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MAGNESIUM [symbol Mg, atomic weight 2...

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Originally appearing in Volume V17, Page 321 of the 1911 Encyclopedia Britannica.
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MAGNESIUM [See also:symbol Mg, atomic See also:weight 24.32 (0 = 16)] , a metallic chemical See also:element. The sulphate or " See also:Epsom salts " (q.v.) was isolated in 1695 by N. See also:Grew, while in 1707 M. B. Valentin prepared See also:magnesia See also:alba from the See also:mother liquors obtained in the manufacture of See also:nitre. Magnesia was See also:con-founded with See also:lime until 1755, when J. See also:Black showed that the two substances were entirely different; and in 1808 See also:Davy pointed out that it was the See also:oxide of a See also:metal, which, however, he was not able to isolate. Magnesium is found widely distributed in nature, chiefly in the forms of silicate, carbonate and chloride, and occurring in the minerals See also:olivine, See also:hornblende, See also:talc, See also:asbestos, See also:meerschaum, See also:augite, See also:dolomite, See also:magnesite, carnallite, kieserite and kainite. The metal was prepared (in a See also:state approximating to purity) by A. A. B. See also:Bussy (Jour. de pharm.

1829, 15, p. 30; 1830, 16, p. 142), who fused the an-hydrous chloride with See also:

potassium; H. Sainte Claire Deville's See also:process, which used to be employed commercially, was essentially the same, except that See also:sodium was substituted for potassium (Comptes rendus, 1857, 44, p. 394), the product being further purified by redistillation. It may also be prepared by See also:heating a mixture of See also:carbon, oxide of See also:iron and magnesite to See also:bright redness; and by heating a mixture of magnesium ferrocyanide and sodium carbonate, the See also:double See also:cyanide formed being then decomposed by heating it with metallic See also:zinc. Electrolytic methods have entirely superseded the older methods. The problem of magnesium reduction is in many respects similar to that of See also:aluminium extraction, but the lightness of the metal as compared, bulk for bulk, with its fused salts, and the readiness with which it See also:burns when exposed to See also:air at high temperatures, render the problem somewhat more difficult. See also:Moissan found that the oxide resisted reduction by carbon in the electric See also:furnace, so that See also:electrolysis of a fusible See also:salt of the metal must be resorted to. See also:Bunsen, in 1852, electrolysed fused magnesium chloride in a See also:porcelain crucible. In later processes, carnallite (a natural double chloride of magnesium and potassium) has commonly, after careful dehydration, been substituted for the single chloride. Graetzel's process, which was at one See also:time employed, consisted in electrolysing the chloride in a metal crucible heated externally, the crucible itself forming the See also:cathode, and the magnesium being de-posited upon its inner See also:surface.

W. Borchers also used an externally heated metal See also:

vessel as the cathode; it is provided with a supporting See also:collar or flange a little below the See also:top, so that the upper See also:part of the vessel is exposed to the cooling See also:influence of the air, in See also:order that a crust of solidified salt may there be formed, and so prevent the creeping of the electrolyte over the top. The carbon anode passes through the See also:cover of a porcelain See also:cylinder, open at the bottom, and provided with a See also:side-See also:tube at the top to remove the See also:chlorine formed during electrolysis. The operation is conducted at a dull red See also:heat (about 76o° C. or 1400° F.), the current See also:density being about 0.64 amperes per sq. in. of cathode surface, and the pressure about 7 volts. The fusing-point of the metal is about 730° C. (1350° F.), and the magnesium is therefore reduced in the See also:form of melted globules which gradually accumulate. At intervals the current is interrupted, the cover removed, and the temperature of the vessel raised considerably above the melting-point of magnesium. The metal is then removed from the walls with the aid of an iron scraper, and the whole See also:mass poured into a See also:sheet-iron See also:tray, where it solidifies. The solidified chloride is then broken up, the shots and fused masses of magnesium are picked out, run together in a See also:plumbago crucible without See also:flux, and poured into a suitable See also:mould. Smaller pieces are thrown into a See also:bath of melted carnallite and pressed together with an iron See also:rod, the bath being then heated until the globules of metal See also:float to the top, when they may be removed in perforated iron ladles, through the holes in which the fused chloride can drain away, but through which the melted magnesium cannot pass by See also:reason of its high surface tension. The globules are then re-melted. F.

Oettel (Zeit. f. Elektrochem., 1895, 2, p. 394) recommends the electrolytic preparation from carnallite; the See also:

mineral should be freed from See also:water and sulphates. Magnesium is a silvery See also:white metal possessing a high lustre. It is malleable and ductile. Sp. gr. 1.75. It preserves its lustre in dry air, but in moist air it becomes tarnished by the formation of a film of oxide. It melts at 632.7° C. (C. T. Hey-See also:cock and F.

