more extensively handled in commerce than the official article, partly on account of its lower price. The British Pharmacopoeia recognizes both varieties. Ammonium alum may be readily distinguished from the official alum by the evolution of an ammoniacal odor upon trituration with potassium or sodium hydroxide or carbonate; moreover, upon heating, ammonium alum leaves a final residue of pure alumina, while the residue from official alum contains potassium sulphate besides.

DRIED ALUM. ALK(SO). Crystallized potassium alum contains 45.52 per cent. of water of crystallization, which may be entirely expelled at a temperature of 200° C. (392° F.). In the official process for preparing dried or burnt alum, the crystals are first fused in a shallow capsule, the heat being then increased and continued until 10 parts have been reduced in weight to 5.5 parts and a white porous mass remains, which is preserved in powder form in tightly stoppered bottles. A temperature exceeding 200° C. (392° F.) must be avoided to prevent decomposition and change of the aluminum sulphate to alumina, with loss of sulphuric acid.

Dried alum, although completely but slowly soluble in water, requires about three or four times as much water for solution as the crystallized alum.

ALUMINUM HYDRATE OR HYDROXIDE. Al(OH), or Al(OH)3. The Pharmacopoeia directs this compound to be prepared by gradually pouring a hot solution of alum into a hot solution of an equal weight of sodium carbonate, repeatedly washing the resulting precipitate with hot water, and finally drying the residue at a temperature not above 40° C. (104° F.). The decomposition is accompanied by the evolution of carbon dioxide, and may be illustrated as follows: 3Na,CO, + AL,K,(SO,), +3H,O= AL(OH),+K2SO, + 3Na2SO,+3CO2; this peculiar reaction is characteristic of certain metals, aluminum, iron in the ferric state, and chromium, the oxides of which exhibit weak basic properties and fail to combine with carbonic acid, but are precipitated as hydroxides when their soluble. salts are acted upon by alkali carbonates.

The object of using hot solutions of the two salts and of adding the alum solution slowly to the alkaline liquid is to prevent the formation of basic aluminum sulphate and to facilitate the complete removal of alkali and sulphuric acid, which would be persistently retained by the precipitated hydroxide if the precipitation took place in the presence of an excess of alum. The use of hot liquids also facilitates the elimination of the carbon dioxide.

Drying the precipitate at a moderate temperature is desirable to insure a smooth product, as a high heat would cause partial decomposition and a gritty powder.

ALUMINUM SULPHATE. Al(SO)+16H,O. This salt is preferably prepared for medicinal purposes by dissolving freshly prepared

aluminum hydroxide in a sufficient quantity of sulphuric acid properly diluted with water. An excess of acid should be avoided, as also an excess of the hydroxide; in the event of the latter, basic sulphates are likely to be formed. 100 Gm. of aluminum hydroxide (obtained from 607.33 Gm. of official alum) require 188.31 Gm. of absolute, or 203.58 Gm. of official sulphuric acid to form a normal salt. The gelatinous hydroxide will dissolve quite readily, and the solution having been filtered is evaporated on a water-bath until a crystalline residue is obtained.

Aluminum sulphate contains about the same percentage of water of crystallization as official alum, but is far more soluble (about 8 times) than the latter.

Besides the official aluminum compounds the following is sometimes used:

SOLUTION OF ALUMINUM ACETATE. This preparation is recognized in the German Pharmacopoeia, and is prepared by adding 360 Gm. of 30 per cent. acetic acid to a solution of 300 Gm. of alumi num sulphate in 8C0 Ce. of water, and afterward introducing, in small portions at a time, a mixture of 130 Gm. of calcium carbonate in 200 Cc. of water. The whole operation must be conducted in a cool place and the mixture be allowed to stand at rest for 24 hours, when the clear liquid may be removed with the aid of a siphon. The solution contains about 7.5 or 8 per cent. of basic aluminum acetate of the composition, Al(OH)2(C,H,O,). The reaction taking place in the foregoing process may be illustrated thus: (A1,(SO,),+16H2O)+4HC2H2O2+3CaCO ̧ = Al(OH),(C2H2O2),+

3CaSO+3CO2+17H2O.

