A less dangerous method of manufacture than that just given, devised by Professor Markoe for the use of the retail pharmacist, consisted in treating phosphorus under the surface of water with bromine and a small quantity of iodine, and at the end of the reaction nitric acid is added. The reaction runs in three distinct stages. First, phosphorus and bromine unite to form phosphorus pentabromide, viz.:

This decomposes in the presence of water, with the formation of phosphoric acid and hydrobromic acid:

The bromine here freed acts on another portion of the phosphorus, and the reactions I, II, and III are repeated in regular order as long as any phosphorus remains unoxidized.

It will be seen that the ultimate result of this process is the same as that of the first process, nitric acid being the real oxidizing agent in both cases. The function of the bromine in the second process is to lessen the violence of the reaction by gently bringing the nitric acid in contact with the phosphorus, insuring a gradual evolution of the gaseous nitrogen dioxide; for be it remembered that most explosions are due either to a sudden evolution or a sudden contraction of a gas. If the process is so planned to permit the evolution of a gas during several hours, no danger ensues, but if the same amount of gas is evolved in a few moments, a violent explosion occurs (p. 422).

In both processes the liquid remaining at completion of reaction is evaporated, partly to remove the last traces of nitric and nitrous acids (also bromine, in last process), and also to concentrate the prod uct to the official strength. The last task is very tedious if carried out in an open evaporating dish, and in the writer's experience he has never succeeded in making an acid of pharmacopoeial strength, the strongest he produced being about 65 per cent. Under the pharmacopoeia of 1880, where a 50 per cent. acid was directed, the Markoe process was feasible for the retail pharmacist, but now that the pharmacopoeia demands the more concentrated form, we are brought back to the original proposition that the pharmacist had better purchase his phosphoric acid. Much of the phosphoric acid of commerce contains arsenic, and the origin of this impurity is of interest as showing how persistently impurities may sometimes cling to chemicals, even though the original substance has undergone several changes. The arsenic in the phosphoric acid is due to arsenic-contaminated phosphorus. The phosphorus obtains it from arsenic-contaminated sulphuric acid, while the sulphuric acid contaminated with arsenic is usually that manufactured from iron pyrites, in which the arsenic originally occurred. It, therefore, behooves every careful pharmacist to test both his phosphoric acid and his sodium phosphate for arsenic by the tests prescribed in the pharmacopoeia. (See Part V.)

ACIDUM PHOSPHORICUM DILUTUM-Diluted Phosphoric Acid (Acid. Phos. Dil.)

An aqueous solution containing not less than 9.5 per cent. nor more than 10.5 per cent. of H3PO (98.06). Preserve it in well-stoppered bottles.

Condensed Recipe.

Mix 100 Gm. phosphoric acid with 765 Gm. distilled water.

Summarized Description.

Clear, colorless, odorless liquid; strongly acid to taste and to litmus; sp. gr. 1.057. For details see U.S.P., p. 21.

For tests for identity, for impurities (metaphosphoric and pyrophosphoric acids; phosphorous, hypophosphorous, sulphuric and nitric acids; heavy metals, arsenic and phosphates) and for assay see U.S.P., p. 21 and also Part V of this book.

Remarks. This body represents the strong acid diluted with a sufficient quantity of water to make it contain 10 per cent. of absolute H2PO1. These two acids are employed as nervous stimulants, although for such purpose that popular remedy, acid phosphates, is more frequently employed. A recipe for this preparation may be found in the National Formulary under the title of liquor acidi phosphorici. (See p. 384.)

It will be observed that the theoretic construction of the phosphoric acid is based on the addition of three molecules of water to one molecule of phosphoric oxide, this yielding H,PO-the so-called orthophosphoric acid. As explained on p. 358, this substance is not the true "ortho-" acid. If the same oxide is combined with one molecule of water, the following reaction occurs:

metaphosphoric acid being formed. This metaphosphoric acid is called glacial phosphoric acid, because on cooling it becomes a transparent, ice-like solid.

Under this division mention may be made of pyrophosphoric acid, HP2O7. The composition of this body will be considered when discussing sodium pyrophosphate (p. 441), and mention is here made merely to call attention to the test distinguishing between the three forms, ortho-, meta-, and pyro-phosphoric acids, these three being also called respectively the tribasic, monobasic, and the tetrabasic phosphoric acids. (See p. 359.) Orthophosphoric acid does not coagulate albumin (white of egg), and forms a yellow precipitate of silver phosphate on the addition of silver nitrate solution. Metaphosphoric acid coagulates albumin, and gives a white precipitate with silver nitrate, while pyrophosphoric acid does not coagulate albumin, and gives a white precipitate with silver nitrate. Silver nitrate reactions also apply to the salts of these acids, but the albumin test can be used only with the free acids.

Dose.-2 mils (30 minims).

Liquor Phosphatum Acidus (N.F.) or solution of acid phosphates is now made by dissolving precipitated chalk and magnesium carbonate in a diluted phosphoric acid, the former method of preparation from bone ash (p. 384) being abandoned. Liquor Phosphatum Compositus (N.F.) is discussed on p. 191.