H. See also:

Neville), and boils at about 11oo°C. Magnesium and- its salts are diamagnetic. It burns brilliantly when heated in air or See also:oxygen, or even in carbon dioxide, emitting a brilliant white See also:light and leaving a See also:residue of magnesia, MgO. The light is See also:rich in the See also:violet and ultra-violet rays, and consequently is employed in See also:photography. The metal is also used in pyrotechny. It also burns when heated in a current of See also:steam, which it decomposes with the liberation of See also:hydrogen and the formation of magnesia. At high temperatures it acts as a reducing See also:agent, reducing See also:silica to See also:silicon, boric See also:acid to See also:boron, &c. (H. Moissan, Comptes rendus, 1892, 114, p. 392). It combines directly with See also:nitrogen, when heated in the See also:gas, to form the nitride Mg3N2 (see See also:ARGON).

It is rapidly dissolved by dilute acids, with the See also:

evolution of hydrogen and the formation of magnesium salts. It precipitates many metals from solutions of their salts. Magnesium Oxide, magnesia, MgO, occurs native as the mineral periclase, and is formed when magnesium burns in air; it may also be prepared by the See also:gentle ignition of the hydroxide or carbonate. It is a non-volatile and almost infusible white See also:powder, which slowly absorbs moisture and carbon dioxide from air, and is readily soluble in dilute acids. On See also:account of its refractory nature, it is employed in the manufacture of crucibles, furnace linings, &c. It is also used in making See also:hydraulic cements. A crystalline form was obtained by M. Houdard (Abst. J. C. S., 1907, ii. p. 621) by fusing the oxide and sulphide in the electric furnace.

Magnesium hydroxide Mg(OH)2, occurs native as the minerals See also:

brucite and nemalite, and is prepared by precipitating solutions of magnesium salts by means of See also:caustic soda or potash. An artificial brucite was prepared by A. de Schulten (Comptes rendus, 1885, 1o1, p. 72) by boiling magnesium chloride with caustic potash and allowing the See also:solution to cool. Magnesium hydroxide is a white amorphous solid which is only slightly soluble in water; the solubility is, however, greatly in-creased by ammonium salts. It possesses an alkaline reaction and absorbs carbon dioxide. It is employed in the manufacture of cements. When magnesium is heated in See also:fluorine or chlorine or in the vapour of See also:bromine or See also:iodine there is a violent reaction, and the corresponding halide compounds are formed. With the exception of the fluoride, these substances are readily soluble in water and are deliquescent. The fluoride is found native as sellaite, and the bromide and iodide occur in See also:sea water and in many mineral springs. The most important of the halide salts is the chloride which, in the hydrated form, has the See also:formula MgCl2.6H20. It may be prepared by dissolving the metal, its oxide, hydroxide, or carbonate in dilute hydrochloric acid, or by mixing concentrated solutions of magnesium sulphate and See also:common salt, and cooling the mixture rapidly, when the less solublesodium sulphate separates first. It is also formed as a by-product in the manufacture' of potassium chloride from carnallite.

The hydrated salt loses water on heating, and partially decomposes into hydrochloric acid and magnesium oxychlorides. To obtain the anhydrous salt, the double magnesium ammonium chloride, MgClz• NH4Cl•6H2O, is prepared by adding ammonium chloride to a solution of magnesium chloride. The solution is evaporated, and the residue strongly heated, when water and ammonium chloride are expelled, and anhydrous magnesium chloride remains. Magnesium chloride readily forms double salts with the alkaline chlorides. A strong solution of the chloride made into a thick See also:

paste with calcined magnesia sets in a few See also:hours to a hard, See also:stone-like mass, which contains an oxychloride of varying See also:composition. Magnesium oxychloride when heated to redness in a current of air evolves a mixture of hydrochloric acid and chlorine and leaves a residue of magnesia, a reaction which is employed in the See also:Weldon-Pechiney and See also:Mond processes for the manufacture of chlorine. Magnesium Carbonate, MgCO3.—The normal salt is found native as the mineral magnesite, and in See also:combination with See also:calcium carbonate as dolomite, whilst hydromagnesite is a basic carbonate. It is not possible to prepare the normal carbonate by precipitating magnesium salts with sodium carbonate. C. See also:Marignac has prepared it by the See also:action of calcium carbonate on magnesium chloride. A salt MgCO3.3H2O or Mg(CO3H)(OH)•2H20 may be prepared from the carbonate by dissolving it in water charged with carbon dioxide, and then reducing the pressure (W. A.

