Part 87

The “blue ground” was at first supposed to be the original home of the diamond, within which it had somehow taken its shape. But no satisfactory explanation was forthcoming as to the state of the carbon before its solidification into the crystalline form. The more general opinion has been in favour of a volcanic origin due to very high temperature; and although the “blue ground” itself is clearly not the ordinary erupted matter of volcanoes due to igneous fusion, the geology points to the district having been the scene of very active and extensive volcanic energies at more than one remote period, for the bed of the Karoo inland sea has been several times covered by level sheets of molten matter extruded somewhere from below; but not through the “pipes,” which were blown out ages afterwards. The strata of basalt and of hornblendic mineral, which extend horizontally over great areas in the Karoo formation, are of igneous origin, as are also some nearly vertical dykes of trap rock, about 7 feet wide, that are found traversing the “blue ground” in certain directions. These intrusive dykes are of course more recent than the formation of the blue ground, and that is itself later than the production of the pipes. The fact of many fragments of crystals being found in the “blue ground” does not comport with the theory that supposes it to be the matrix; and besides this, many of the diamonds show scratches, and as these are producible only by other diamonds, it would appear that they must all have travelled in company, some part of their journey at least.

Carbon in any form is quite infusible at the highest temperature we have hitherto been able to produce, although an incipient softening under the influence of the electric arc has been suspected. Professor Dewar, an English chemist, basing his data on analogies with other substances, and on purely theoretical grounds, has calculated that the melting temperature of carbon is near 3,600° C. (6,512° F.), and that it cannot remain in a liquid state at a temperature exceeding 5,527° C., when its vapour would have a tension equivalent to a pressure of 15 tons on the square inch. So far as these deductions are correct, both the melting point of carbon and the boiling point of its liquid must lie within the range of temperature expressed by 3,600° C. and 5,527° C. The most intense heat we can produce is that developed in the electric arc discharge, and an eminent French chemist and metallurgist, M. Moissan, by employing special arrangements and very powerful currents, has thus been able to obtain in his “electric furnace” a temperature estimated at 3,500° C., which nearly approaches the lower of the above-mentioned limits, and he has thereby produced many new and unexpected chemical combinations of refractory elements. Among the most striking of his results is the formation artificially of real crystalline diamonds. He found that carbon is freely dissolved by several of the metals in fusion at the temperature of the electric furnace. When the carbon separated from the metals, as they cooled and became solid, it was always in the condition of graphite. The carbons of the electric poles were readily attacked by molten iron, and it was from the solution of carbon in iron that Moissan prepared his diamonds. The fact of carbon thus combining with iron was of course no discovery, as the reader already knows; and the resulting combination was found, on allowing the metal to cool, to be simply cast iron, the greater part of the carbon separating out in the graphitic form. But M. Moissan, having studied the conditions of the Kimberley mines, and recognizing the probability of the diamonds having taken their origin at very great depths, where the pressure due to the weight of superincumbent strata would be immense, was struck with the idea of _pressure_ being in some way a factor in their formation; and it occurred to him that the carbon might separate from its liquid condition in the iron in the crystalline, and not in the graphitic form, if the solidification could be effected under great pressure. The apparently insurmountable difficulty of applying an enormous pressure to a small quantity of molten iron (half a pound) yielded to the experimenter’s ingenuity. He took advantage of the circumstance that cast iron at the moment of solidification expands, a property upon which depends its use for many purposes. If then the fused mass were suddenly cooled on the outside, we should have a shell of solid iron enclosing a nucleus of still fluid metal, which, on cooling in its turn, would tend to expand, and by so doing would exert a great pressure within the shell by which it was confined. At first Moissan plunged his glowing crucible into cold water, but a method of more rapidly cooling it was to immerse it in melted lead. It seems a strange proceeding to cool the crucible by surrounding it with hot metal, yet the _difference_ of the temperatures was sufficient to produce the desired effect, the cooling contact of water not really operating on the intensely heated body, which becomes separated from the liquid by a coating of _steam_. When the mass of iron was dissolved off, diamonds of all kinds were found in the residue, and, though extremely small, some crystals were perfect in shape and colour; every variety that occurred in the mines being found reproduced in tiny size. There was also some graphite in the residue. Many more crystals of “pure water” were obtained by the lead-cooling than by the water-cooling, as the former process gave some flawless cubes and octahedra. The largest of the set was only 1/50 inch across, and although of perfect form when first extracted, within the course of three months it had spontaneously split up into fragments.

