Part 5

Lavoisier’s idea, that oxygen was the necessary constituent of all acids, began about this time to lose ground. For Davy had proved the elementary nature of chlorine; and hydrochloric acid, one of the strongest, was thus seen to contain no oxygen, and Davy expressed the view, founded on his observation, that iodic “acid,” I2O5, was devoid of acid properties until dissolved in water, and that the essential constituent of all acids was hydrogen, not oxygen. The bearing of this theory on the dualistic theory is, that while, _e.g._, sulphuric acid was regarded by Berzelius as SO3, containing no hydrogen, and was supposed to be separated as such at the positive pole of a battery, Davy’s suggestion led to the opposite conclusion that the formula of sulphuric acid is H2SO4, and that by the current it is resolved into H2 and SO4. Faraday’s electrolytic law, that when a current is passed through electrolytes in solution the elements are liberated in quantities proportional to their equivalents, led to the abandonment of the dualistic theory. For when a current is passed in succession through acidified water, fused lead chloride, and a solution of potassium sulphate, the quantities of hydrogen and oxygen from the water, of lead and chlorine from the lead chloride, and the potassium of the sulphate are in accordance with Faraday’s law. But in addition to the potassium there is liberated at the same pole an equivalent of hydrogen. Now, if Berzelius’s theory be true, the products should be SO3 and K2O, but if the opposite view be correct, then K2 is liberated first and by its subsequent action on water it yields potash and its equivalent of hydrogen. This was pointed out first by Daniell, professor at King’s College, London, and it was regarded as a powerful argument against Berzelius’s system. In 1833, too, Graham investigated the phosphoric acids, and prepared the salts of three, to which he gave the names, _ortho-_, _pyro-_, and _meta-_ phosphoric acids. To understand the bearing of this on the doctrine of dualism it must be remembered that P2O5, pentoxide of phosphorus, was at that date named phosphoric acid. When dissolved in water it reacts with bases, forming salts—the phosphates. But the quantity of water necessary was not then considered essential; Graham, however, showed that there exist three series of salts—one set derived from P2O5,3H2O, one from P2O5,2H2O, and a third from P2O5,H2O. His way of stating the fact was that water could play the part of a base; for example, the ordinary phosphate of commerce possessed, according to him, the formula P2O5,2Na2O,H2O, two-thirds of the “water of constitution” being replaced by oxide of sodium. Liebig, then professor at Giessen (1803–1873), founded on these and on similar observations of his own the doctrine of poly-basic acids—acids in which one, two, three, or more atoms of hydrogen were replaceable by metals. Thus, instead of writing, as Graham did, P2O5,2Na2O,H2O, he wrote, PO4Na2H; and for orthophosphoric acid PO4H3. The group of atoms (PO4), therefore, existed throughout the whole series of orthophosphates, and could exist in combination with hydrogen, with hydrogen and metals, or with metals alone. Similarly the group (P2O7) was characteristic of pyrophosphates and (PO3) of metaphosphates, for P2O5,2H2O=(P2O7)H4; and P2O5,H2O=2(PO3)H.

The first clear ideas of the structure of the molecule were, however, gained from the study of the compounds of carbon. It was difficult to apply the dualistic theory to them. For few of them are electrolytes, and therefore their products of electrolysis, being non-existent, could not be classified. Nevertheless, Gay-Lussac regarded alcohol, C2H6O, as a compound of C2H4, ethylene, and H2O, water; and oxalic acid (anhydrous), C2O3, as one of CO2 with CO. The discovery of “isomeric compounds,” _i.e._, of compounds which possess the same ultimate formula and yet differ entirely in their properties, forced upon chemists the necessity of attending to the structure of the molecule; for only by such a supposition could the difference between two isomeric bodies be explained. In 1823 Liebig discovered that silver fulminate and silver cyanate both possessed the empirical formula AgCNO; in 1825 this was followed by the discovery by Faraday that oil gas contains a hydrocarbon identical in composition with ethylene, C2H4, yet differing from it in properties; and in 1829 Wöhler, professor in Göttingen (1800–1882), discovered that urea, a constituent of urine, could be produced by heating ammonium cyanate, NH4CNO, a substance of the same formula. It therefore became clear that the identity of a compound must depend on some other cause than its ultimate composition.

In 1833 Liebig and Wöhler took an important step in elucidating this question by their investigations on benzoic acid and acid obtainable by distilling a resin named gum benzoin. They showed that this acid, C7H6O2, could be conceived as consisting of the group C7H5O, to which they gave the name “benzoyl,” in combination with OH; that benzoic aldehyde, C7H6O, might be regarded as its compound with hydrogen; that it also formed compounds with chlorine, and bromine, and sulphur, and replaced hydrogen in ammonia (C7H6O,NH2). They termed this group, benzoyl, a “compound element” or a “radical.” This research was followed by one by Robert Bunsen, professor at Heidelberg, born in 1811, and recently (1899) dead, which bore reference to cacodyl, a compound of arsenic, carbon and hydrogen, in which the idea of a radical was confirmed and amplified.