See also:

Davis, Jour. See also:Soc. Chem. Ind. 1906, 25, p. 788). The carbonate is not easily soluble in dilute acids, but is readily soluble in water containing carbon dioxide. Magnesia alba, a white bulky precipitate obtained by adding sodium carbonate to Epsom salts,is a mixture of Mg(CO3H) (OH) •2H20,Mg(CO3H) (OH) and Mg(OH)2. It is almost insoluble in water, but readily dissolves in ammonium salts. Magnesium See also:Phosphates.—Byadding sodium phosphate to magnesium sulphate and allowing the mixture to stand, hexagonal needles of MgHPO4.7H2O are deposited. The normal phosphate, Mg3P2O8, is found in some guanos, and as the mineral wagnerite. It may be prepared by adding normal sodium phosphate to a magnesium salt and boiling the precipitate with a solution of magnesium sulphate.

It is a white amorphous powder, readily soluble in acids. Magnesium ammonium phosphate, MgNH4PO4.6H2O, is found as the mineral struvite and in some guanos; it occurs also in urinary calculi and is formed in the putrefaction of urine. It is prepared by adding sodium phosphate to magnesium sulphate in the presence of See also:

ammonia and ammonium chloride. When heated to See also:loo° C., 'it loses five molecules of water of See also:crystallization, and at a higher temperature loses the See also:remainder of the water and also ammonia, leaving a residue of magnesium pyrophosphate,. Mg2P2O7. Magnesium Nitrate, Mg(NO3)2.6H2O, is a colourless, deliquescent, crystalline solid obtained by dissolving magnesium or its carbonate in nitric acid, and concentrating the solution. The crystals melt at 9o° C. Magnesium Nitride, Mg3N2, is obtained as a greenish .yellow amorphous mass by passing a current of nitrogen or ammonia over heated magnesium (F. Briegleb and A. Geuther, See also:Ann., 1862, 123, p. 228; see also W. Eidmann and L.

Moeser, Ber., 1901, 34, p. 390). When heated in dry oxygen it becomes incandescent, forming magnesia. Water decomposes it with liberation of ammonia and formation of magnesium hydroxide. The chlorides of See also:

nickel, See also:cobalt, See also:chromium, iron and See also:mercury are converted into nitrides when heated with it, whilst the chlorides of. See also:copper and See also:platinum are reduced to,the metals (A. Smits, Rec. Pays Bas, 1896, 15, p. 135). Magnesium sulphide, MgS, may be obtained, mixed with some unaltered metal and some magnesia, as a hard See also:brown mass by heating magnesia, in See also:sulphur vapour. It slowly decomposes in moist air. Magnesium sulphate, MgSO4, occurs (with IH2O) as Kieserite. A hexahydrate is also known.

The salt may be obtained from Kieserite: formerly it was prepared by treating magnesite or dolomite with sulphuric acid. Organic Compounds.—By heating magnesium filings with methyl and See also:

ethyl iodides A. Cahours (Ann. chim. phys., 1860, 58, pp. 5, 19) obtained magnesium methyl, Mg(See also:CH3)2, and magnesium ethyl, Mg(C2H5)2, as colourless, strongly smelling, See also:mobile liquids, which are spontaneously inflammable and are readily decomposed by water. The compounds formed by the action of magnesium on alkyl iodides in the See also:cold have been largely used in synthetic organic Grignard See also:chemistry since V. Grignard (Comptes rendus, 1900 et Reagent. seq.) observed that magnesium and alkyl or aryl. halides combined together in presence of anhydrous See also:ether at See also:ordinary temperatures (with the See also:appearance of brisk boiling) to form compounds of the type RMgX(R = an alkyl or aryl See also:group and X = halogen). These compounds are insoluble in ether, are non-inflammable and exceedingly reactive. A. V. See also:Baeyer (Ber., 1902, 35, p. 1201) regards them as oxonium salts containing tetravalent oxygen (C2H5)2 :0 :(MgR) (X), whilst W. Tschelinzeff (Ber., 1906, 39, p.

773) considers that they contain two molecules of ether. In preparing the Grignard reagent the commencement of the reaction is accelerated by a trace of iodine. W. Tschelinzeff (Ber., 1904, 37, p. 4534) showed that the ether may be replaced by See also:

benzene containing a small quantity of ether or anisole, or a few drops of a See also:tertiary amine. With unsaturated alkyl halides the products are only slightly soluble in ether, and two molecules of the alkyl See also:compound are brought into the reaction. They are very unstable, and do not react in the normal manner. (V. Grignard and L. Tissier, Comptesrendus, 1901, 132, p. p oducts formed by the action of the Grignard reagent with the various types of organic compounds are usually thrown out of solution in the form of crystalline precipitates or as thick See also:oils, and are then decomposed by See also:ice-cold dilute sulphuric or acetic acids, the magnesium being removed as a basic halide salt. A pplications.—For the formation of See also:primary and secondary See also:alcohols see See also:ALDEHYDES and See also:KETONES. Formaldehyde behaves abnormally with magnesium benzyl bromide (M.

