Organic Chemistry.

The chemistry of colloids has now assumed such importance that it may be considered as a separate branch of the science. It has its own technical journal and deals largely with the chemistry of organic products. All living matter is built up of colloids, and hæmoglobin, starch, proteins, rubber and milk are examples of colloidal substances or solutions. Among inorganic substances, many sulphides, silicic acid, and the amorphous hydroxides, like ferric hydroxide, frequently act as colloids.

Law of Mass Action.—Berthollet about the beginning of the last century was the first chemist to study the effect of mass, or more correctly, the concentration of substances on chemical action. His views summarized by himself are as follows: “The chemical activity of asubstance depends upon the force of its affinity and upon the mass which is present in a given volume.” The development of this idea, which is fundamentally correct, was greatly hindered by the fact that Berthollet drew the incorrect conclusion that the composition of chemical compounds depended upon the masses of the substances combining to produce them, a conclusion in direct contradiction to the law of definite proportions, and since this view was soon disproved by Proust and others, Berthollet’s law in its other applications received no immediate attention. Mitchell, however, pointed out in the Journal (16, 234, 1829) the importance of Berthollet’s work, and Heinrich Rose in 1842 again called attention to the effect of mass, mentioning as one illustration the effect of water and carbonic acid in decomposing the very stable natural silicates. Somewhat later several other chemists made important contributions to the question of the influence of concentration upon chemical action, but it was the Norwegians, Guldberg and Waage, who first formulated the law of mass action in 1867.

This law has been of enormous importance in chemical theory, since it explains a great many facts upon a mathematical basis. It applies particularly to equilibrium in reversible reactions, where it states that the product of the concentrations on the one side of a simple reversible equation bears a constant relation to the products of the concentrations on the other side, provided that the temperature remains constant. In cases of this kind where two gases or vapors react with two solids, the latter if always in excess may be regarded as constant in concentration, and the law takes on a simpler aspect in applying only to the concentrations of the gaseous substances. For example, in the reversible reaction

3Fe + 4H2O ⇄ Fe3O4+ 4H2,

which takes place at rather high temperatures, a definite mixture of steam and hydrogen at a definite temperature will cause the reaction to proceed with equal rapidity in both directions, thus maintaining a state of equilibrium, provided that both iron and the oxide are present inexcess. If, however, the relative concentrations of the hydrogen and steam are changed, or even if the temperature is changed, the reaction will proceed faster in one direction than in the other until equilibrium is again attained.

The principle of mass action also explains why it is sometimes possible for a reversible reaction to become complete in either direction. For instance, in connection with the reaction that has just been considered, if steam is passed over heated iron and if hydrogen is passed over the heated oxide, the gaseous product in each case is gradually carried away, and the reaction continually proceeds faster in one direction than in the other until it is complete, according to the equations

3Fe + 4H2O → 3Fe3O4+ 4H2, andFe3O4+ 4H2→ 3Fe + 4H2O.

3Fe + 4H2O → 3Fe3O4+ 4H2, and

Fe3O4+ 4H2→ 3Fe + 4H2O.

Many other well-known and important facts, both chemical and physical, depend upon this law. It explains the circumstance that a vapor-pressure is not dependent upon the amount of the liquid that is present; it also explains the constant dissociation pressure of calcium carbonate at a given temperature, irrespective of the amounts of carbonate and oxide present; in connection with the ionic theory, it furnishes the reason for the variable solubility of salts due to the presence of electrolytes containing ions in common; and it elucidates Henry’s law which states that the solubilities of gases are proportional to their pressures.

Ostwald, more than any other chemist, has been instrumental in making general applications of this law, and he made particularly extensive use of it in connection with analytical chemistry in a book upon this subject which he published.

The Phase Rule.—In 1876 Willard Gibbs of Yale published a paper in the Proceedings of the Connecticut Academy of Science on the “Equilibrium of Heterogeneous Substances,” and two years later he published an abstract of the article in the Journal (16, 441, 1878). He had discovered a new law of nature of momentous importance and wide application which is called the“Phase-Rule” and is expressed by a very simple formula.

The application of this great discovery to chemical theory was delayed for ten years, partly, perhaps, because it was not sufficiently brought to the attention of chemists, but largely it appears because it was not at first understood, since its presentation was entirely mathematical.