There was evidently no danger of M. Moissan’s manufacture of diamonds from coke causing consternation at Kimberley; though it would not be without interest to speculate upon the consequences had the French _savant_ achieved the greater triumph of turning out carbon crystals in every respect equal to the productions of nature’s own laboratories. What a drop there would have been in the shares of the De Beers Mines Consolidated! What heaviness of heart would have fallen upon those great ladies who exult in the exclusive possession of priceless tiaras and precious necklaces flashing with the resplendent gems! From a scientific point of view, M. Moissan’s fabrication of even those minute crystals, which so soon spontaneously crumbled into fragments, is a distinct and valuable success; for, notwithstanding their diminutive size and instability, they show us that art has so far succeeded in imitating the processes of nature, that some of her secrets have been revealed. Though we know the exact chemical composition of all kinds of crystallized minerals, very very few of these have we been able to imitate artificially. Nor is this to be wondered at; for nature’s resources are immense compared with ours: she can command temperatures unlimited by which to form her solutions or liquefactions; prodigious pressures to keep them close; and time immeasurable—geological time—in which to let them cool, and their particles freely coalesce into geometric forms. Human agency, being obviously unable to reproduce, even on the smallest scale, such conditions as attended the deposition and slow cooling of the earth’s crust, may not hope to rival the products of the planet’s prime. So the fair owners of the earth-born gems may possess their souls in peace, free from any fear of the chemists’ crucibles; and the Kimberley Diamond Companies are not likely to suffer panics from the results of scientific researches, and probably will continue to pay their handsome dividends for time indefinite.

But curiously enough, a discovery of the latest years of our century has revealed the existence of diamonds in a region not mapped by the most advanced of geographers—a region which indeed cannot be defined by degrees of latitude and longitude. In the recesses of an unquestionable meteorite—one of those celestial lumps of iron of which mention has been made in the earlier pages of this volume—real diamonds have been found. These quite resembled the products of M. Moissan’s experiments, being extremely small, but including clear and perfectly shaped crystals, associated with black ones, and also with much graphite in more or less definite forms. So very limited, however, could be the quantity of diamonds obtainable from this hitherto unsuspected source, that even if they rivalled in quality the finest stones from the South African mines, it might be difficult to form a “Company” for their _exploitation_. Still, there is the possibility of some one falling in with a little meteorite containing some mature full-sized carbon crystals, and such a one might be considered equally fortunate with the finder of the famous Australian nugget “Welcome” (£25,000). The association of diamonds with the ferruginous matter of the “blue ground” in the Kimberley pipes, their crystallization out of iron in M. Moissan’s experiments, and their presence in iron meteorites, would seem to point to special relations between the two elements, iron and carbon. Some of these relations are exemplified in another way by the profound modification effected in the physical properties of iron, by its combination with a very small quantity of carbon, as in some kinds of steel; or again, by the differences between white cast iron and grey cast iron, as determined by the condition of the carbon in each.

NEW METALS.

The chemistry of the nineteenth century can boast of a series of discoveries more brilliant and more numerous than ever belonged to any other science within a like period. And the advantage to the world must have been great, for chemistry more directly than any other branch of knowledge ministers to the useful arts and promotes the comfort and well-being of society. The science itself, as it now exists, is almost the creation of the present age. But its recent developments cannot be here discussed; nor, of the immense number of new products with which it has enriched the world, can more than a very few be brought under the reader’s notice in the remaining pages of the present work. Among the most striking of the remarkable series of discoveries by which Sir Humphrey Davy penetrated the mysteries of matter was the isolation of the alkali metals—a circumstance which marks an important era in the history of chemistry. That the alkalies were oxides of unknown metals had indeed been previously surmised by chemists, from the fact of their behaving like metallic oxides in neutralizing and combining with acids to form the class of compounds called _salts_. All attempts to decompose these alkalies had proved fruitless until Davy separated the metal _potassium_ from potash, in 1807. When, however, this alkali had once been proved a compound, more correct ideas were introduced into chemical science; the nature of other alkalies and earths was explained in like manner, and new and powerful re-agents were placed in the hands of the chemist.

