(2) A city of Lydia, about 40 miles north-east of Smyrna, which stood on the southern bank of the river Hermus, at the foot of Mount Sipylus, from which circumstance it was often called for distinction’s sake Magnesia ad Sipylum. It is probable from its name that it was founded, like the city of the same name in the valley of the Maeander, by Magnesian colonists from Thessaly; but we have no authority for the fact. Nor is any mention of the town found in history till 190 b.c., when the Syrian king, Antiochus the Great, was defeated under its walls by the Roman consul I.. Scipio, who derived from his success in this campaign the surname of Asiaticus. Magnesia became a city of importance under the Roman dominion, and, though nearly destroyed by an earthquake in the reign of Tiberius, was restored by that emperor, and continued to flourish throughout the period of the Roman empire. It was one of the few towns in this part of Asia Minor that retained its prosperity under the Turkish rule; and Manisa (q.v.) is at the present day a large and flourishing town with considerable trade.

A. Silicates.— (1) Olivine, SiO 2 ∙2MgO, occasionally met with in transparent crystals (“chrysolite ”), but more frequently embedded as an admixture in lava, basalt, and other rocks; also in meteorites. (2) Augite and (3) hornblende, both MgO.SiO 2 ∙ The latter more frequently than the former forms independent rocks; both occur abundantly as components of more ordinary mixed rocks, all basalts containing augite as a principal component. (4) Serpentine, SiO 2 2H 2 O.3(SiO 2 .2MgO) + 2Aq (meaning 2H 2 O loosely combined),—a very common mineral; there are mountain masses consisting almost of it alone. In all these silicates, and more especially in (2) and (3), the MgO is partly replaced by FeO, CaO, MnO, and other isomorphous oxides, these latter not unfrequently even predominating over the MgO itself. Their names, in fact, represent genera of minerals comprising each numerous species. Asbestos (so remarkable on account of its exceptional structure, which lends itself for the production of fire-proof textile fabrics, paper, pasteboard, &c.) must be mentioned here, as some varieties fall under hornblende, others under chrysotil (included in serpentine). More purely magnesian are—(5) talc, 4SiO 2 .3MgO.H 2 O (Rammelsberg), and (6) meerschaum, 3SiO 2 ∙2MgO + 4(1)H 2 O (same authority).

B. Carbonates.— Of these the most abundant are the dolomites, all compounds of the carbonates of lime and magnesia, a;Ca.(1-tf)Mg.O.CO 2 , where x may assume almost any value down to zero, which it actually has in “magnesite,” MgOCO 2 ∙ Bitter spar, MgOCO 2 , is very similar to and isomorphous with Iceland spar, CaOCO 2 ∙

C. Soluble Salts, known chiefly as occurring in the famous salt-deposits of Stassfurth in Germany :—kieserite, MgOSO 3 .H 2 O (Epsom salt minus 6H 2 O); carnallite, MgCl 2 .KCl +6H 2 O; kainite, a hydrated compound of chlorides and sulphates of magnesium and potassium.

Any of these minerals may be used, and A (4), B, and C are actually used for the preparation of magnesium compounds. Starting from magnesite, we need only subject it to gentle ignition to obtain the oxide MgO; treatment with dilute sulphuric or muriatic acid produces the sulphate (Epsom Salt, q.vj) or the chloride, as a solution, contaminated in general,—chiefly with iron, also with alumina, and perhaps lime. The two former, after peroxidation of the iron by chlorine, may be eliminated by digestion with powdered magnesite, and filtered off. The acidified muriate solution, on evaporation and cooling, deposits transparent crystals, MgCl 2 .6H 2 O, very hygroscopic, very readily soluble in water, and (like the anhydrous salt) soluble also in alcohol. When dehydrated by heat it loses acid (HC1), so that the residue is largely contaminated with oxychlorides, MgCl 2 .icMgO. Certain of these are industrially important as cements, used by dentists and otherwise. A concentrated solution of the chloride, made into a paste with the ignited (CO 2 -free) oxide, in a few hours hardens into a stone susceptible of polishing (Sorel). To prepare the anhydrous chloride, the crystallized or dissolved hydrated salt is evaporated with sal-ammonia to dryness, and the residue (a stable double salt) is ignited in platinum. Pure chloride, MgCl 2 , remains as a clear liquid, solidifying on cooling into a compact aggregate of flexible crystal-leaves of a mother-of-pearl lustre. From the chloride the metal can be produced, either by electrolysing the fused substance, or (more conveniently) by Caron and Deville’s method: 600 grammes of the chloride, 480 grammes of powdered fluor-spar, and 230 grammes of finely cut-up sodium are mixed, and thrown, small portions at a time, into a red-hot crucible, which is then well covered. When the reaction, which is violent, is over, the contents are well stirred with an iron rod to cause the small metal-beads to unite into bigger lumps, which, after cooling, are picked from the broken-up mass. Sonstadt (who, along with Mellor, was the first to prepare magnesium industrially) substitutes for plain MgCl 2 the double chloride of potassium and magnesium, obtainable synthetically from the hydrated components without the use of sal-ammoniac, or simply by ignition of pure carnallite. To purify the crude metal Sonstadt and Mellor distil it “per descensum” out of an iron crucible provided with a pipe piercing the bottom and reaching up to almost the lid; an atmosphere of coal-gas is established to prevent oxidation of the metal in this operation, as also in the subsequent casting of the ingot.