It was Rooseboom, a Dutch chemist, who first applied the phase-rule. It soon attracted profound attention, and the name of Willard Gibbs attained world-wide fame among chemists. When Nernst, who is perhaps the most eminent physical chemist of the present time, was delivering the Silliman Memorial Lectures at Yale a few years ago, he took occasion to place a wreath on the grave of Willard Gibbs in recognition of his achievements.

To understand the rule, it is necessary to define the three terms, introduced by Gibbs,phase,degrees of freedomandcomponent.

By the first term, is meant the parts of any system of substances which are mechanically separable. For instance, water in contact with its vapor has two phases, while a solution of salt and water is composed of but one. The degrees of freedom are the number of physical conditions, including pressure, temperature and concentration, which can be varied independently in a system without destroying a phase. The exact definition of a component is not so simple, but in general, the components of a system are the integral parts of which it is composed. Any system made up of the compound H2O, for instance, whether as ice, water or vapor, contains but one component, while a solution of salt and water contains two. Letting P, F, and C stand for the three terms, the phase-rule is simply

F = C + 2 − P

that is, the number of degrees of freedom in a system in equilibrium equals the number of components, plus two, minus the number of phases. The rule can be easily understood by means of a simple illustration. In a system composed of ice, water and water vapor, there are three phases and one component and therefore

F = 1 + 2 − 3 = 0

Such a system has no degrees of freedom. This means that no physical condition, pressure or temperature can be varied without destroying a phase, so that such a system can only exist in equilibrium at one fixed temperature, with a fixed value for its vapor-pressure.

J. William Gibbs

For instance, if the system is heated above the fixed temperature, ice disappears and if the pressure is raised, vapor is condensed. If this same system of water alone contains but two phases, for instance, liquid and vapor, F = 1 + 2 − 2 = 1, or there is one degree of freedom. In such a system, one physical condition such as temperature can be varied independently, but only one, without destroying a phase. For instance, the temperature may be raised or lowered, but for every value of temperature there is a corresponding value for the vapor-pressure. One is a function of the other. If both values are varied independently, one phase will disappear, either vapor condensing entirely to water or the reverse. Finally if the system consists of one phase only, as water vapor, F = 2, or the system is divariant, which means that at any given temperature it is possible for vapor to exist at varying pressures.

The illustration which has been given relates to physical equilibrium, but the rule is applicable to cases involving chemical changes as well. In comparing the phase-rule with the law of mass action, it will be noticed that both have to do with equilibrium. The great advantage of the former is that it is entirely independent of the molecular condition of the substances in the different phases. For instance, it makes no difference so far as the application of the rule is concerned, whether a substance in solution is dissociated, undissociated or combined with the solvent. In any case, the solution constitutes one phase. On the other hand, the rule is purely qualitative, giving information only as to whether a given change in conditions is possible. The law of mass action is a quantitative expression so that when the value of the constant is once known, the change can be calculated which takes place in the entire system if the concentration of one substance is varied. The law, however, requires a knowledge of the molecular condition ofthe reacting substances, which may be uncertain or unknown, and chiefly on this account it has, like the phase-rule, often only a qualitative significance.

The phase rule has served as a most valuable means of classifying systems in equilibrium and as a guide in determining the possible conditions under which such systems can exist. As illustrations of its practical application, van’t Hoff used it as an underlying principle in his investigations on the conditions under which salt deposits have been formed in nature, and Rooseboom was able by its means to explain the very complicated relations existing in the alloys of iron and carbon which form the various grades of wrought iron, steel and cast iron.

Thermochemistry.—This branch of chemistry has to do with heat evolved or absorbed in chemical reactions. It is important chiefly because in many cases it furnishes the only measure we have of the energy changes involved in reactions. To a great extent, it dates from the discovery by Hess in 1840 of a fundamental law which states that the heat evolved in a reaction is the same whether it takes place in one or in several stages. This law has made it possible to calculate the heat values of a large number of reactions which cannot be determined by direct experiment.

