1911 Encyclopædia Britannica/Nitrogen

NITROGEN [symbol N., atomic weight 14·01, O＝16]. A non-metallic chemical element, first isolated in 1772 by D. Rutherford, who showed that on removing oxygen from air a gas remained, which was incapable of supporting combustion or respiration. Nitrogen forms approximately 79% by volume (or 77% by weight) of the atmosphere; actual values are: % by volume—79·07 (Regnault), 79·20 (Dumas); % by weight—76·87 (Regnault), 77·00 (Dumas), 77·002 (Léwy), 76·900 (Stas), 77·010 (Marignac). No absolutely accurate determinations appear to have been made recently. Free nitrogen is also found in some natural waters and has been recognized in certain nebulae. In the combined state nitrogen is fairly widely distributed, being found in nitre, Chile saltpetre, ammonium salts and in various animal and vegetable tissues and liquids. It is invariably present in soils, where compounds are formed by nitrifying bacteria.

Nitrogen may be obtained from the atmosphere by the removal of the oxygen with which it is there mixed. This may be effected by burning phosphorus in a confined volume of air, by the action of an alkaline solution of pyrogallol on air, by passing air over heated copper, or by the action of copper on air in the presence of ammoniacal solutions.

It is also prepared by heating ammonium nitrite (or a mixture of sodium nitrite and ammonium chloride): NH₄NO₂＝2H₂O+N₂; by heating a mixture of ammonium nitrate and chloride (the chlorine which is simultaneously produced being absorbed by milk of lime or by a solution of sodium hydroxide): 4NH₄NO₃+2NH₄Cl＝5N₂ +Cl₂+12H₂O; by heating ammonium dichromate (or a mixture of ammonium chloride and potassium dichromate): (NH₄), Cr₂O₇＝Cr₂O₃+4H₂O+N₂; by passing chlorine into a concentrated solution of ammonia (which should be present in considerable excess): 8NH₃+3Cl₂＝6NH₄Cl+N₂; by the action of hypochlorites or hypobromites on ammonia: 3NaOBr+2NH₃＝3NaBr+3H₂O+N₂; and by the action of manganese dioxide on ammonium nitrate at 180-200° C. It is also formed by the reduction of nitric and nitrous oxides with hydrogen in the presence of platinized asbestos at a red heat (G. v. Knorre and K. Arndt, Ber., 1899, 32, p. 2136); by the oxidation of hydroxylamine (ibid., 1900, 33, p. 30); and by the electrolysis of hydrazine and its salts (E. Ch. Szarvasy, Jour. Chem. Soc., 1900, 77, p. 603)

The chief importance of nitrogenous compounds depends upon their assimilation by living plants, which, in their development, absorb these compounds from the soil, wherein they are formed mainly by the action of nitrifying bacteria. Since these compounds are essential to plant life, it becomes necessary to replace the amount abstracted from the soil, and hence a demand for nitrogenous manures was created. This was met in a very large measure by deposits of natural nitre and the products of artificial nitrières, whilst additional supplies are available in the ammoniacal liquors of the gas-manufacturer, &c. The possible failure of the nitre deposits led to attempts to convert atmospheric nitrogen into manures by processes permitting economic success. Combination can be made in five directions, viz. to form (1) oxides and nitric acids, (2) ammonia, (3) readily decomposable nitrides, (4) cyanides, (5) cyanamides. The first three will be treated here; for the others see Prussic Acid and Cyanamide.

