
Monamines containing C n H, 2n+1 radicles. — The following are known :— [table] [table]
These amines are produced by the following methods :— 1. By heating the haloid compounds of C n H 2 , l+1 radicles with a solution of ammonia in alcohol :—
C n H 2n+ I 1 + NH 3 = N(C⅛ +1 )H 3 I.
Iodide of radicle. R'-yl-ammonium Iodide.
The resulting compound is formed on the type of the salts of ammonia, and like these bodies is decomposed by fixed alkalies :—
N(C n H 2n+1 )H 3 I + KHO = N(C n H 2n+1 )H 2 + OH 2 + KI.
R⅛>- 1 ∏≡ ’ a ,, 11m ,
At the same time, according to the proportion of the reagents, the temperature, <fcc., in the first reaction, more or less of the secondary and tertiary monamines are produced :—
2C n H 2n+1 I + 2NH 3 = N(C⅛ +1 ) 2 H 2 I + NH 4 I
Secondary monamine hydriodide.
3C⅛ +1 I + 3NH 3 = N(C n H 2n+1 ) 3 HI + 2NH 4 I.
Tertiary monaminβ hydriodide.
Amines containing different radicles are obtained thus :—
N(C 2 H 5 )H 2 + CH 3 I = N(CH 3 )(C 2 H 5 )H 2 I Ethylamine. Methyl iodide. Methyl-ethylammonium
The secondary and tertiary amines are in all these cases liberated by the action of KHO.
2. Cyanic and cyanuric acids (p. 554) form two classes of ethereal salts analogous to the cyanides and isocyanides (p. 555); thus—
N≡C-OR' O=C=N-R'
ũyanate. Isocyanate.
Isocyanates are prepared by distilling ethereo-potassium salts of H 2 SO 4 containing the necessary radicles with potassiumcyanate: R'KSO 4 + CO".NK = CO".NR' + K 2 SO 4 . Isocyanurates are obtained by a similar reaction, substituting potassium cyanurate for cyanate.
Both these classes of ethereal salts furnish primary amines on distillation with caustic alkali :—
CO .NR ’ + 2KHO = NRΗ 2 + K 2 CO 3 .
3. Nitro-substitution derivatives of the paraffins are reduced by nascent hydrogen, ammonium sulphide, &c., with the formation of amines (see reduction of nitroethane to ethylamine, p. 572).
With the exception of the methylamines (which are gaseous), the amines of the present class are limpid liquids, having powerful ammoniacal odours and highly basic properties, restoring the colour of red litmus, and uniting with acids to form salts analogous to those of ammonium :— [table]
Like ammonium salts, these amines form yellow crystalline double salts when their hydrochlorides are mixed with a solution of platinic chloride. The general formula of these salts (A representing amine) is 2AHC1, PtCl 4 . Tertiary amines unite directly with iodides of C n H 2 , +1 radicles, forming compounds of the type of haloid ammonium salts :—
n(c 2 h 5 ) s + c⅛ι = n(c 2 h 5 ) 4 i ae. τ ≡sr
These compounds do not again furnish the tertiary amine on treatment with KHO, but on heating with Ag 2 O and water bodies such as tetrethylammonium hydroxide, N(C 2 H 5 ) 4 HO, are produced. These bodies are highly caustic bases analogous to ammonium hydroxide, but more like the fixed alkalies in their general behaviour.
The amines of the present group are susceptible of metamerism and isomerism. For example :— Metameric... N^(C 3 H 7 )H 2 , N(CH 3 )(C 2 H 5 )H , N(CH 3 ) 3 .
Propylamine. Metħyl-ethylamine. Trimethylamine.
Isemeric...N(C 3 H 7 )H 2 , N[CH(CH 3 ) 2 ]H 2 . Propylamine. Pβeudopropylamine.
By the action of nitrous acid primary amines are con verted into nitrites, which are decomposed on heating, with the formation of the corresponding alcohol :—
N(C n H 2n+1 )H 2 ,HNO 2 = C n H 2n+1 .HO + OH 2 + N 2 .
