CHAPTER X. PHOSPHORIC ACID AND PHOSPHATES. Ir is well known that phosphorus was first obtained from urine by Brandt (1669). This substance occurs naturally only in the form of the tribasic phosphoric acid, and in combination with organic substances. In these two forms it is met with in all organized bodies. It is an invariable ingredient of all food, and enters largely into the composition of all parts of the animal and human body. In the urine phosphorus occurs only as the acid, and, in combination with soda, lime, and magnesia, forms a regular constituent of that fluid. Phosphoric Acid. Composition, Equivalent, Formula. The common or tribasic phosphoric acid has the composition PO+3HO. This formula expresses the equivalent of 1 equiv. phosphorus 8 ditto oxygen 3 ditto hydrogen 1 equivalent of phosphoric acid, = The theory which assumes P to be a double atom, and the single atomic weight 15718, uses P2 as the symbol for the above equivalent of phosphorus. This is an explanation to the reader, should he find himself embarrassed by the formulæ of different authors. Physical properties. Phosphoric acid may be obtained in crystals forming quadrangular or hexagonal prisms of a transparency like glass, or as a syrupy fluid. When heated to 320° F., it begins to lose water, and, at a temperature of 415-4° F., is transformed into pyrophosphoric acid, PO,+2HO. When heated still more, it loses another equivalent of HO, and is then metaphosphoric acid, PO + HO. When exposed to red heat in an open platinum capsule, it is volatilized. Chemical properties. Phosphoric acid is easily soluble in water and alcohol or spirits of wine, and its solutions exhibit a strongly acid reaction. Solutions of albumen, chloride of barium or calcium, do not cause any precipitate in solutions of phosphoric acid; the solutions of caustic baryta, strontia, or lime, when added in excess, produce a white precipitate. With the basic oxydes phosphoric acid has great affinity, and forms with them the phosphates, which may contain one, two, or three equivalents of almost any of the common bases, instead of any of the three equivalents of water. Thus we may have in the urine: Phosphates of Alkalies. Phosphate of soda with two equivalents of base and one of water, PO,, 2NaO+HO (alkaline) may occur abnormally in the alkaline urine of chlorosis. Ammonio-phosphate of soda, PO,+NaO+NH2O+HO +8HO. Acid phosphate of soda, PO,+NaO+2HO. Phosphates of the Alkalies and Earths. Ammonio-phosphate of magnesia, PO+2MgO+NHO +12HO. Phosphates of the Alkaline Earths. Phosphate of lime (acid), PO,+2CaO+HO. Of these, the phosphates of the alkalies are easily soluble in water; the others are scarcely soluble, or altogether insoluble. They all dissolve in nitric or hydrochloric acid. The phosphates of the alkaline earths, when newly precipitated, are, moreover, soluble in acetic acid. The solution of an earthy phosphate in an acid, when neutralized with an alkali, throws down a precipitate of the original phosphate, which is insoluble in an excess of the alkali. The insoluble phosphates are, as a rule, soluble in an excess of a solution of any salt from which they have been precipitated by phosphate of soda;1 this solution becomes very turbid by heating, and clears again on cooling.2 With the common soluble phosphates, nitrate of silver produces a yellow precipitate (PO,+3AgO), which is soluble in nitric acid, and in ammonia. Acetate of lead produces a white precipitate (PO,+3PbO), which is soluble in nitric acid, but insoluble in acetic acid and in ammonia. If chlorides are present, the precipitate contains chloride of lead in chemical combination. Chloride of barium and chloride of calcium produce white precipitates with the soluble phosphates (PO+3BaO and PO+3CaO), each readily soluble in hydrochloric, nitric, or acetic acid. When a solution of phosphate of lime in acetic acid is allowed to stand some time, the phosphate has a great inclination to fall down from this solution in a crystalline state, particularly when the mixture is warmed a little, and when the phosphate is prevalent. Phosphate of lime is somewhat soluble in water containing carbonic acid, and in salts of ammonia, even when free ammonia is present. From its solution in acetic acid, or from its solution in hydrochloric acid when mixed with acetate of soda (which is virtually a solution in acetic acid, because hydrochloric acid, combining with the soda, sets acetic acid free, which is now the solvent for the phosphate of lime), oxalate of ammonia throws down the whole amount of lime as oxalate of lime. From its solution in hydrochloric or nitric acid, the entire amount of lime may be precipitated by means of sulphuric acid and alcohol. These reactions are the bases for the quantitative analysis of lime in ashes and the earth of bones. A mixture of sulphate of magnesia, or chloride of magnesium, with chloride of ammonium and ammonia, produces a crystalline precipitate in soluble phosphates, which has the composition PO+2MgO, NH0+12HO, is easily soluble in all acids, somewhat soluble in pure water, and perfectly insoluble in water containing ammonia, even if a large amount of any salt of ammonia should be present. This precipitate, after exposure to red heat, is of the composition PO+2MgO, and is the qualitative and quantitative test for phosphoric acid (in absence of arsenic acid) in all combinations, which are soluble in water, the watery solution of which does, Enderlin, Ann. d. Chem. und Pharm.,' 1844, p. 320. 