Davy first obtained potassium by exposing to the action of the voltaic current a fragment of potash which had become moist on the surface by exposure to the air. The battery was formed of the then unprecedented combination of two hundred pairs of 6–inch plates on Wollaston’s plan, which was constructed for the Royal Institution of London. The heat produced by the passage of the current fused the potash, and globules of metallic potassium were separated at the negative wire. This method yielded the metal in very small quantities only, and at a great cost. Gay Lussac and Thenard soon afterwards found that potassium could be obtained more cheaply and in greater abundance when fused potash was made to flow over iron-turnings heated to whiteness in a gun-barrel, and the hydrogen and potassium vapour were passed into a cooled receiver, in which the latter body was condensed. The metal is now obtained by heating potassium carbonate with charcoal. For this purpose it suffices to heat crude tartar in a covered vessel from which air is excluded. The tartar is first calcined in a crucible until all combustible vapour has been driven off. The charred mass, which now consists of potassium carbonate mixed with finely-divided carbon, is then broken into lumps and quickly introduced into a wrought-iron retort, which is heated in a furnace to nearly a white heat. A receiver in the form of a flat iron box, 12 in. long, 5 in. wide, and ¼ in. deep, is adapted to the neck of the retort, and is kept cooled by the application of a wet cloth on the outside. The potassium thus obtained is not pure, and it must be distilled in an iron retort, as otherwise a powerfully detonating compound is apt to be formed by a portion of the metal combining with carbonic oxide.

Immediately after his discovery of potassium Davy obtained sodium in the same manner, and Gay Lussac and Thenard also procured it by the same process they used for the sister metal. Sodium is now extracted on the manufacturing scale for use as an agent in the reduction of two other metals, of which we shall have to speak. A mixture of dried sodium carbonate, powdered charcoal, and chalk is heated in wrought-iron cylinders, about 4 ft. long, 5 in. internal diameter, and ½ in. thick. The chalk takes no part in the chemical action, but is added in order to give the sodium carbonate when it fuses a pasty consistence, and thus prevent the separation of the charcoal. A number of these iron cylinders are set in a reverberatory furnace; but they are coated with fire-clay and enclosed in earthenware tubes, to prevent their destruction by the intense heat. To one end of each cylinder a receiver is adapted, of the form and dimensions already described for potassium. The other extremity is closed by an iron plug, luted with fire-clay. When the charge in a cylinder is exhausted, a fresh one is introduced by removing the plug, taking out the residue, and inserting a new supply of the mixture made up in a canvas bag. The operation is therefore continuous, and the metal obtained is nearly pure, as sodium does not exhibit the same tendency as potassium to form compounds with carbonic oxide.

Potassium and sodium are extremely soft metals; they are lighter than water, upon which they float, at the same time rapidly decomposing that compound by displacing half the hydrogen, which is set on fire by the heat. The instant a piece of potassium touches the surface of water, a violet flame bursts forth; but with sodium no flame appears unless the metal is dropped on warm water, or prevented from swimming about. Since these metals are thus capable of displacing hydrogen from its combination with oxygen at ordinary temperatures, it follows that they must have a powerful affinity for oxygen; and, indeed, they can only be preserved in rock oil, for they rapidly combine with the oxygen of the air. The great attraction of these metals for oxygen, and for chlorine and other similar bodies, induces the chemist to employ them for separating such bodies from their combination with other metals. Sodium is generally employed for this purpose, as being far cheaper than potassium.

Among the sixty-nine elementary or undecomposable substances which, variously combined, constitute the whole material of our planet, so far as we are acquainted with it, no fewer than fifty-six are metals. Of these fifty-six metals very few are found in a free or uncombined state, like the gold described in the last article. On the contrary, the whole of the metallic elements of the globe, with insignificant exceptions, exist in nature in a state of combination with one or more of the other thirteen non-metallic substances. In this condition they form the stony masses which are termed the ores of the more common metals, and they constitute also the earths, the metallic bases of which were, until recent times, unsuspected and unknown. Davy followed up his discovery of the metals of potash and soda by experimental demonstrations that the earths _alumina_, _magnesia_, and others, were really oxides of metals; and when the nature of these substances had once been established, chemists soon devised means for readily obtaining their metallic bases in an isolated form. The new metals which have been thus isolated all deserve the attention of the chemist; and the general reader will probably also regard with interest the processes by which two of these new metals, for which practical applications have been found, are extracted, and the properties which have caused them to be produced on the commercial scale. These are _aluminium_, the metallic base of common clay; and _magnesium_, the metallic base of common magnesia, and Epsom salts, and a constituent of dolomite, or magnesian limestone.