The alcohols produced by this reaction are isomeric and not identical with those of the radicle from which the amine is derived in cases where such isomerism is possible.
The action of nascent hydrogen on the nitriles gives rise to the formation of amines :—
C n H 2π+1 .CN + 2H 2 = N(CH 2 C n H 2n+1 )H 2 .
The nitriles can be obtained from the alcohols by combining the latter with sulphuric acid, converting the sulpho-acid thus produced into a potassium salt, and distilling this with potassium cyanide (p. 566). By combining all these reactions the homologous series of alcohols can be ascended. To give an illustration :— Methyl alcohol is combined with sulphuric acid, and the resulting sulphoacid converted into potassium methylsulphate. This latter salt is mixed with potassium cyanide and submitted to dry distillation, when methyl cyanide (acetonitrile) is produced; this when acted on by nascent hydrogen is converted into ethylamine, which by the action of nitrous acid is transformed into ethyl alcohol. By a similar series of reactions ethyl alcohol can be converted into propylamine, which by the action of HNO 2 is converted into pseudopropyl alcohol. Isocyanides or Carbamines. — These compounds, metameric with the cyanides of hydrocarbon radicles (p. 555), may be regarded as formed on the type of ammonium salts; thus (X' representing the acid radicle)—
NΗ 4 X' N t C it R'
Ammonium salts. carbamine.
Most of the reactions made use of in preparing the cyanides give rise to the simultaneous formation of isocyanides (see p. 555); the latter bodies are in excess when the iodides of the radicles are made to act upon silver cyanide. Isocyanides are also obtained by heating primary amines with chloroform and caustic potash :—
NR'H 2 + CHC] 3 + 3KHO = N ∣ + 3KC1 + 3OH 2 .
This reaction affords a ready means of distinguishing primary amines, as the isocyanides are at once detected, even in the smallest traces, by their powerful and nauseous odours. The carbamines reduce many metallic oxides (such as HgO, Ag 2 O, &c.) with the formation of cyanates of the contained hydrocarbon radicles. They possess properties decidedly more basic than their metamerides, combining energetically with acids to form salts. When heated in sealed tubes the isocyanides are partially transformed into the corresponding nitriles.
Monami∙nes containing C w H 2λ . 1 radicles are represented by allylamine, N(C 3 H 3 )H 2 .
Monamines containing C B H 2n _ 7 radicles. — These bases belong to two isomeric series corresponding to the alcohols of the benzyl series and to the phenols. This relationship is exemplified by the following formulae :—
C 6 H 4 (NH 2 ).CH 3 C 6 H 5 .CH 2 (NH 2 )
Toluidine. isomeric with Benzylamine
C 6 H 3 (NH 2 ),(CH 3 ) 2 C 0 H 4 .CH,.CH λ XH 2 )
Xylidine isomeric with Xylylamine.
The following are the more important amines of the present group :— [table]
The amines of the benzyl series are obtained by the action of the chlorides of the corresponding radicles on ammonia. Toluidine and its homologues are prepared by reducing the nitro-derivatives of the corresponding hydrocarbons by means of acetic acid and iron (ferrous acetate), ammonium sulphide, &c. :—
C 6 H 5 (NO 2 ) + 3SH 2 = C 6 H 5 (NH 2 ) + 2OH 2 + 3S.
Nitrobenzene. Amidobenzene (aniline).
These latter amines, which may be regarded as amidobenzene, in which 1, 2, 3, <kc., H atoms are replaced by R', R 2 ', R 3 ', &c., are susceptible of the isomeric modification dependent on the relative positions of the replaced H atoms. Thus we have—
C 6 (NH 2 )(CH 3 )HHHH C 6 (NH 2 )H(CH 3 )HHH
1 2 3456 12 3 456
Orthotoluidine. Metatoluidine.
C 6 (NH 2 )HH(CH 3 )HH
1 2 3 4 5 6
Paratoluidine.