2 Dr. G. Owen Rees, Guy's Hospital Rep.,' vol. i, p. 402. however, not become turbid by admixture of a solution of chloride of ammonium and ammonia. In very dilute solutions the precipitate forms only slowly. When the solution contains tartaric acid and oxyde of iron, some tartrate of magnesia and oxyde of iron may easily be mixed with the precipitate. Chloride of iron produces in solutions of phosphates a yellowish-white precipitate of PO5+ Fe,О, which is soluble in hydrochloric acid, in an excess of chloride of iron, in acetate of iron, and in ammonia. This precipitate is, however, quite insoluble in acetic acid, and will, for this reason, form even when its solution in hydrochloric acid is mixed with acetate of soda, as already explained, or when the solution in hydrochloric acid of the phosphate of an alkaline earth is mixed with a small quantity of chloride of iron, and with acetate of soda. If the solution of any phosphate in hydrochloric acid, after any excess of the acid has been neutralized by a little ammonia or carbonate of soda, is mixed first with acetate of soda, and then with chloride of iron in slight excess (which may be recognised by the fluid assuming a reddish colour), and is then heated to ebullition, a reddishbrown precipitate is obtained, which contains the whole amount of oxyde of iron and all the phosphoric acid present. It is filtered hot, and the precipitate and filter are washed with hot water. This test forms the basis of the volumetrical analysis of phosphoric acid by Professor Liebig, and of the method for removing all phosphoric acid and iron from solutions in which the quantity of lime and magnesia and isomorphous oxydes has yet to be determined. On mixing a solution of the nitrate of protoxyde of mercury with a solution of phosphate of soda, a white flocculent precipitate of phosphate of protoxyde of mercury is immediately produced, which on being allowed to stand in the fluid, rapidly becomes crystalline. A solution of corrosive sublimate, however, may be mixed with the alkaline phosphate without any turbidity being produced. If to a mixture of the two first-mentioned salts we add a solution of chloride of sodium before the precipitate has had time to become crystalline, the latter will immediately decompose with the chloride of sodium, corrosive sublimate and phosphate of soda being produced: the precipitate disappears, and the fluid becomes perfectly clear. This test is the basis of Professor Liebig's method for ascertaining the amount of protoxyde of mercury contained in a solution of its nitrate. (Vide p. 59 et seq.) Method of ascertaining the amount of Phosphoric Acid in the Urine. I have already described the test upon which this method, also by Professor Liebig, is based. By adding to the urine, containing an unknown amount of phosphoric acid, some acetate of soda, and then a nearly neutral solution of chloride of iron of known strength, until the entire amount of phosphoric acid is precipitated, and a trace of the solution of chloride of iron can be discovered to have been added in excess, we ascertain from the amount of iron in the test-fluid the amount of phosphoric acid precipitated in the form of PO+FO. The excess of chloride of iron remaining in solution causes the latter to give a blue precipitate when mixed with some ferrocyanide of potassium. Before this test could be applied, a filtration of the fluid from the precipitate would be required. But this is avoided by the following simple manipulation: a piece of bibulous paper is saturated with a solution of ferrocyanide of potassium, and spread over a white china plate, or a glass disc lying on a piece of white paper. A drop of the solution to be tested is now allowed to fall on a second slip of dry bibulous paper, spread over the one saturated with the ferrocyanide. The solution thus filtered mixes with the ferrocyanide of potassium in the moistened paper, and, when it contains an excess of chloride of iron, causes a blue spot to appear on the paper. In the performance of this analysis there is, however, one caution to be observed; namely, to take the first blue test obtained as the mark of the completion of the analysis. For if any excess of chloride of iron is present in the fluid only for a very short time, the precipitate PO+ FeO takes up more iron, and the blue spot is now no longer obtained with the fluid. A few drops of the chloride of iron added to the fluid will immediately allow it to reappear. Thus an excess of the test-fluid might be added, which would unduly increase the apparent amount of phosphoric acid. This error is avoided by the observance of the above caution. Preparation of the solution of Iron of known strength. 15.556 grammes of pure iron (piano-forte string) are dissolved in as much hydrochloric acid as necessary. After addition of a little nitric acid, the solution is evaporated to |