Thermochemistry has been developed by a comparatively few men who have contributed a surprisingly large number of results. Favre and Silbermann, beginning shortly after 1850, improved the apparatus for calorimetric determinations, which is called the calorimeter, and published many results. At about the same time Julius Thomsen, and in 1873 Berthelot, began their remarkable series of publications which continued until recently. Thomsen’s investigations were published in 1882 in4volumes. It is probably safe to say that the greater part of the data of thermochemistry was obtained by these two investigators. The bomb calorimeter, an apparatus for determining heat values by direct combustion, was developed by Berthelot. The recent work of Mixter at Yale, published in the Journal, and of Richards at Harvard should be mentioned particularly. Mixter’s work in this field began in 1901 (12, 347). Using an improved bomb calorimeter, he has developed amethod of determining the heats of formation of oxides by combustion with sodium peroxide. By this same method as well as by direct combustion in oxygen, he has obtained results which appear to equal or excel in accuracy any which have ever been obtained in his field of work. Richards’s work has consisted largely of improvements in apparatus. He developed the so-called adiabatic calorimeter which practically eliminates one of the chief errors in thermal work caused by the heating or cooling effect of the surroundings. This modification is being generally adopted where extremely accurate work is required.

One hundred years ago qualitative tests for a few organic compounds were known, the elements usually occurring in them were recognized, and some of them had been analyzed quantitatively, but organic chemistry was far less advanced than inorganic, and almost the whole of its enormous development has taken place during our period.

Berzelius made a great advance in the subject by establishing the fact, which had been doubted previously, that the elements in organic compounds are combined in constant, definite proportions. In 1823 Liebig brought to light the exceedingly important fact of isomerism by showing that silver fulminate had the same percentage composition as silver cyanate, a compound of very different properties. Isomeric compounds with identical molecular weight as well as the same composition have since been found in very many cases, and they have played a most important part in determining the arrangements of atoms in molecules. They have been found to be very numerous in many cases. For instance, three pentanes with the formula C5H12are known, all that are possible according to theory, and in each case the structure of the molecule has been established. On theoretical grounds it has been calculated that 802 isomeric compounds with the formula C13H28are possible, while with more complex formulas the numbers of isomers may be very much greater.

A particularly interesting case of isomerism was observed by Wöhler in 1828, when he found that ammonium cyanate changes spontaneously into urea

(NH4CNO → N2H4CO).

This was the first synthesis of an organic compound from inorganic material, and it overthrew the prevailing view that vital forces were essential in the formation of organic substances. A great many natural organic compounds have been made artificially since that time, and some of them, such as artificial alizarin, indigo, oil of wintergreen, and vanillin, have more or less fully replaced the natural products. The preparation of a vast number of compounds not known in nature, many of which are of practical importance as medicines, dyes, explosives, etc., has been another great achievement of organic chemistry.

The development of our present formulas for organic compounds, by means of which in many cases the relative positions of the atoms can be shown with the greatest confidence, has been gradual. Formulas based on the dualistic idea of Berzelius were used for some time, type-formulas, with the employment of compound radicals, came later, the substitution of atoms or groups of atoms for others in chemical reactions came to be recognized, but one of the most important steps was the recognition of the quadrivalence of carbon and the general application of valency to atoms by Kekulé about 1858. This led directly to the use of modern structural formulas which have been of the greatest value in the theoretical interpretation of organic reactions. It was Kekulé also who proposed the hexagonal ring-formula for benzene, C6H6, which led to exceedingly important theoretical and practical developments. The details of the formulas for many other rings and complex structures have been established since that time, and there is no doubt that the remarkable achievements in organic chemistry during the past sixty years have been much facilitated by the use of these formulas.

Many important researches in organic chemistry have been carried out in the United States, and the activity in this direction has greatly increased in recent years. Inthis connection the large amount of work of this kind accomplished in the Sheffield Laboratory, at present under the guidance of Professor T. B. Johnson, should be mentioned.

It has happened that comparatively few publications on organic chemistry have appeared in the Journal, but it may be stated that the preparation of chloroform and its physiological effects were described by Guthrie (21, 64, 1832). Unknown to him, it had been prepared by Souberain, a French chemist, the previous year, but the former was the first to describe its physiological action. Silliman gave a sample to Doctor Eli Ives of the Yale Medical School, who used it to relieve a case of asthma. This was the first use of chloroform in medical practice (21, 405, 1832). Guthrie also described in the Journal (21, 284, 1832) his new process for converting potato starch into glucose, a method which is essentially the same as that used to-day in converting cornstarch into glucose. Lawrence Smith (43, 301, 1842et seq.), Horsford (3, 369, 1847et seq.), Sterry Hunt (7, 399, 1849), Carey Lea (26, 379, 1858et seq.), Remsen (5, 179, 1873et seq.), and others have contributed articles on organic chemistry.