All these amines are basic colourless liquids, more or less oily, and possessing highly characteristic odours. They combine with acids forming crystalline salts, which are decomposed by caustic alkalies with the liberation of the amine.
Phxnylamine or aniline, the first member of the present series of primary amines, may be regarded as the first homologue of both the above isomeric series. It derives its name from the indigo-plant (Indigofera Anil), as it was first obtained by distilling indigo with caustic potash. Aniline is found in small quantities in coal-tar oils, but is manufactured on the large scale by reducing nitrobenzene with iron and acetic acid. Aniline is, when pure, a colourless, oily liquid, having a peculiar odour; it boils at 182° C. j and solidifies at - 8° C. The replacement of the phenylic hydrogen by halogens gradually destroys the basic character of aniline. Thus, chloraniline, C 6 H 4 C1(NH 2 ), and dichloraniline, C 6 H 3 Cl 2 (NH 2 ), form crystalline salts with acids: but trichloraniline, C 6 H 2 C1 3 (NH 2 ), possesses no basic properties. By the action of the iodides of C n H 2n+1 radicles upon aniline, the amido-hydrogen can be replaced by these radicles, giving rise to the formation of secondary and tertiary mcnamines containing different radicles, such, for example, as eth,yl-aniline, (C 6 H 5 )(C 2 H 5 )NH, produced by the action of C 2 H 5 I upon aniline, and subsequent treatment with KHO; ethyl-amyl-anilvne, (C 6 H 6 )(C 2 H 5 )(C 5 H 11 )N, formed by the action of C 5 H n I upon ethyl-aniline, &c. These tertiary amines combine directly with C n II 2 ,, +1 I, producing stable iodides—
N(C 6 H 5 )(C⅛ +1 )(C m H 2w+1 )(C p H 2p+1 )I, which are not decomposed by KHO, but by the action of Ag 2 O and water are converted into alkaline hydroxides, analogous to NH 4 HO. It will be seen that the iodides of C m H 2n+1 radicles can thus be employed to determine whether the amine is primary, secondary, or tertiary. When aniline hydrochloride and methyl alcohol are heated together in a sealed tube, the amido-hydrogen is partially or entirely replaced by methyl, giving rise to the formation of methyl and dimethyl-aniline hydrochlorides. When both atoms of amido-hydrogen have been thus replaced, if the heating is continued the phenylic hydrogen becomes replaced by methyl giving rise to dimethyl-toluidine and its homologues. This interesting intra-molecular substitution was discovered by Hofmann. The nature of the transformation will be seen from the following formulae :—
N(C β H 5 )(CH 3 )(CH 3 ) N[C 6 H 1 (CH,)](CH 3 )(CH 3 )
Dimethylaniline. Dimethyltoluidine.
N[c 6 H 3 (CH 3 )J(CH s XCH 31 Dimethylxylidine.
(The last atom of phenylic hydrogen does not appear to be removed by this reaction.)
Nitrous acid acts upon aniline in a manner similar to its action upon the C n H 2n+1 primary monamines, converting it into the corresponding alcohol (phenol) :—
C 6 H 5 (NH 2 ) + HNO 2 = C 6 H 5 (HO) + OH 2 + n 2 .
If the aniline is diluted with alcohol, however, intermediate products are formed—
2C 6 H 5 (NH 2 ) + HNO 2 = C 12 H 11 N 3 + 2OH 2 Aniline. Azodiphenyldiamine.
C 12 H π N 3 + HNO 2 = C 12 H 8 N 4 + 2OH 2 . Azodiphenyldiamine. Diazo-diphenyldiamine.
If salts of aniline are employed, salts of these azo-derivatives are produced.
Aniline combines with aldehydes with the elimination of water, producing pkenaldines :—
2C 6 H 5 NH 2 + CH 3 ,COH - OH 2 = C 14 H 16 N 2 .