The combination of nitrogen with oxygen was first effected by Cavendish in 1785, who employed a spark discharge. The process was developed by Madame Lefebre in 1859; by Meissner in 1863, who found that moist gases gave a better result; and by Prim in 1882, who sparked the gases under pressure; it was also used by Lord Rayleigh in his isolation of argon (q.v.). It was not, however, a commercial success, and the same result attended Siemens and Halske’s application of the silent discharge. More effective was the electric arc. In 1892 Sir W. Crookes showed that the arc brought about combination; and in 1897 Lord Rayleigh went into the process more fully. But the first careful working-out of the conditions was made in 1900 by A. McDougall and F. Howles, who, employing a high tension alternating arc, showed that the effectiveness depended upon the temperature. The commercial manufacture of nitric acid was attempted by C. S. Bradley and D. R. Lovejoy at Niagara Falls, who passed atmospheric air, or air enriched with oxygen, about a high tension arc made as long as possible; but the company (the Atmospheric Products Company) was a failure. Better results have attended the process of K. Birkeland and S. Eyde, which is being worked on a large scale at Notodden, Norway. The arc is produced by leading a current of about 5000 volts equatorially between the poles of an electromagnet; this produces what is practically a disk of flame, 61/2 ft. in diameter and having a temperature of about 3000°. The disk really consists of a series of successive arcs which increase in size until they burst. The first product of the reaction is nitric oxide, which on cooling with the residual gases produces nitrogen peroxide. The cooled gases are then led into towers where they meet a stream of water coming in the contrary direction. Nitric acid (up to 59%) is formed in the first tower, and weaker acids in the successive ones; the last tower contains milk of lime which combines with the gases to form calcium nitrite and nitrate (this product, being unsuitable as a manure, is decomposed with the acid, and the evolved gases sent back). It was found advantageous not to work for acid but for a basic calcium nitrate (normal calcium nitrate being very deliquescent); for this purpose the acid is treated with the requisite amount of milk of lime. In the process of the Badische Anilin- und Soda-Fabrik, the arc, which is said to be 30 to 50 ft. long or more, is formed in a long tube, and the gases are sent round the arc by obliquely injecting them. A 30% acid is said to be formed. I. Moscicki and J. von Kowalski have patented a process wherein the arc is formed at two vertical concentric copper electrodes and rotated by an electromagnet; it is worked at Vevey, Switzerland. The Rankin process, of which very little is known, produces the arc with much lower current.

The conversion of nitrogen into ammonia by electricity has received much attention, but the commercial aspect appears to have been first worked out by de Hemptinne in 1900, who used both the spark and silent discharge on mixtures of hydrogen and nitrogen, and found that the pressure and temperature must be kept low and the spark gap narrow. J. Schlutius in 1903 employed Dowson gas as a source of hydrogen, and induced combination by means of platinum and the silent discharge. Several non-electrical processes have been devised. In 1862 Fleck passed a mixture of steam, nitrogen and carbon monoxide over red-hot lime, whilst in 1904 Woltereck induced combination by passing steam and air over red-hot iron oxide (peat is used in practice). In de Lambilly's process air and steam is led over white-hot coke. and carbon dioxide or monoxide removed from the escaping gases according as ammonium formate or carbonate is wanted. The residual gas is then passed through a tube containing porous materials, such as wood- or bone-charcoal, platinized pumice or spongy platinum, then mixed with steam and again forced through the tube. The reactions are represented as

(1) N₂+3H₂+2CO+2H₂O＝2H⋅CO₂NH₄  (Ammonium formate).

(2) N₂+3H₂+2CO₂+2H₂O＝2HO⋅CO₂NH₄ (Ammonium carbonate).

The best temperature for the first reaction is between 80° C. and 130° C. and for the second between 40° C. and 60° C. In another process, which originated with C. Kaiser (Abst. J.C.S., 1907, ii. p. 862), calcium is heated in a current of hydrogen, and nitrogen passed over the hydride so formed; this gives ammonia and calcium nitride, the latter of which gives up its nitrogen as ammonia and reforms the hydride when heated in a current of hydrogen.

The fixation of nitrogen as a nitride has not been attended with commercial success. H. Mehner patented heating the oxides of silicon, boron or magnesium with coal or coke in an electric furnace, and then passing in nitrogen, which forms, with the metal liberated by the action of the carbon, a readily decomposable nitride.

For an extended bibliography see Bulletin No. 63 of the Bureau of Soils, U.S. Department of Agriculture (Washington, 1910).

Nitrogen is a colourless, tasteless and odourless gas, which is only very slightly soluble in water. It is slightly lighter than air. Lord Rayleigh in 1894 found that the density of atmospheric nitrogen was about 1/2% higher than that of chemically prepared nitrogen, a discovery which led to the isolation of the rare gases of the atmosphere (see Argon). The values obtained are shown below.,

Atmospheric Nitrogen.	Chemical Nitrogen.
0·97209	0·96727	Lord Rayleigh, Chem. News, 1897, 76, p. 315.
0·9720	0·9671	A. Leduc, Comptes rendus, 1896, 123, p. 805.