Tiffeneau, Comptes rendus, 1903, 137, P. 573), forming ortho-tolylcarbinol, CH3•See also:

C6H4•See also:CH2OH, and not benzylcarbinol, See also:C6H;See also:CH2-CH2OH (cf. the reaction of form-aldehyde on phenols: O. Manasse, Ber. 1894, 27, p. 2904). Acid See also:esters yield carbinols, many of which are unstable and readily pass over into unsaturated compounds, especially when warmed with acetic anhydride: R•CO2R'(R")2-RiC-OMgX--(R")2R:C-OH. Formic ester yields a secondary See also:alcohol under similar conditions. Acid chlorides behave in an analogous manner to esters (Grignard and Tissier, Comptes rendus, 1901, 132, p. 683). Nitriles yield ketones (the nitrogen being eliminated as ammonia), the best yields being given by the aromatic nitriles (E. Blaise, ibid., 1901, 133, p. 1217): R-CN—>RR':C:NMgI-- R-CO•R'.

Acid amides also react to form ketones (C. Geis, ibid., 1903, 137, 575) : R•CONH2-RR':C(OMgX).NHMgX+R'H—>R-CO•R'; the yield increases with the complexity of the organic residue of the acid See also:

amide. On passing a current of dry carbon dioxide over the reagent, the gas is absorbed and the resulting compound, when decomposed by dilute acids, yields an organic acid, and similarly with carbon oxysulphide a thio-acid is obtained: RMgX->R.CO2MgX-->R.See also:CO2H; See also:COS--)CS(OMgX)•R-->R•CSOH. A Klages (Ber., 1902, 35, pp. 2633 et seq.) has shown that if one uses an excess of magnesium and of an alkyl halide with a ketone, an See also:ethylene derivative is formed. The reaction appears to be perfectly See also:general unless the ketone contains two ortho-substituent See also:groups. Organo-metallic compounds can also be prepared, for example SnBr4 +4MgBrC6H 6 = 4MgBr2+Sn (C6Hs) 4. For a See also:summary see A. See also:McKenzie, B. A. See also:Rep. 1907.

Detection.—The magnesium salts may be detected by the white precipitate formed by adding sodium phosphate (in the presence of ammonia and ammonium chloride) to their solutions. The same reaction is made use of in the quantitative determination of magnesium, the white precipitate of magnesium ammonium phosphate being converted by ignition into magnesium pyrophosphate and weighed as such. The atomic weight of magnesium has been determined by many observers. J. See also:

Berzelius (Ann. chim. phys., 1820, 14, p. 375), by converting the oxide into the sulphate, obtained the value 12.62 for the See also:equivalent. R. F. Marchand and T. Scheerer (Jour. prakt. Chem., 185o, 5o, p. 358), by ignition of the carbonate, obtained the value 24.00 for the atomic weight, whilst C.

Marignac, by converting the oxide into the sulphate, obtained the value 24.37. T. W. See also:

Richards and H. G. See also:Parker (Zeit. anorg. Chem., 1897, 13, p. 81) have obtained the value 24.365 (O =16). See also:Medicine.—These salts of magnesium may be regarded as the typical saline purgatives. Their aperient action is dependent upon the minimum of irritation of the bowel, and is exercised by their See also:abstraction from the See also:blood of water, which passes into the bowel to See also:act as a diluent of the salt. The stronger the solution administered, the greater is the quantity of water that passes into the bowel, a fact to be See also:borne in mind when the salt is administered for the purpose of draining superfluous fluid from the See also:system, as in See also:dropsy. The oxide and carbonate of magnesium are also invaluable as antidotes, since they form insoluble compounds with oxalic acid and salts of mercury, See also:arsenic, and copper.

The result is to prevent the See also:

local corrosive action of the See also:poison and to prevent absorption of the metals. As alkaloids are insoluble in alkaline solutions, the oxide and carbonate—especially the former—may be given in alkaloidal poisoning. The compounds of magnesium are not absorbed into the blood in any appreciable quantity, and therefore exert no remote actions upon other functions. This is fortunate, as the result of injecting a solution of a magnesium salt into a vein is rapid poisoning. Hence it is of the utmost importance to avoid the use of salts of this metal whenever it is necessary—as in diabetic See also:coma—to increase the alkalinity of the blood rapidly. The usual doses of the oxide and carbonate of magnesium are from See also:half a drachm to a drachm.

End of Article: MAGNESIUM [symbol Mg, atomic weight 24.32 (0 = 16)]

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