Paraniline a polymeride of aniline, having the composition C 12 H 14 N 2 , is known.
Allied to the present group of amines is naphthalidine, C 10 H 7 (NH 2 ), produced by the reduction of nitronaphthalene.
Natural Alkaloids. — These compounds are organic bases existing already formed in plants. With the exception of pψm'dme(C 5 H 11 N) and canine, which are secondary mona-mifles, the alkaloids at present known are tertiary compounds. The following are some of the more important of this group of amines :— [table] [table]
The constitution of the radicles contained in these various bases has not hitherto been made clear. Among the cinchona alkaloids—of which the valuable medicine quinine is a member—some interesting cases of physical isomerism occur. Many bases have been obtained by the action of reagents on the opium alkaloids. A base isomeric with coni ne has been prepared artificially by heating normal butyl aldehyde with an alcoholic solution of ammonia so as to produce dibutyraldine : 2C 3 H 7 .COH + NH 3 -OH 2 = C 8 H 17 NO, and then submitting this latter substance to dry distillation: C 8 H 17 NO - OH 9 = C ft H 1fi N.
Bases from, various Sources. — The following are a few bases of animal origin :— [table]
Many bases have been obtained by the destructive distillation of organic matter containing nitrogen. These compounds are tertiary monamines:— [table]
Others are derived from oil produced by the destructive distillation of bones, <fcc., and coal-tar naphtha :— [table]
Hydramides (p. 568), when boiled with potash solution, are converted into basic compounds. Amarine, C 21 H 18 N 2 , is a substance formed in this manner from hydrobenzamide.
Diamines are formed by reactions analogous to those employed in the preparation of monamines :—
C⅛Br 2 + 2NH 3 = N 2 (C n H 2 „)H 4 .2HBr R" dibromide. R"-diamine dihydrobromide.
As in the preparation of monamines, secondary diamines are at the same time formed :—
2C n H 2n Br 2 + 4NH 3 = N 2 (C n H 2n ) 2 H 2 . 2HBr + 2NH 4 Br.
R" dibromide. Di-R"-diamine dihydrobromide.
The action of KHO upon the primary diamine dihydrobromide is different from its action upon the corresponding monamine compound, inasmuch as an oxide of the diamine is formed :—
N 2 (C n H 2n )H 6 Br 2 + 2KHO - N 2 (C n H 2n )H 6 O + 2KBr + OH 2
R"-diamine dihydro-R-"diammonium oxide bromide.
The diamines corresponding to benzene and its homologues are produced by the reduction of the dinitroderivatives of the respective hydrocarbons—
C 6 H 4 (NO 2 ) 2 + 6H 2 = C 6 H 4 (NH 2 ) 2 + 4OH 2 .
Dinitrobenzene. Phenylene diamine.
Intermediate products are sometimes formed—
C β H 4 (NO 2 ) 2 + 3H o = C 6 H 4 (NO 2 )(NH 2 ) + 2OH 2 .
Dinitrobenzene. Nitranifine.
The diamines have not been investigated with the amount of detail that has been bestowed upon the monamines. By the action of the iodides of C n H 2n+1 radicles upon ethene and diethene diamines a large number of bases have been obtained, in which the ammonic hydrogen is more or less replaced by these radicles. Diamines form two classes of salts, monacid and diacid.
Triamdnes. — The following belong to this class :— [table]
Aniline Colours. — A large number of the colours obtained from aniline are triamines, and may be conveniently referred to here, ]iosαηiline (magenta or fuchsine) is the base of a splendid red dye which serves as the starting point of numerous other colours. It is prepared by oxidizing a mixture of aniline and toluidine with arsenic acid :—C 6 H 7 N + 2C 7 H 9 N - 3H 2 = C 20 H 19 N 3 (Rosaniline). Mauve is a purple dye produced by the oxidation of aniline by means of sulphuric acid and potassium dichromate. The base of this colour is mauυeiηe (C 26 H 24 N 4 ), and it is interesting as being the first of the aniline dyes practically employed in the arts.