Nitrogen is a very inert gas: it will neither burn nor support the combustion of ordinary combustibles. It combines directly with lithium, calcium and magnesium when heated, whilst nitrides of the rare earth metals are also produced when their oxides are mixed with magnesium and heated in a current of nitrogen (C. Matignon, Comptes rendus, 1900, 131, p. 837). Nitrogen has been liquefied, the critical temperature being −149° C. and the critical pressure 27·54 atmospheres. The liquefied gas boils at −195·5° C., and its specific gravity at its boiling point is 0·8103 (E. C. C. Baly and F. G. Donnan, Jour. Chem. Soc., 1902, 81, p. 912).

Compounds.

Nitrogen combines with hydrogen to form ammonia, NH₃, hydrazine, N₂H₄, and azoimide, N₃H (qq.v.); the other known hydrides, N₄H₄ and N₅H₅, are salts of azoimide, viz. NH₄⋅N₃ and N₂H₄N₃H.

Nitrogen trichloride, NCl₃, discovered by P. L. Dulong in 1811 (Schweigg. Journ., 1811, 8, p. 302), and obtained by the action of chlorine or sodium hypochlorite on ammonium chloride, or by the electrolysis of ammonium chloride solution, is a very volatile yellow oil. It possesses an extremely pungent smell, and its vapour is extremely irritating to the eyes. It is a most dangerous explosive (see D. L. Chapman and L. Vodden, Jour. Chem. Soc., 1909, 95, p. 138). Chlorine azide, Cl⋅N₃, was discovered by F. Raschig in 1908 (see Azoamide); the corresponding iodine compound had been obtained in 1900 by A. Hantzsch (Ber., 33, p. 522). For the so-called nitrogen iodide see Ammonia.

Nitrogen forms five oxides, viz. nitrous oxide, NH₂O, nitric oxide, NO, nitrogen trioxide, N₂O₃, nitrogen peroxide, NO₂, and nitrogen pentoxide, N₂O₂, whilst three oxyacids of nitrogen are known: hyponitrous acid, H₂N₂O₂, nitrous acid, HNO₂, and nitric acid. HNO₃ (q.v.). The first four oxides are gases, the fifth is a solid. Nitrous oxide, N₂O, isolated in 1776 by J. Priestley, who obtained it by reducing nitrogen peroxide with iron, may be prepared by heating ammonium nitrate at 170-260° C., or by reducing a mixture of nitric and sulphuric acid with zinc. It is a colourless gas, which is practically odourless, but possesses a sweetish taste. It is somewhat soluble in water. When liquefied it boils at −89·8° C., and by further cooling may be solidified, the solid melting at −102·3° (W. Ramsay, Chem. News, 1893, 67, p. 140). It does not burn, but Supports the combustion of heated substances almost as well as oxygen. It is used as an anaesthetic, principally in dentistry, producing when inhaled a condition of hysterical excitement often accompanied by loud laughter, whence it is sometimes called “laughing gas.”

Nitric oxide, NO, first obtained by Van Helmont, is usually prepared by the action of dilute nitric acid (sp. gr. 1·2) on copper. This method does not give a pure gas, varying amounts of nitrous oxide and nitrogen being present (see Nitric Acid). In a purer condition it may be obtained by the action of sulphuric acid on a mixture of potassium nitrate and ferrous sulphate, or of hydrochloric acid on a mixture of potassium nitrate and ferric chloride. It is also formed by the action of concentrated sulphuric acid on sodium nitrite in the presence of mercury. It is a colourless gas which is only sparingly soluble in water. It may be liquefied, its critical temperature being −93·5°, and the liquid boils at −153·6° C. It is not a supporter of combustion, unless the substance introduced is at a sufficiently high temperature to decompose the gas, when combustion will continue at the expense of the liberated oxygen. If the gas be mixed with the vapour of carbon disulphide, the mixture burns with a vivid lavender-coloured flame. Nitric oxide is soluble in solutions of ferrous salts, a dark brown solution being formed, which is readily decomposed by heat, with evolution of nitric oxide. It combines with oxygen to form nitrogen peroxide. Nascent hydrogen reduces it to hydroxylamine (q.v.), whilst solutions of hypochlorites oxidize it to nitric acid. In some instances it reacts as a reducing agent, e.g. silver oxide is reduced to metallic silver at 170° C., lead dioxide to the monoxide and manganese dioxide to sesquioxide.