Until near the middle of the nineteenth century, it was believed that plants, like animals, used organic matter for food, and depended chiefly upon the humus of the soil for their growth. This view was held even long after it was known that plant leaves absorb carbon dioxide and give off oxygen, and after the ashes of plants had been accurately analyzed.

This incorrect view was overthrown by the celebrated German chemist, Liebig, who made many investigations upon the subject, and, properly interpreting previous knowledge, published a book in 1840 upon the application of chemistry to agriculture and physiology in which he maintained that the nutritive materials of all green plants are inorganic substances, namely, carbon dioxide, water, ammonia (nitrates), sulphates, phosphates, silica, lime, magnesia, potash, iron, and sometimes common salt. He drew the vastly important conclusion that the effectivefertilization of soils depends upon replenishing the inorganic substances that have been exhausted by the crops.

The fundamental principles set forth by Liebig have been confirmed, and it has been found that the fertilizing constituents most commonly lacking in soils are nitrogen compounds, phosphates, and potassium salts, so that these have formed the important constituents of artificial fertilizers. Liebig himself found that humus is valuable in soils, because it absorbs and retains the soluble salts.

The foundation established by Liebig in regard to artificial fertilizers has led to an enormous application of these materials, much to the advantage of the world’s food supply.

It was Liebig’s belief, in accordance with the prevailing views, that decay and putrefaction as well as alcoholic and other fermentations were spontaneous processes, and when the eminent French chemist, Pasteur, in 1857, explained fermentation as directly caused by yeast, an epoch-making discovery which led to the explanation of decay and putrefaction by bacterial action and to the germ-theory of disease, the explanation was violently opposed by Liebig and other German chemists. Pasteur’s view prevailed, however, and since that time it has been found that various kinds of bacteria are responsible for the formation of ammonia from nitrogenous organic matter and also for the change of ammonia into the nitrates that are available as plant-food.

The long-debated question as to the availability of atmospheric nitrogen for plant-food was settled in 1886 by the discovery of Hellriegel that bacteria contained in nodules on the roots, especially of leguminous plants, are capable of bringing nitrogen into combination and furnishing it to the plants.

No more than an allusion can be made to agricultural experiment stations where soils, fertilizers, foods and other products are examined, and where other problems connected with agriculture are studied.

The late S. W. Johnson of Yale studied with Liebig and subsequently did much service for agricultural chemistry in this country, by his investigations, his teaching, and his writings. His book, “How Crops Grow,” publishedin 1868, gave an excellent account of the principles of agricultural chemistry. He did much to bring about the establishment of agricultural experiment stations in this country, and for a long time he was the director of the Connecticut Station.

In the Journal, as early as 1827, Amos Eaton (12, 370) published a simple method for the mechanical analysis of soils to determine their suitability for wheat-culture, and Hilgard, between 1872 and 1874, described an elaborate study of soil-analysis. J. P. Norton, a Yale professor, in 1847 (3, 322) published an investigation on the analysis of the oat, which was awarded a prize of fifty sovereigns by a Scotch agricultural society, while Johnson, Atwater, and others have contributed articles on the analysis of various farm products.

One hundred years ago sulphuric acid was manufactured on a comparatively very small scale in lead chambers. In 1818, an English manufacturer of the acid introduced the modern feature of using pyrites in the place of brimstone, while the Gay-Lussac tower in 1827 and the Glover tower in 1859 began to be applied as great improvements in the chamber process. Within about twenty years the contact process, employing platinized asbestos, has replaced the old chamber process to a large extent. It has the advantage of producing the concentrated acid, or the fuming acid, directly.

During our period the manufacture of sulphuric acid has increased enormously. Very large quantities of it have been used in connection with the Leblanc soda process in its rapid development. It came to be employed extensively for absorbing ammonia in the illuminating-gas industry, which was in its infancy one hundred years ago. New industries such as the manufacture of “superphosphates” as artificial fertilizers, the refining of petroleum, the manufacture of artificial dyestuffs and many other modern chemical products have greatly increased the demand for it, while its employment in the production of nitric and other acids, and for many other purposes not already mentioned, has been very great.