The numerous other dyes derived from benzene, phenol, and naphthalene cannot be discussed here.
It has been recently observed by Dr Otto Witt that the peculiar properties of organic colouring substances are in most cases due to the combined presence of two groups or radicles attached to a carbon nucleus. The one which is the colour-producing group is termed the chromophore, the other is a salt-forming group, ι.e , a group which confers upon a molecule either acidity or basicity. A substance containing a chromophore does not become a colour until the salt-forming group is introduced into the molecule; such potential colouring matters are therefore termed chromogens. The chromophoric group exercises its influence more powerfully in the salts of colouring substances. The principal salt-forming groups combined with aromatic nuclei are HO and NH 2 , so that most chromogens of this class give rise to two colours according as one or other of the foregoing radicles enters into their composition.
Phosphines, Arsines, Stibines, and Bismuthines.
These compounds are analogues of the amines, and contain the above elements in place of nitrogen. Some of the methods employed in their production are seen in the following reactions :— [table]
The corresponding tertiary stibines and bismuthines are produced by a similar reaction.
The compounds of the presen group are possessed of great affinity for oxygen, many of them taking fire spontaneously when exposed to the air. Arsendimethyl or cacodyl, As 2 (CH 3 ) 4 , is a spontaneously inflammable liquid, having a most repulsive odour, produced by the action of methyl iodide on sodium arsenic alloy, or by distilling a mixture of potassium acetate and arsenious anhydride. The analogy between these compounds and the amines is still further exemplified by the following reactions :— [table]
XΓV. Amides.
The compounds of this class are most conveniently regarded as derived from acids by the substitution of NH 2 for the HO contained in the carboxyl group. It will be thus evident that monobasic acids can yield only one amide—a monamide of the form R'CO(NH 2 ). Polybasic acids, on the other hand, can have their hydroxyl partially or entirely replaced by amidogen, thus yielding normal and acid amides. The latter are known as amic acids : -- From Dibasic Acids. [table]
From Τribasix Acids. [table]
Amides are primary, secondary, or tertiary, according as one-third, two-thirds, or all the hydrogen of the ammonia is replaced by acid radicles (see p. 574).
Amides are produced—(1.) By distilling the normal ammonium salts of the corresponding acids :—
CH 3 .CO(ONH 4 ) - OH 2 = CH 3 .CO(NH 9 ) Ammonium acetate. Acetamide.
C 2 H 4 .[CO(ONH 4 )] 2 - 2OH i - C 2 H 4 .[CO(NH 2 )] 2 ; Ammonium succinate. Succinamide.
(2.) Also by the action of ammonia on the acid halides :—
CH 3 ,COC1 + 2NH 3 = CH 3 .CO(NH 2 ) + NH 4 C1. Acetyl chloride. Acetamide.
By the further action of the acid halide upon the primary monamide secondary and tertiary monamides are produced :—
CH 3 .COC1 + CH 3 .CO(NH 2 ) = (CH 3 CO) 2 (NH) + HC1 Acetyl chloride. Acetamide. Diacetamide.
(CH 3 .CO) 2 (NH) + CH 3 .COC1 = (CH 3 CO) 3 N + HC1. Diacetamide. Acetyl chloride. Triacetamide.
Diacetamide is also formed by the action of dry HC1 upon acetamide and triacetamide by heating acetonitrile (methyl cyanide) with acetic anhydride.
The formation of amides containing dyad radicles is exemplified by the equation—
C 2 H 4 (COC1) 2 + 4NH 8 = C 2 H 4 .[CO(NH 2 )] 2 + 2NH 4 C1. Succinyl chloride. Succinamide.