Nitrogen trioxide, N₂O₃, was first mentioned by J. R. Glauber in 1648 as a product of the reaction between nitric acid and arsenious oxide. Sir W. Ramsay (Jour. Chem. Soc., 1890, 5, p. 590), by distilling arsenious oxide with nitric acid and cooling the distillate, obtained a green liquid which consisted of nitrogen trioxide and peroxide in varying proportions, and concluded that the trioxide could not be obtained pure. He then tried the direct combination of nitric oxide with liquid nitrogen peroxide. A dark blue liquid is produced, and the first portions of gas boiling off from the mixture correspond fairly closely in composition with nitrogen trioxide. H. B. Baker (Jour. Chem. Soc., 1907, 91, p. 1862) obtained nitrogen trioxide in the gaseous form by volatilizing the liquid under special conditions. L. Francesconi and N. Sciacca (Gazz., 1904, 34 (i.), p. 447) have shown that liquid nitric oxide and oxygen, or gaseous nitric oxide and liquid oxygen, mixed in all proportions and yielded nitrogen trioxide, whilst gaseous nitric oxide mixed with excess of oxygen always gave the trioxide if the mixture was kept below −110° C. They also state that nitrogen trioxide is stable at ordinary pressure up to −21° C. N. M. v. Wittorf (Zeit. anorg. Chem., 1904, 41, p. 85) obtained blue crystals of the trioxide (melting at −103° C.) on saturating liquid nitrogen peroxide with nitric oxide and cooling the mixture. The liquid prepared by Baker is green in colour, and has a specific gravity 1·11 at ordinary temperature, but below −2° C. becomes of a deep indigo blue colour. It forms a mass of deep blue crystals at the temperature of liquid air. It is exceedingly soluble in concentrated sulphuric acid.

Nitrogen peroxide, NO₂ or N₂O₄, may be obtained by mixing oxygen with nitric oxide and passing the red gas so obtained through a freezing mixture. The production of this red gas when air is mixed with nitric oxide was mentioned by R. Boyle in 1671. Nitrogen peroxide is also prepared by heating lead nitrate and passing the products of decomposition through a tube surrounded by a freezing mixture, when the gas liquefies. At low temperatures it is a colourless crystalline solid which melts at −10·14° C. (W. Ramsay, Chem. News, 1900, 61, p. 91). As the temperature increases the liquid becomes yellowish, the colour deepening with rise of temperature until at +15° C. it has a deep orange tint. The liquid boils at about 22° C. This change of colour is accompanied by a change in the vapour density, and is explained by the fact that nitrogen peroxide consists of a mixture of a colourless compound N₂O₄, and a red-brown gas NO₂, the latter increasing in amount at the expense of the former as the temperature is raised (G. Salet, Comptes rendus, 1868, 67, p. 488; see also E. and L. Natanson, Wied. Ann., 1885, 24, p. 454; 1886, 27, p. 606). M. Berthelot and J. Ogier (Bull. Soc. Chim., 1882 [2], 37, p. 434; 38, p. 60) have also shown that the specific heat of the gas decreases with increase of temperature until it reaches a minimum at about 198-253° C. Cryoscopic determinations of the molecular weight of nitrogen peroxide dissolved in glacial acetic acid show that it corresponds to the molecular formula N₂O₄ at low temperatures (W. Ramsay, Jour. Chem. Soc., 1888, 53, p. 621). Nitrogen peroxide is the most stable oxide of nitrogen. It is decomposed by water, giving at 0° C. a mixture of nitric and nitrous acids: 2NO₂+H₂O＝HNO₃+HNO₂. It combines with sulphuric acid to form nitro-sulphonic acid, SO₂(OH)(NO₂). It does not support the combustion of a taper, but burning phosphorus and red-hot carbon will continue to burn in the gas. It converts many metallic oxides into mixtures of nitrates and nitrites, and attacks many metals, forming nitrates and being itself reduced to nitric oxide. It is an energetic oxidizing agent.

Nitrogen pentoxide, N₂O₅, was first obtained in 1849 by H. Sainte-Claire-Deville (Ann. Chim. Phys., 1850 [3], 28, p. 241) by the action of dry chlorine on silver nitrate: 4AgNO₃+2Cl₂＝4AgCl+2N₂O₅ +O₂. It may also be obtained by distilling nitric acid over phosphorus pentoxide. It crystallizes in large prisms which melt at 29-30° C. to a yellowish liquid, which boils at 45-50° C. with rapid decomposition. It is very unstable, decomposing slowly even at ordinary temperatures. It dissolves in water, forming nitric acid.