Aluminium was first isolated by Œrsted, in 1827, by decomposing its chloride by means of potassium. The chlorine leaves the aluminium to combine with the potassium, and thus the former is set free. Wöhler effected some improvements in Œrsted’s process, and he first obtained the metal in malleable globules. It is, however, to Deville that we are indebted for the invention, in 1854, of a process which admitted of application on a manufacturing scale. He obtains chloride of aluminium by mixing alumina (the oxide of the metal) with powdered charcoal made into a paste with oil, and heating the mixture in a tubular earthenware retort, like those sometimes used in the manufacture of coal-gas, while a current of dry chlorine is made to pass through the vessel. The charcoal combines with the oxygen, forming carbonic oxide, a permanent invisible gas; and the aluminium unites with the chlorine, giving rise to aluminium chloride, which, being volatile, sublimes into a chamber lined with glazed tiles, in which it condenses as a yellow translucent mass. The metal is reduced from the chloride in the following manner: A tube of hard glass, about an inch and a half in diameter, is placed over a furnace, or chaffing-dish, as shown in Fig. 336, where D C is the tube, and G G an iron pan for containing the red-hot charcoal. Into the part of this tube marked E, about half a pound of dry aluminium chloride has previously been introduced, and is kept in its place by plugs of asbestos. A current of dry hydrogen gas, perfectly free from air, is passed through the tube; the gas being generated in the vessel, A, and in B passed over some substance which removes from it all moisture. The aluminium chloride is then gently heated by placing red-hot charcoal beneath it, so that any hydrochloric acid it may contain may be expelled. A long narrow porcelain tray, or “boat,” containing pieces of sodium, F, is then introduced into the tube; and, the current of hydrogen being still maintained, heat is applied to the part of the tube containing the sodium, and the aluminium chloride is made to distil over by a regulated heat. As it passes over the sodium, it is reduced with a vivid glow. The aluminium is set free, and collects in the tray with the double chloride of sodium and aluminium which is produced by the reaction. The tray is removed and more strongly heated in a porcelain tube through which a current of hydrogen is passing, and the metal is thus obtained in globules.

Messrs. Bell, of Newcastle, undertook the manufacture of aluminium by a system founded on this process. The first step is the preparation of pure alumina, which may be obtained by igniting ammonia alum, or by precipitating from a solution of alum free from iron, or from sodium aluminate made from the mineral called _bauxite_. The precipitate of hydrated alumina, mixed with charcoal and common salt, is made into balls and dried. These balls, which are about as large as an orange, are placed in upright earthenware retorts, which are heated to redness, while a current of dry chlorine is passed through them. The volatile double chloride of aluminium and sodium distils over, and is condensed in chambers lined with earthenware. This substance is mixed with powdered fluor-spar, or with _cryolite_ (itself a compound of aluminium), which serves as a flux; and small pieces of sodium are interspersed throughout the mixture. The proportions are ten parts of the double chloride, five of fluor-spar, and two of sodium. This mixture is thrown upon the hearth of a reverberatory furnace, and the doors are shut to exclude air. A very intense action occurs: the chlorine, quitting the aluminium, seizes on the sodium, and their combination is attended by an enormous increase of temperature. The fused aluminium is run off from the furnace together with the slags which are produced by the operation. In this way, with a furnace having a hearth 16 ft. square, about 16 lbs. of aluminium can be obtained in one operation.

Rose, the eminent German metallurgist, prefers to obtain aluminium from cryolite, which is a compound of sodium, aluminium, and fluorine, found in large quantities in Greenland. It is powdered and mixed with common salt, and with the mixture a certain quantity of sodium cut into small pieces is uniformly mingled. The whole is thrown into a heated crucible, previously lined with a fused mixture of cryolite and salt, and more of the same mixture is poured upon the contents of the crucible, which is then covered and exposed to a red heat for two hours. The aluminium generally collects into buttons, which may be easily melted together by heating them in a crucible with common salt.

It will be obvious, from the preceding account of the processes of extracting aluminium, that the cost of the metal must depend upon that of sodium; and the same remark will apply to the case of magnesium. It is interesting to observe how the price of the alkaline metals has decreased as improved processes have been devised, and as the scale of production has increased with the commercial demand for the article. Prepared by Gay Lussac and Thenard’s process, these metals were produced in but small quantities, and were sold at £5 per oz. When the mode of reducing them by charcoal came into operation, the price fell to 30_s._ per oz.; and the researches of Deville so far improved the processes, that in 1854 sodium could be procured for 5_s._ per oz. Mr. Gerhard, of Battersea, subsequently manufactured sodium, so that it can now be retailed at less than 1_s._ per oz. The price of aluminium before Deville’s investigations was about 24_s._ per oz., but now the metal can be purchased at about one-eighth of that cost. [1875.]