(3.) Amides are formed by the action of ammonia on the ethereal salts of the corresponding acids :— [table]
All the known amides are white crystalline solids, many of them possessing both acid and basic properties Boiled with acids or alkalies the primary amides regenerate their parent acids :— [table]
Distilled with P 2 O 5 primary amides are converted into nitriles :— [table]
Amic acids are intermediate between the amides and their parent acids, from which they are produced by submitting the acid ammonium salts to dry distillation :— [table]
Diamides of dibasic acids must be distinguished from secondary monamides containing dyad acid radicles. The latter are termed imides. Thus,— [table]
Ethereal salts of amic acids have been obtained by various methods. The following are a few examples :— [table]
The salts formulated above, containing the radicle in the carboxyl group, are neutral; those in which the radicle replaces amido-hydrogen are acid. The ammonium salt of carbamic acid is formed by the direct action of carbon dioxide on ammonia: CO 2 + 2NH 3 = CO"(ONH 4 )(NH 2 ). The neutral ethereal salts of carbamic acid are known as urethanes. Alkalamides are compounds intermediate between amines and amides, that is, containing both positive and negative radicles. The following are examples :— [table]
It has been previously mentioned that cyanic acid forms two classes of compounds analogous to the nitriles and carbamines. Sulphocyanic acid forms two analogous series. The ethereal salts of the iso-series derived from these acids can be formulated as alkalimides :— [table]
Allyl zsosulphocyanate, N(CS)"(C 3 H 5 ), is interesting as being obtained from oil of mustard.
Carbamide and Urea, CO(NH 2 ) 2 .—These compounds are the neutral amides of carbonic acid. Urea is the chi er solid constituent of human urine, and is interesting a» being the first organic compound synthesized (Wöhler, 1828). It was first produced artificially by heating a. solution of ammonium cyanate: CNO(NH 4 ) = CO(NH 2 ) 2 .
Urea may also be obtained by the action of ammonia on carbon oxydichloride: COC1 2 + 4NH 3 = CO(NH 2 ) 2 + 2NH 4 C1; by the action of ammonia on ethyl carbonate: CO(OC 2 H 5 ) 2 + 2NH 3 = CO(NH 2 ) 2 + 2C 2 H 5 .I1O; and by the action of heat on ammonium carbamate and carbonate: 2CO(NH 2 )(ONH 4 ) = CO(NH 2 ) 2 + (NH 4 ) 2 CO 3 , and. (NH 4 ) 2 CO 3 - 2OH 2 = CO(NH 2 ) 2 .
Urea crystallizes in long white 4-sided prisms, very soluble in water and alcohol. It combines with acids, forming compounds like CO(NH 2 ) 2 .HC1, CO(NH 2 ) 2 .HNO 3 ^ and with metallic oxides, forming such compounds as CO(NH 2 ) 2 .2HgO . Heated with water to a very high temperature, it decomposes into carbon dioxide and ammonia: CO(NH 2 ) 2 + OH 2 = CO 2 + 2NH 3 ; by the action of nitrous acid it is decomposed into carbon dioxide,, nitrogen, and water: CO(NH 2 ) 2 + N 2 O 3 = CO 2 + 2OH 2 + 2N 2 . Urea is also decomposed by chlorine with the formation of cyanuric acid, &c.: 6CO(NH 2 ) 2 + 3Cl 2 = 2H 3 C 3 N 3 O 3 + 4NH 4 C1 + N 2 + 2HC1.
Urea combines with aldehydes, with the elimination of water, giving rise to a series of compounds termed ureides. By the action of heat on ammonium sulphocyanate, sulphο-urea, CS(NH 2 ) 2 , is obtained.
Compound ureas are derived from urea by the replacement of amido hydrogen by hydrocarbon radicles. Som¢ of the methods employed in their production are shown by the following equations :— [table]
Uric Acid (C 5 N 4 H 4 O 3 ) is contained in human urine∙ r serpents’ excrement, guano, &c. It is a dibasic acid yielding a large number of derivatives by oxidation. &c. It has never been obtained artificially.
Unclassified Organic Compounds.
These are substances of which the formulae have not yet been satisfactorily determined, owing to their complexity of composition and to the difficulty of obtaining them in a state of purity. The names and sources of some of these compounds are given below.