Hyponitrous acid, H₂N₂O₂, was first obtained in the form of its salts by E. Divers in 1871 (Chem. News, 23, p. 206) by reducing a solution of potassium nitrite with sodium amalgam, and subsequent precipitation as silver salt. Hyponitrites also result when hydroxyamido-sulphonates, e.g. HO⋅NH⋅SO₃Na, are hydrolysed by caustic alkalis (E. Divers and T. Haga, Jour. Chem. Soc., 1889, 55, p. 760), or when benzsulphohydroxamic acid, C₆H₅SO₂⋅NH⋅OH, is treated in the same manner (O. Piloty, Ber., 1896, 29, p. 1560). They may also be prepared by the action of mercuric or cupric oxides on alkaline solutions of hydroxylamine (A. Hantzsch, Ann., 1896, 292, p. 317); by the action of hydroxylamine sulphate on alkaline nitrites in the presence of lime or calcium carbonate, the mixture being rapidly heated to 60° C.; or by the hydrolysis of dimethyl nitroso-oxyurea, (CH₃)₂N⋅CO⋅N(NO)⋅OH (A. Hantzsch, Ber., 1897, 30, p. 2356). The free acid, which crystallizes in brilliant scales, is best prepared by decomposing the silver salt with an ethereal solution of hydrochloric acid. It is very explosive, dissolves readily in water and behaves as a dibasic acid. It does not liberate iodine from potassium iodide, neither does it decolorize iodine solution. Bromine oxidizes it to nitric acid, but the reaction is not quantitative. In acid solution, potassium permanganate oxidizes it to nitric acid, but in alkaline solution only to nitrous acid. It decomposes slowly on standing, yielding water and nitrous oxide. The silver salt is a bright yellow solid, soluble in dilute sulphuric and nitric acids, and may be crystallized from concentrated solutions of ammonia. It slowly decomposes on exposure or on heating. The calcium salt, CaN₂O₂⋅4H₂O, formed by the action of calcium chloride on the silver salt in the presence of a small quantity of nitric acid, is a lustrous crystalline powder, almost insoluble in water but readily soluble in dilute acids.” It is decomposed by sulphuric acid, with evolution of nitrous oxide.

Nitrous acid, HNO₂, is found to some extent in the form of its salts in the atmosphere and in rain water. The pure acid has not yet been obtained, since in the presence of water it decomposes with formation of nitric acid and liberation of nitric oxide: 3HNO₂＝HNO₃+2NO+H₂O. Its salts may be obtained in some cases by heating the corresponding nitrates, but the method does not give good results. Sodium nitrite, the most commonly used salt of the acid, is generally obtained by heating the nitrate with metallic lead; by heating sodium nitrate with sulphur and sodium hydroxide, the product then being fractionally crystallized (Read, Holliday & Sons): 3NaNO₃+S+2NaOH＝Na₂SO₄+3NaNO₂+H₂O; by oxidizing atmospheric nitrogen in an electric arc, keeping the gases above 300° C., until absorption in alkaline hydroxide solution is effected (German Pat. 188188); or by passing air, or a mixture of oxygen and ammonia, over heated metallic oxides (ibid., 168272). The salts of the acid are colourless or faintly yellow. In aqueous solution the free acid acts as an oxidizing agent, bleaching indigo and liberating iodine from potassium iodide, or it may act as a reducing agent since it readily tends to pass into nitric acid: consequently it discharges the colour of acid solutions of permanganate's and chromates. The acid finds considerable use in organic chemistry, being employed to discriminate between the different types of alcohols and of amines, and also in the production of diazo, azo and diazo-amino compounds. It may be recognized by the blue colour it gives with diphenylamine sulphate and by its reaction with potassium iodide-starch paper.