The idea of a radical having thus become established, Jean Baptiste Andrée Dumas, professor in Paris (1800–1884), propounded the theory of “substitution,” _i.e._, that an element such as chlorine or oxygen (which, be it noticed, is electro-negative on Berzelius’s scale) could replace hydrogen in carbon compounds, atom for atom, the resulting compound belonging to the same “type” as the one from which it was derived. And Laurent, warden of the mint at Paris (1807–1853), and Gerhardt, professor at Montpelier and at Strasburg (1816–1856), emphasized the fact that one element, be it what it may, can replace another without fundamentally altering its chemical character, and also that an atom of hydrogen can be replaced by a group of atoms or radical, behaving for the occasion like the atom of an element. It is to Laurent and Gerhardt that we owe the definition of an atom—_the smallest quantity of an element which can be present in a compound_; an equivalent—_that weight of an element which combines with or replaces one part by weight of hydrogen_; and a molecule—_the smallest quantity which can exist in a free state, whether of an element or a compound_. They recognized, too, that a molecule of hydrogen, chlorine, etc., consists of two atoms.

In 1849 Wurtz, professor in Paris (1817–1884), and Hofmann, then professor in the College of Chemistry in London, afterwards at Berlin (1818–1892), discovered a series of compounds allied to ammonia, NH3, in which one or more atoms of hydrogen were replaced by a group or radical, such as methyl (CH3), ethyl (C2H5), or phenyl (C6H5). Wurtz referred such compounds to the ammonia “type.” They all resemble ammonia in their physical properties—smell, taste, etc.—as well as in their power of uniting with acids to form salts resembling ammonium chloride (NH4Cl), and other ammonium compounds. Shortly afterwards Williamson, professor at University College, London, added the “water type,” in consequence of his researches on “mixed ethers”—bodies in which the hydrogen of water might be regarded as replaced by organic radicals. Thus we have the series:

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

H H H / \ / H——Cl; H——O——H; H——N ; C \ / \ H H H

Instances of penta, hexa, and even hepta-valency are not wanting.

This was the key to unlock the structure of chemical compounds; and Frankland’s views, just stated, are still held by chemists. The determination of the constitution of compounds, chiefly those of carbon, occupied the attention of chemists, almost exclusively, until 1880. The plan of action is much the same as that of a mechanician who wishes to imitate a complicated mechanism. He must first dissect it into groups of mechanical contrivances; these are next constructed; and they are finally built together into the complete machine. In certain cases the atoms of carbon are arranged in “chains,” as, for example, in pentyl alcohol:

H3C——C——C——C——C——O——H H2 H2 H2 H2

each atom being tetrad, and its “affinities,” or powers of combination, saturated either with hydrogen or with those of neighboring atoms of carbon; in others they are in the form of a “ring,” as in benzene, the formula of which was first suggested by Kekulé, viz.:

H H C——C / \ HC CH; \ / C==C H H

or in both, as in ethyl benzene,

H H C——C / \ H H HC C——C——CH. \ / H H C==C H H

One or more atoms of nitrogen, or of oxygen, may form part of the circle, as in pyridine:

H H H H C——C C C / \ / N CH and furfurane, O == , \ / \ C==C C C H H H H

and so on. By means of conceptions such as these many interesting compounds have been built up out of the elements which they contain; _e.g._, urea and uric acid, constituents of urine; theobromine and caffeine, the essential principles of cocoa and tea; alizarine and indigo, valuable dyestuffs; and several of the alkaloids, bitter principles contained in plants, of great medicinal value.

They have led, too, to the discovery of many brilliant colors, now almost universally employed, to the exclusion of those less brilliant, because less pure, derived from plants, and in one or two cases from animals; the manufacture of gun-cotton, dynamite, and similar high explosives; and to the development of the candle industry; the sugar manufacture; to improvement in tanning, in brewing, and in the preparation of gas and oils for illuminating purposes. In short, it may be said that the industrial progress of the latter half of the century has been due to the theoretical views of which a short sketch has just been given.