The metal magnesium has the colour of silver, and remains unchanged in dry air; in ordinary air it tarnishes a little more readily than zinc does. It is malleable and ductile, but has little tenacity. The specific gravity is 1 ’ 75; thus it is considerably lighter than even aluminium, whose specific gravity is 2 ’ 6. It fuses and distils at about the same temperatures as zinc. It is generally sold in the form of thin ribbon, being used for the easy production of highly intense light. The ribbon kindles readily in a candle flame, and then continues burning most intensely by itself, the solid oxide produced radiating out abundant light. A wire 0 ’ 3 millimetre ( = ’ 012 inch) thick gives out the light of 74 stearine candles weighing 100 grammes (⅜ lb avoirdupois) a piece (Bunsen and Roscoe). It used to be employed for photographing at night, but is now superseded by the electric light. Magnesium has the exceptional property of combining (at a bright red heat) directly with nitrogen gas into Mg 3 N∙ 2 , (Briegleb and Geuther), a greenish-yellow powder, which, when heated in steam, yields magnesia and ammonia, NH 3 .

Of magnesium salts the most important is the sulphate, Epsom Salt (<∕.w.). This salt serves as a raw material for the preparation of two or three medicinally important substances, especially magnesia alba. When epsom salt solution is precipitated by carbonate of soda, the salt MgOCO 2 , first formed, loses carbonic acid and takes up water, forming a precipitate of the approximate composition 3MgCO 3 .Mg(OH) 2 +2 to 3 times H 2 O. When prepared by cold precipitation it forms a very light, when in the heat a somewhat denser, white powder (magnesia alba levis and ponderosa). This substance holds an important position in the history of chemistry, having served, in the hands of Black, to prove finally the individual existence of magnesia as something distinct from lime, and helped him in establishing the true relation between “caustic” and “mild” alkalies. Before his time the “causticity” was supposed to go into the “mild” substance from the fire; Black showed that it is owing to the loss of a ponderablo substance, which he called “fixed air,” that is, carbonic acid. It was in this memorable research that for the first time in chemistry the balance was used for the precise determination of quantitative relations.

Magnesia, the oxide MgO, is produced by the gentle ignition of magnesia alba. It is a white powder, absolutely infusible and non-volatile, and not reducible by charcoal and is used medicinally. The gently ignited oxide combines very slowly with water into the practically insoluble hydrate MgO.H 2 O. Magnesia when boiled with sugar-water dissolves into a solution of saccharate, and in this form is sometimes administered medicinally in lieu of plain magnesia. Magnesia alba (also the normal carbonate) dissolves rather largely in carbonic acid water. According to R. Wagner, one part of MgOCO 2 dissolves in 760 parts of water saturated by carbonic acid under 1 atmosphere pressure. Under 6 atmospheres pressure it requires only 76 parts of carbonic acid water for its solution. Dinneford’s “fluid magnesia” is a solution of such bicarbonate of magnesia. The bicarbonate solution, when allowed to stand in air, deposits crystals of hydrated normal carbonate, MgCO 3 + 3 or 5H 2 O.

Magnesia preparations play a great part in therapeutics. The oxide and basic carbonate (also the dissolved forms of saccharate and bicarbonate) are used in small doses as anti-acids, in larger ones as very mild purgatives, for children more especially. For the latter purpose, however, the sulphate is generally preferred as acting far more energetically. The nauseous bitter taste of the salt can be concealed, to some extent, by acidification of its solution with dilute sulphuric acid. Citrate of magnesia, being exceptionally free of the ti Bittererde ” taste, was introduced some thirty years ago by the French as a pleasant substitute for Epsom salt, and it has since come much into fashion everywhere, although, weight for weight, it is far less efficient than the sulphate. The preparation of dry soluble citrate offering difficulties, the French originally dispensed it exclusively in the dissolved form of “Limonade au citrate de magnésie,” a flavoured, decidedly acid solution of the salt, rendered effervescent by addition of some bicarbonate of soda immediately before corking up. In England it is generally preferred to offer the dry ingredients of the “limonade” in the form of “granular effervescent citrate of magnesia.” Magnesia alba is pounded up with an excess of citric acid crystals and a few drops of water to produce a paste of amorphous acid citrate, which is dried at a temperature below 30“ C. At higher temperatures the salt would pass into a crystalline, insoluble, and consequently therapeutically valueless modification. The citrate is mixed with bicarbonate of soda, citric acid, and sugar, made into a “dough” with alcohol, granulated, and dried. The granules, when thrown into water, dissolve with effervescence. We must not omit to state here that much of what is sold under the name is a mere concoction in which Epsom salt figures as “citrate.”

To test a solution for magnesia, remove whatever can be precipitated by means of sulphuretted hydrogen or sulphide of ammonium; then eliminate lime, baryta, and strontia by precipitation with carbonate of ammonia in the presence of sal ammoniac. The magnesia remains dissolved, and can be precipitated (and detected) by addition of phosphate of ammonia (or soda) and free ammonia; the salt PO 4 MgNH 4 .6HnO gradually separates out as a crystalline precipitate. This method of course fails when the magnesia is present from the first as phosphate; but we cannot here enter into a consideration of this or any other exceptionally difficult case. (w. d.)