Nitrosyl chloride, NOCl, is obtained by the direct union of nitric oxide with chlorine; or by distilling a mixture of concentrated nitric and hydrochloric acids, passing the resulting gases into concentrated sulphuric acid and heating the so-formed nitrosyl hydrogen sulphate with dry salt: HNO₃+3HCl＝NOCl+C1₂ +H₂O; NOCl + H₂SO₄＝HCl + NO⋅SO₄H; NO⋅SO₄H + NaCl＝NOCl+NaHSO₄ (W. A. Tilden, Jour. Chem. Soc., 1860, p. 630). It is also prepared by the action of phosphorus pentachloride on potassium nitrite or on nitrogen peroxide. It is an orange-coloured gas which may be readily liquefied and by further cooling may be solidified. The liquid boils at −5° C. and the solid melts at −65° C. It forms double compounds with many metallic chlorides, and finds considerable application as a means of separating various members of the terpene group of compounds. It is readily decomposed by water and alkaline hydroxides, yielding a mixture of nitrite and chloride. On treatment with silver fluoride it yields nitrosyl fluoride, NOF (O. Ruff, Zeit. anorg. Chem., 1905, 47, p. 190). Nitroxyl fluoride, NO₂F, is formed by the action of fluorine on nitric oxide at the temperature of liquid oxygen (H. Moissan and P. Lebeau, Comptes rendus, 1905, 140, pp. 1573, 1621). It is a gas at ordinary temperature; when liquefied it boils at −63·5° C. and on solidification melts at −139° C. Water decomposes it into nitric and hydrofluoric acids. Nitramide, NH₂NO₂, is obtained by the action of sulphuric and nitric acids on potassium imidosulphonate, or by the action of ice-cold sulphuric acid on potassium nitro-carbamate (J. Thiele and A. Lachmann, Ann., 1895, 288, p. 297): NO₂⋅NK⋅CO₂K+H₂SO₄＝NH₂NO₂+K₂SO₄+CO₂. It crystallizes in prisms or leaflets which melt at 72-75° C. and are readily soluble in water and in all organic solvents except ligroin. It is somewhat volatile at ordinary temperature, and its aqueous solution possesses a strongly acid reaction. It is very unstable, decomposing into nitrous oxide and water when mixed with copper oxide, lead chromate or even powdered glass. On reduction it gives a strongly reducing substance, probably hydrazine. According to A. Hantzsch (Ann., 1896, 292, pp. 340 et seq.) hyponitrous acid and nitramide are to be regarded as stereo-isomers, being the anti- and syn- forms of the same compound. Thiele, however, regards nitramide as imidonitric acid, HN:NO(OH).

Nitrogen sulphide, N₄S₄, first obtained by W. Gregory (Jour. pharm., 1835, 21, p. 315) by the action of ammonia on sulphur chloride, has been investigated by O. Ruff and E. Geisel (Ber., 1904, 37, 1573; 1905, 38, p. 2659). who also obtained it by dissolving sulphur in liquid ammonia. It is a reddish-yellow crystalline solid, insoluble in water and melting at 178° C. It explodes readily when melted or subjected to shock. Dry hydrochloric acid gives ammonia but no nitrogen; with ammonia it gives N:SNH₂ and S:S(NH₂)₂; and with secondary amines it forms thiodiamines, S(NR₂)₂, nitrogen and ammonia being liberated. When heated with CS₂ to 100° C. under pressure, it forms liquid nitrogen sulphide, N₂S₅, a mobile red liquid which solidifies to an iodine-like mass of crystals which melt at 10-11° C. Water, alkalis and acids decompose it into sulphur and ammonia (W. Muthmann, Zeit. anorg. Chem., 1897, 13, p. 200).

For sulphonic acids containing nitrogen see Ammonia.

Numerous determinations of the atomic weight of nitrogen have been made by different observers, the values obtained varying somewhat according to the methods used. These methods have been purely chemical (either gravimetric or volumetric), physical (determinations of the density of nitrogen, nitric oxide, &c.), or physicochemical. P. A. Guye has given a critical discussion of the relative accuracy of the gravimetric and physico-chemical methods, and favours the latter, giving for the atomic weight a value less than 14·01. The more important papers dealing with the subject are: J. Stas, (Œuvres complètes, i. pp. 342 et seq.; Lord Rayleigh, Proc. Roy. Soc. (1894), 55, p. 340; (1904) 73, p. 153; G. Dean, Jour. Chem. Soc. (1901), 79, p. 147; R. W. Gray, Jour. Chem. Soc. (1906), 88, p. 1174; A. Scott, Proc. Chem. Soc. (1905), 21, p. 309; P. A. Guye, Chem. News (1905), 92, pp. 261 et seq.; (1906) 93, p. 13 et seq.; D. Berthelot, Comptes rendus (1907), 144, p. 269.