Such formulas, however, can evidently not represent the true constitution of matter, inasmuch as the atoms are imagined to lie on a plane, whereas it is evident that they must occupy space of three dimensions and possess the attributes of solidity. The conception which led to the formulation of such views was due first to Pasteur, in his later years director of the institute known by his name at Paris, and more directly to LeBel and Van’t Hoff, now professor at Berlin, independently of each other. In 1848 Pasteur discovered that it was possible to separate the two varieties of tartaric acid from each other; and that that one which rotated the plane of polarized light to the right gave crystals with an extra face, unsymmetrically disposed with regard to the other faces of the crystal. The variety, the solution of which in water was capable of producing left-handed rotation, also possessed a similar face, but so placed that its reflection in a mirror reproduced the right-handed variety. Pasteur also showed that a mixture of these acids gave crystals not characterized by an unsymmetrically placed face; and also that the solution was without action on polarized light. These observations remained unexplained, until LeBel and Van’t Hoff, in 1874, simultaneously and independently devised a theory which has, up till now, stood the test of research. It is briefly this: Imagine two regular tetrahedra, or three-sided pyramids, standing each on its triangular base. An idea can best be got by a model, easily made by laying on a table three lucifer matches so as to form an equilateral triangle, and erecting a tripod with three other matches, so that each leg of the tripod stands on one corner of the triangle. At the centre of such a tetrahedron, an atom of carbon is supposed to be placed. Marsh gas, CH4, is supposed to have such a structure, each corner, or solid angle of the structure (of which there are four), being occupied by an atom of hydrogen. This represents the solid or _stereochemical_ formula of methane or marsh gas. Now, suppose one of the atoms of hydrogen in each of these structures to be replaced by chlorine, the group (OH), or any other monovalent element or group. It is evident that if not exactly similar (owing to the replacement not having been made at similar corners in each), the two structures can be made similar by turning one of them round, until the position of the substituting atom or group (which we will term X) coincides in position with X in the stationary one. If two such replacements be made, say, with X and Y in each, coincidence can again be made to take place; but the same is not the case if X, Y, and Z replace three atoms of hydrogen in the structure; for there is one way of replacement which is the optical image of the other, and represents the other’s reflection in a mirror.

Now, it is found that when the four corners of such a structure are occupied by four separate atoms or groups, or when (as the expression goes) the body contains an “asymmetrical carbon atom,” if the substance or one of its derivations can be obtained in a crystalline form, the crystals are also asymmetric, _i.e._, arc develops a face which is the mirror-reflection of a similar face developed on the other variety; and if a beam of polarized light be passed through the solution of the substance, its plane is rotated to the left if one variety be used, and, if the other, to the right. This hypothesis of LeBel’s and Van’t Hoff’s has had an enormous influence on the progress of organic chemistry. By its means Fischer, now professor at Berlin, has explained the reason of the existence of the enormous number of bodies analogous to grape and cane sugar, and has prepared many new varieties; and it appears likely that the terpenes, a class of bodies allied to turpentine, and comprising most of the substances to which the odor of flowers is due, may thereby find their explanation. It may be mentioned in passing that Pasteur, having found that ordinary mould destroyed one variety of tartaric acid rather than the other in a mixture of the two, and made use of this observation in order to prepare the unattached variety in a state of purity, was led to study the action of organisms more or less resembling mould; and that this has led to the development of the science of bacteriology, which has had an enormous influence on our views regarding fermentation in general, and guides the work of our physicians, our surgeons (witness Lister’s antiseptic treatment), our sanitary engineers in their estimate of the purity of drinking-water and of the disposal of sewage, of our manufacturers of beer and spirits, of wine-growers, and more recently of farmers. All these processes depend upon the action of organisms in producing chemical changes, whether in the tissues of the body, causing or curing disease, or in the production of flavored alcohol from sugar, or in the manufacture of butter and cheese, or in preparing the land for the reception of crops. We also owe to the genius of Van’t Hoff the most important advance of recent times in the region of physical chemistry. It has been observed by Raoult, professor at Grenoble, that the freezing-point of a solvent as a general rule is lowered to the same extent if there be dissolved in it quantities of substances proportional to their molecular weights. Thus, supposing 1.80 grams of grape-sugar be dissolved in 100 grams of water and the solution cooled below 0° with constant stirring, ice separates suddenly in thin spicules, and the temperature rises to −0.185°. If 3.42 grams of cane-sugar be similarly dissolved in 100 grams of water, the freezing-point of the solution is again −0.185°. Now, 1.80 and 3.42 are respectively the hundredth part of the molecular weights of grape-sugar (C6H12O6) and cane-sugar (C12H22O11). Similarly, Raoult found that quantities proportional to molecular weights dissolved in a solvent depress the vapor pressure of that solvent equally, or, what comes to the same thing, raise its boiling-point by an equal number of degrees. But ordinary salts, such as sodium chloride, potassium nitrate, etc., dissolved in water, give too great a depression of the freezing-point and too high a boiling-point. Next, it has been observed by botanists, Devries, Pfeffer, and others, who had examined the ascent of sap in plants, that if a vessel of unglazed porcelain, so treated as to cause a film of cupric ferrocyanide (a slimy red compound) to deposit in the pores of its walls, be filled with a weak (about 1 per cent.) solution of sugar or similar substance, and plunged in a vessel of pure water, water entered through the pores. By attaching a monometer to the porous vessel the pressure exerted by the entering water could be measured. Such pressure was termed “osmotic pressure,” referring to the “osmosis” or passage through the walls of the vessel. Such prepared walls are permeable freely to water, but not to sugar or similar bodies. Van’t Hoff pointed out that the total pressure registered is proportional to the amount of substance in solution, and that it is proportional to the absolute temperature, and he showed, besides, that the pressure exerted by the sugar molecules is the same as that which would be exerted at the same temperature were an equal number of molecules of hydrogen to occupy the same volume as the sugar solution. This may be expressed by stating that when in dilute solution sugar molecules behave _as if_ they were present in the gaseous state. Here again, however, it was noticed that salts tended to give a higher pressure; it was difficult to construct a semi-permeable diaphragm, however, which would resist the passage of salt molecules, while allowing those of water to pass freely. Lastly, Arrhenius, of Stockholm, had shown that the conductivity of salt solutions for electricity may be explained on the assumption that when a salt, such as KNO3 is dissolved in water, it dissociates into portions similar in number and kind to those it would yield if electrolyzed (and if no secondary reactions were to take place). Such portions (K and NO3, for example) had been named ions by Faraday. The conductivity of such solutions becomes greater, per unit of dissolved salt, the weaker the solution, until finally a limit is reached, after which further dilution no longer increases conductivity. Now Van’t Hoff united all these isolated observations and showed their bearing on each other. Stated shortly, the hypothesis is as follows: When a substance is dissolved in a large quantity of a solvent, its molecules are separated from each other to a distance comparable with that which obtains in gases. They are, therefore, capable of independent action; and when placed in a vessel the walls of which are permeable to the solvent, but not to the dissolved substance (“semi-permeable membrane”), the imprisoned molecules of the latter exert pressure on the interior surface of these walls as if they were gaseous. Van’t Hoff showed the intimate connection between this phenomenon and the depression of freezing-point and the use of vapor pressure already alluded to. He pointed out further that the exceptions to this behavior, noticed in the case of dissolved salts, are due to their “electric dissociation,” or “ionization,” as it is now termed; and that in a sufficiently dilute solution of potassium nitrate, for example, the osmotic pressure, and the correlated depression of freezing-point and rise of boiling-point, are practically equal to what would be produced were the salt to be split into its ions, K and NO3. These views were vigorously advocated by Ostwald, professor at Leipzig, in his _Zeitschrift für physikalische Chemie_, and he and his pupils have done much to gather together facts in confirmation of this theory, and in extending its scope.