Mucilage, abundant in linseed, mallow root, &c.; allied to starch and gum; soluble in cold water.
Gum-tragacanth, a kind of mucilage, insoluble in water,, but becoming soft and gelatinous.
Pectin (C 32 H 48 O 32 ?), a transparent jelly contained in most ripe fruits; soluble in water, precipitated by alcohol. It is probably allied to the carbohydrates (p. 564).
Resins, allied to the terpenes. Common resin (colophony),. from pine, contains two acids, abietic acid (C 44 H 64 O 5 ) and pinic acid (C 20 H 30 O 2 ), the former crystalline, the latter amorphous. Other well-known resins are lac, mastic r copal, sandarac, dragon's blood, <kc. Amber is probably a fossil resin.
Caoutchouc (India-rubber) and gutta-percha are the hardened juices of species of Ficus, Euphorbia, Isonandra, &c. These substances appear to consist of a mixture of terpenes (p. 559).
Balsams are natural mixtures of resins with volatile oils.
Bitumens result from the decay of vegetable matter out of contact with the air. Pit-coal, lignite, and jet are allied to bitumen. Petroleum and rock oil have been referred to as paraffins (p. 556).
The remaining compounds we have to notice are of animal origin.
Albuminoid Substances or Proteids. — These substances contain carbon, hydrogen, nitrogen, oxygen, and small quantities of sulphur (0 ’ 8 to 1 ’ 6 per cent.). Serum of blood is an albuminous substance, soluble in water, precipitated by alcohol, and coagulated by heating. Eggalbumin is the chief constituent of white of egg; it differs from blood serum in being precipitated by ether. Globulin is present in blood serum, connective tissue, &c.; it is insoluble in water, soluble in dilute acids and alkalies. Myosin, allied to globulin, exists in muscular tissue, and vitellιn in yolk of egg. Casein is an alkaline albuminate which occurs in milk, coagulable by certain animal membranes (such as rennet or stomach of calf), with the formation of cheese. Fibrin is the proteid which causes the clotting of blood; it is an elastic substance of filamentous structure, insoluble in water at ordinary temperatures. Peptones result from the action of gastric juice upon albuminous substances; they are soluble in water, and are not coagulated by boiling.
Haemoglobin (which, with the other substances to be mentioned, is allied to the proteids) is the chief constituent of the red corpuscles of the blood of vertebrate animals; it can be obtained in the crystalline form from the blood of certain animals, and contains, in addition to the usual elements composing proteids, about 0 4 per cent, of iron. Haematin results from the decomposition of haemoglobin.
Pepsine is the albuminous constituent of gastric juice, which converts proteids into peptones.
Mucin (Mucus) is a secretion from the salivary glands, and from all mucous membranes of the animal body.
Gelatin and Chondrin. — The first of these substances is produced by the action of boiling water on membranous tissue, skin, tendons, and bones. Isinglass is nearly pure gelatin from the swim-bladder of the sturgeon and other fishes. Common size and glue are examples of impure gelatin. Gelatin is soluble in water and precipitated by alcohol; the aqueous solution is completely precipitated by tannic acid (p. 570), this last reaction being the basis of the process of tanning. When hides are steeped in an infusion of oak bark or of other substances yielding tannic acid, the gelatine of the hide by the action of the tannin is converted into leather. Chondrin is a gelatinous substance, obtained from the cartilage of the joints and ribs, which differs in some respects from gelatin.
Keratin is a name given to several substances left after the complete exhaustion of horn, nails, feathers, hair, <fcc., by various solvents.
Protagon is a phosphoretted fatty substance, forming the chief constituent of nerve tissue. Boiled with baryta water, it decomposes into stearic acid, glycerin, phosphoric acid, and other substances, among which is included trimethyl-oxethyl-amm onium hydrate or neurine, N(CH 3 ) 3 (C 2 H 4 .HO)HO. (κ. me.) Index of Principal Subjects. [table]