The manufacture of nitric acid has been greatlyextended during our period on account of its employment for producing explosives, artificial dyestuffs, and for many other purposes. Chile saltpeter became available for making it about 1852. This acid has been manufactured recently from atmospheric nitrogen and oxygen by combining them by the aid of powerful electric discharges. This process has been used chiefly in Norway where water-power is abundant, as it requires a large expenditure of energy. A still more recent method for the production of nitric acid depends upon the oxidation of ammonia by air with the aid of a contact substance, such as platinized asbestos.

The production of ammonia, which was very small a hundred years ago, has been vastly increased in connection with the development of the illuminating-gas industry and the employment of by-product coke ovens. This substance is very extensively used in refrigerating machines and also in a great many chemical operations, including the Solvay soda process. Ammonium salts are of great importance also as fertilizers in agriculture. The conversion of atmospheric nitrogen into ammonia on a commercial scale is a recent achievement. It has been accomplished by heating calcium carbide, an electric-furnace product made from lime and coke, with nitrogen gas, thus producing calcium cyanamide, and then treating this cyanamide with water under proper conditions. Another method devised by Haber consists in directly combining nitrogen and hydrogen gases under high pressure with the aid of a contact substance.

Leblanc’s method for obtaining sodium carbonate from sodium chloride by first converting the latter into the sulphate by means of sulphuric acid and then heating the sulphate with lime and coal in a furnace was invented as early as 1791, but it was not rapidly developed and did not gain a foothold in England until 1826 on account of a high duty on salt up to that time. Afterwards the process flourished greatly in connection with the sulphuric acid industry upon which it depended, and with the bleaching-powder industry which utilized the hydrochloric acid incidentally produced by it, and, of course, in connection with soap manufacture and many other industries in which the soda itself was employed.

About 1866 the Solvay process appeared as a rival to the Leblanc process. This depends upon the precipitation of sodium bicarbonate from salt solutions by means of carbon dioxide and ammonia, with the subsequent recovery of the ammonia. It has displaced the older process to a large extent, and it is carried on extensively in this country, for instance, at Syracuse, New York.

Other processes for soda depend upon the electrolysis of sodium chloride solutions. In this case caustic soda and chlorine are the direct products, and the chlorine thus produced and liquefied by pressure in steel cylinders, has become an important commercial article.

In earlier times wood-ashes were the source of potash and potassium salts. Wurtz in the Journal (10, 326, 1850) suggested the availability of New Jersey greensand as a source of potash and showed how this mineral could be decomposed, but it does not appear that this mineral has ever been utilized for the purpose. About 1861 the German potash-salt deposits began to be developed, and these have since become the chief source of this material. At present many efforts are being made to obtain potassium compounds from other sources, such as brines, cement-kiln dust, and feldspar and other minerals but thus far the results have not satisfied the demand.

This account of chemical progress has given only a limited view of small portions of the subject, because the amount of available material is so vast in comparison with the space allowed for its presentation. Since the Journal has published comparatively little organic chemistry, it was decided to make room for a better presentation of other things by giving only a brief discussion of this exceedingly active and important branch of the science. For similar reasons industrial and metallurgical chemistry, and other branches besides, in spite of their great growth and importance, have been neglected, except for some incidental references to them, and some account of a few of the more important industrial chemicals.

It appears that we have much reason to be proud of theadvances in chemistry that have been made during the Journal’s period, and of the part that the Journal has taken in connection with them, and there seems to be no doubt that this progress has not diminished during more recent times.

The present tendency of chemical research is evidently towards a still greater development of organic chemistry, and an increased application of physics and mathematics to chemical theory and practice.

The very great improvements that have been made in chemical education, both in the number of students and the quality of instruction, during the period under discussion, and particularly in rather recent times, gives promise for excellent future progress.

153. It appears that the most accurate experimental demonstration ever made of this law was that of E. W. Morley, published in the Journal (41, 220, 276, 1891). He showed that 2·0002 volumes of hydrogen combine with one volume of oxygen.

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