ACIDUM HYPOPHOSPHOROSUM.-Hypophosphorous Acid (Acid. Hypophos.)

An aqueous solution containing not less than 30 per cent. nor more than 32 per cent. of HPH2O2 (66.06). Preserve it in glass-stoppered bottles.

Summarized Description.

For

Colorless or faint-yellow, odorless liquid; strongly acid to taste and to litmus; sp. gr. 1.13; on heating it decomposes into phosphine and phosphorous acid. details see U.S.P., p. 15.

For tests for identity, for impurities, see diluted hypophosphorous acid. For assay see U.S.P., p. 16 and also Part V of this book.

Remarks. This official contains 30 to 32 per cent. of absolute hypophosphorous acid. The usual method of manufacture is by a modification

of the Fothergill process, potassium hypophosphite being treated with enough tartaric acid to precipitate all the potassium in the form of insoluble potassium bitartrate. Among other processes which have been suggested is the treatment of calcium hypophosphite with oxalic acid or ammonium oxalate in which case the following reaction occurs:

the calcium oxalate being then filtered from the solution of the acid. A third process is by treatment of barium hypophosphite with sulphuric acid, precipitated barium sulphate being removed by filtration.

The pharmacopoeia writes the formula of hypophosphorous acid HPH2O2. This is because it has but one replaceable hydrogen. Its composition has been the subject of considerable discussion, but now its graphic formula is supposed to be

The valence of phosphorus is five, just as in phosphoric acid. This is totally different from what we usually find in various acids of the same element, and constitutes one of the weak points in the valence theory.

ACIDUM HYPOPHOSPHOROSUM DILUTUM-Diluted Hypophosphorous Acid (Acid. Hypophos. Dil.)

An aqueous solution containing not less than 9.5 per cent. nor more than 10.5 per cent. of HPH2O2 (66.06). Preserve it in well-stoppered bottles.

Condensed Recipe.

Mix 100 Gm. hypophosphorous acid with 210 Gm. distilled water. For details see U.S.P., p. 160.

For tests for identity, for impurities (barium, oxalic acid, heavy metals, arsenic) and for assay see U.S.P., p. 16 and also Part V of this book.

Summarized Description.

Colorless, odorless liquid; strongly acid to taste and to litmus; sp. gr. 1.042. For details see U.S.P., p. 16.

Remarks. This 10 per cent. acid is directed in the present pharmacopoeia to be made by diluting the official stronger acid. Diluted hypophosphorous acid is official because a constituent of syrup of hypophos phites. It is rarely used alone, but when it is, its action is similar to that of the better-known hypophosphites of the alkalies.

Dose.-5 decimils (8 minims).

BIBLIOGRAPHY

Hydrochloric Acid.-(History) B. Valentine, Wootton, 1, 1910, 228; Glauber, Wootton, 1, 1910, 263 and 369.

Hydrobromic Acid.-(Manufacture) Wade, Dr. Circ., 21, 1877, 108; Fothergill, Dr. Circ., 21, 1877, 83; Maisch, A.J.P., 49, 1877, 42.

Hydriodic Acid.—(Manufacture) Dunn, A.Ph.A., 16, 1868, 383. (Darkening) Hansmann, A.J.P., 72, 1900, 217.

Nitric Acid.-(History) Gmelin-Kraut, 11, 1907, 290. (Manufacture) Kuhlmann, Ph. Jl., [2], 4, 1862, 155.

Nitro-hydrochloric Acid.-(History) Kopp, 3, 1845, 348. (Preparation) Priwozink, C. A., 5, 1911, 41; Goldschmidt, A., 205, 1880, 372. (Composition) Moore, Jl. Am. Ch. Soc., 33, 1911, 1091.

Sulphuric Acid.-(History) Lassar-Cohn, "Things Chemical," January, 1907; Arny, Am. Dr., 39, 1901, 3. (Structure) Berzelius, Thorpe's Dict., 5, 1913, 341;

Oddo and Anelli, Ch. Zt., 35, 1911, 837. (Lead chamber process) Lunge and Berl, Zt. angew. Ch., 19, 1906, 807, 857, and 881; Raschig, Zt. angew. Ch., 23, 1910, 2241; Berl, Zt. angew. Ch., 23, 1910, 2250; Raschig, Jl. Soc. Ch. Ind., 30, 1911, 166; Divers, Jl. Soc. Ch. Ind., 30, 1911, 594. (Contact process) Herreshoff, Jl. Am. Ch. Soc., 30, 1908, 44; Woehler, Foss and Pleuddermann, B., 39, 1906, 3458. (Consumption) Wedge, 8th Inter. Cong. App. Ch., 2, 1912, 241; Anon., Ch. Trade Jl., 57, 1915, 6; Baekeland, Jl. Ind. Eng. Ch., 7, 1915, 978.