The Compounds of Cerium.

pow

CERIUM OXALATE. Ce(CO),+9HO,. This salt is prepared from the mineral cerite by a somewhat complicated process, on account of the presence of two other metals, lanthanum and didymium, which are intimately associated with cerium as silicates. The dered mineral is digested with sulphuric acid, the mass dried and treated with diluted nitric acid and hydrogen sulphide to remove copper and other metals. The cerite metals are next precipitated by means of oxalic acid, and the mixed oxalates, after the addition of magnesium carbonate, are calcined and the residue dissolved in a small quantity of concentrated nitric acid. The solution is poured into a large quantity of water containing about one-half per cent. of sulphuric acid, whereby the cerium is precipitated as yellow ceric sulphate, while lanthanum and didymium, together with the magnesium, remain in solution. The ceric sulphate is dissolved in sulphuric acid and reduced to cerous sulphate, by means of sodium

thiosulphate, after which it is precipitated, as cerous oxalate, with oxalic acid and dried.

Pure cerium oxalate is white, but the commercial article is frequently of a pink color, due to the presence of didymium, which may be detected by heating the suspected salt to redness, when a reddish-yellow residue of ceric oxide should be obtained, didymium imparting a brown color, as stated in the official test.

Among the non-official salts of cerium, the nitrate, Ce(NO3)3+ 6H,O, has been used to some extent. It may be conveniently made by decomposing cerous sulphate with barium nitrate, and possesses the advantage of being freely soluble in water and alcohol.

CHAPTER XLVII.

THE COMPOUNDS OF IRON.

THERE is no class of inorganic compounds, excepting the official preparations of the alkalies, more extensively employed in medicine than those of iron; they must therefore be considered as among the most important in the study of pharmacy. The Pharmacopoeia recognizes, besides iron in the metallic form, no less than 38 different preparations of the same, of which 12 are liquid. Chemists have grouped all compounds of iron into two classes, designated as ferrous and ferric compounds respectively, which differ from each other in striking physical and chemical properties; this distinction has also been maintained in the official titles of the iron salts and their solutions. Ferrous compounds, in which iron is bivalent, are, when not anhydrous, of a green color, with one exception, the yellow oxalate, and form a blue precipitate of ferrous ferricyanide, Fe,(Fe(CN)6)2, known as Turnbull's Blue, with solution of potassium ferricyanide; ferric compounds, in which iron is trivalent, on the other hand, are characterized by a reddish- or yellowish-brown color and form a blue precipitate of ferric ferrocyanide, Fe,(Fe(CN)。), known as Prussian Blue, with solution of potassium ferrocyanide.

The following is a list of the official preparations of iron, divided, for convenience, into three classes:

Ferric Hydrate with Magnesia,
Ferric Hypophosphite,
Ferric Valerianate,

Iron and Ammonium Citrate,
Iron and Ammonium Tartrate,
Iron and Potassium Tartrate,
Iron and Quinine Citrate,
Soluble Iron and Quinine Citrate,
Iron and Strychnine Citrate.
Soluble Ferric Phosphate,
Soluble Ferric Pyrophosphate,
Solution of Ferric Acetate,
Solution of Ferric Chloride,
Solution of Ferric Citrate,
Solution of Ferric Nitrate,
Solution of Ferric Subsulphate,
Solution of Ferric Sulphate,
Solution of Iron and Ammonium
Acetate,

Tincture of Ferric Chloride,
Iron Plaster,

Troches of Iron,

Bitter Wine of Iron,

Wine of Ferric Citrate,

IRON. Fe. The kind of metallic iron recognized in the Pharmacopoeia is that occurring in the form of soft, bright wire. It should be free from rust and the commercial article, as it has usually been coated with grease or paraffin oil to protect it from moisture, must be thoroughly cleaned before it is used for pharmaceutical purposes. The kind of iron wire known in the trade as card-teeth, obtained as clippings and waste from the manufacturers of cotton cards, is usually preferred on account of its convenient form and general good quality; sometimes, however, card-teeth of a very inferior grade are sold and require careful garbling and subsequent washing to remove grease and dirt.

REDUCED IRON. This preparation represents more or less pure metallic iron in a fine state of division, obtained by reduction of ferric oxide with hydrogen gas. Ferric hydroxide (see Ferric Hydrate) is first dried, whereby it is changed to oxyhydrate, and then placed in an iron reduction tube so arranged that the same can be heated to dull redness, while a current of hydrogen gas, previously washed and dried by being passed through a moderately strong solution of potassium permanganate and afterward sulphuric acid, is constantly passed through it. The reducing action of hydrogen on ferric oxide may be illustrated by the following equation: FeO3+ H=Fe2+3H ̧Ó. The supply of hydrogen is kept up as long as any oxygen is left, as shown by the escape of aqueous vapor from the tube. When reduction is complete the tube and contents are allowed to cool slowly, while a slow stream of hydrogen is continued until the temperature has been reduced to that of the air; this is necessary, otherwise the hot, finely divided iron will be readily re