It must be understood that the ions K and NO3 are not, strictly speaking, atoms; they are _charged_ atoms; the K retains a +, and the NO3 a − charge. On immersing into the solution the poles of a battery, one charged + and the other −, the + K atoms are attracted to the − pole, and are there discharged; as soon as they lose their charge they are free to act on the water, when they liberate their equivalent of hydrogen. Similarly, the − NO3 groups are discharged at the + pole, and abstract hydrogen from the water, liberating an equivalent quantity of oxygen. Thus the phenomenon of electrolysis, so long a mysterious process, finds a simple explanation. The course of ordinary chemical reactions is also readily realized when viewed in the light of this theory. Take, for example, the ordinary equation:

AgNO3.Aq + NaCl.Ag = AgCl + NaNO3.Aq;

_i.e._, solutions of silver nitrate and sodium chloride give a precipitate of silver chloride, leaving sodium nitrate in solution. By the new views, such an equation must be written:

+ − + − + − Ag.Aq + NO3.Aq + Na.Aq + Cl.Aq = AgCl + Na.Aq + NO3.Aq.

The compound, silver chloride, being insoluble in water, is formed by the union of the ions Ag and Cl, and their consequent discharge, forming an electrically neutral compound; while the sodium ions, charged positively together with the NO3 ions, negatively charged, remain in solution.

One more application of the principle may be given. Many observers—Andrews, Favre, and Silbermann, but especially Julius Thomsen, of Copenhagen, and M. Berthelot, of Paris—have devoted much labor and time to the measurement of the heat evolved during chemical reactions. Now, while very different amounts of heat are evolved when chlorine, bromine, or iodine combine respectively with sodium or potassium, the number of heat units evolved on neutralizing sodium or potassium hydroxide with hydrochloric, hydrobromic, hydriodic, or nitric acids is always about 13,500. How can this fact be explained? It finds its explanation as follows: These acids and bases are ionized in solution as shown in the equation:

+ − + − + − H.Aq + Cl.Aq. + Na.Aq + OH.Aq = H.OH + Na.Aq + Cl.Aq.