Elixir of Vitriol.-(History) Anon., Ph. Jl., [2], 5, 1863, 109. (Manufacture) Anon., A.J.P., 13, 1842, 174.

Phosphoric Acid.—(Manufacture) Hüttner, A.Ph.A., 57, 1909, 250; Markoe, A.Ph.A., 23, 1875, 677. (Arsenic in) Anon., Dr. Circ., 44, 1900, 151 and 174. Hypophosphorous Acid.-(Manufacture) Heikel, A.J.P., 80, 1908, 583. (Use) Anon., J. A.M.A., 62, 1914, 1347.

CHAPTER XXV

THE METALS-POTASSIUM

THE elements thus far considered are grouped in the class which are called non-metals or metalloids. As already mentioned, however, a sharp distinction between the positive and the negative elements is a matter of considerable difficulty, and even so it is difficult to decide whether a given element is metallic or non-metallic. In choosing a definition for the metals, about all that can be said is that those substances are metals which possess a metallic lustre and which are good conductors of electricity. Their oxides, when they dissolve in water, usually have an alkaline reaction, although several of the elements which we are now to consider will be found to possess oxides which are more acid than basic. It, therefore, resolves the definition to the one statement that metals are good conductors of electricity.

The alkaline metals are the group of elements, lithium, sodium, and potassium, to which is added the radical, ammonium (NH4), it possessing properties so similar to the three metals just mentioned that its compounds are considered with them. It might be added, too, that the rare elements rubidium and cæsium belong to this group, but since they have no pharmaceutic application, nothing need be said of them in this work. These alkaline metals possess many characteristics in common. They all possess the valence 1, are soft and easily oxidizable substances, and are the only metals whose carbonates are soluble in water. This fact is used in analytic chemistry for separating the salts of these four metals from the other metals, as will be seen in Part V.

POTASSIUM

Symbol, K. Atomic weight, approximately 39

This element is found in carnallite, which is a mixture of potassium. and magnesium chlorides, and in sylvite, which is chiefly potassium chloride these two ores are obtained from the mines of Stassfurt, Germany. Another source of potassium is suint, or the washings of wool. The freshly sheared wool is roughly cleaned by washing in water, the washings are then concentrated, and from them is obtained wool-fat by extraction with the proper solvent. In the residue is found potassium in sufficiently large amounts to be a commercial source of the element. Potassium carbonate is also extracted from the residues in the manufacture of beet-sugar (vinasse), while large amounts are obtained from ashes of

wood by lixiviation. Lately attempts have been made to obtain potassium salts from sea weed.

Potassium was isolated by Sir Humphrey Davy in 1807 in his classic researches in electricity. He submitted potassa to electrolysis, placing a small lump on a piece of platinum foil, which was then connected with the negative pole of the battery. Bringing the positive pole in contact with the potassa the latter dissociated and small globules of metallic potassa were formed.

The electrolysis of potassium in Davy's day was too expensive a process for commercial application, and the metal was first produced commercially by Brunner by another process-by heating potassium carbonate with carbon in a retort to 1000°C., when potassium and carbon monoxide were produced, as shown in the following equation:

This process is accompanied with danger, inasmuch as carbon monoxide forms with potassium, as well as with several other metals (copper, for example), well-defined solid carbonyl derivatives. In the isolation of potassium the carbonyl, KCO, clogs up the delivery tube leading from the retort, and with the stoppage of this outlet explosion of the retort occurs. This danger has been avoided by arranging the apparatus, whereby a plunger is made to fit in the delivery tube, and by passing this plunger up and down the obstructions were removed. It should have been stated that the delivery tube terminates in a vessel containing liquid petroleum or benzin, in order that the potassium will be condensed in a medium by which it is not affected, and which will protect it from atmospheric action.

Potassium is now being made on a large scale by electrolysis of potassium hydroxide, the reduction in the cost of the production of electricity since Davy's day permitting the commercial application of Davy's

process.

Potassium is a silvery metal lighter than water and of the consistence of wax, hence it can be readily cut with a knife. A freshly cut surface, after exposure to air, becomes white and granular, due to the formation of the potassium carbonate. Metallic potassium combines so readily with oxygen that when exposed to the air the action is so energetic that the metal burns. Likewise, when the metallic potassium is thrown on the surface of water it combines with the same, forming potassium hydroxide, as shown in the following equation:

and the heat of the reaction is sufficient to cause an ignition of the hydrogen, which burns with explosive violence. Hence great care should be taken in the handling of potassium, and particular caution must be exercised in combining same with water. For this reason the metallic potassium is always kept under substances devoid of oxygen, such as benzin or petroleum.

The statements just given emphasize the care to be used in handling potassium. In weighing it one rapidly cuts from the piece the oxidized exterior and then presses the pure metal of the interior rapidly between filter-paper to remove the last traces of the liquid in which it has been kept. Thus it can be weighed upon a watch-glass, although in cases of large quantities it is best to weigh same under petroleum, leaving the drying of the metal until it has been weighed.