CHAPTER XIX

PHOSPHORUS AND THE OTHER ELEMENTS OF THE FIFTH GROUP

This

NITROGEN is the lightest and most widely distributed representative of the elements of the fifth group, which form a higher saline oxide of the form R2O, and a hydrogen compound of the form RH3. Phosphorus, arsenic, bismuth, and antimony belong to the uneven series of this group, and vanadium, niobium, and tantalum to the even series. The latter do not form compounds with hydrogen, like the elements of the even series in general (Chapter XV.); whilst phosphorus, arsenic, and antimony, on the other hand, are analogous to nitrogen, even in their property of forming RH3. Phosphorus is the most widely distributed of these elements. There is hardly any mineral substance composing the mass of the earth's crust which does not contain some-it may be small amount of phosphorus compounds in the form of the salts of phosphoric acid. The soil and earthy substances in general usually contain from one to ten parts of phosphoric acid in 10000 parts. amount, which appears so small, has, however, a very important significance in nature. No plant can attain its natural growth if it be planted in an artificial soil completely free from phosphoric acid. Plants equally require the presence of potash, magnesia, lime, and ferric oxide, among basic, and of carbonic, sulphuric, nitric, and phosphoric anhydrides, among acid, oxides. In order to increase the fertility of a more or less poor soil, the above-named nutritive elements are introduced into it by means of fertilisers. Direct experiment has proved that these substances are inevitably necessary to plants, but that they must be all present simultaneously and in small quantities, and that an excess, like an insufficiency, of one of these elements is necessarily followed by a bad harvest, or the impossibility of a perfect growth, if the sum of all other conditions (light, heat, water, air) is normal. The phosphoric compounds of the soil accumulated by plants pass into the organism of animals, in which these substances are assimilated in many cases in large quantities. Thus

the chief component part of bones is calcium phosphate, Ca,P20 ̧, on which the hardness of bones depends.1

Phosphorus was first extracted by Brand in 1669, by the ignition of evaporated urine. After the lapse of a century Scheele, who knew of the existence of a more abundant source of phosphorus in bones, pointed out the method which is now employed for the extraction of this element. Calcium phosphate in bones permeates a nitrogenous organic substance, which is called ossein, and forms a gelatin. When bones are treated exclusively for the extraction of phosphorus, neglecting the gelatin, they are burnt, in which case all the ossein is burnt away. When, however, it is desired to preserve the gelatin, the bones are immersed in cold dilute hydrochloric acid, which dissolves the calcium phosphate and leaves the gelatin untouched; calcium chloride and acid calcium phosphate, CaH (PO4)2, are then obtained in the solution. When the bones are directly burnt in an open fire, their mineral components only are left, as an ash, containing about 90 per cent. of calcium phosphate, Ca,(PO4)2, mixed with a small amount of calcium carbonate and other salts. This mass is treated with sulphuric acid, and then the same substance is obtained in the solution as was obtained from the unburnt bones immersed in hydrochloric acid-i.e. the

1 Dry bones contain about one-third of gelatinous matter and about two-thirds of ash, chiefly calcium phosphate.

The salts of phosphoric acid are also found in the mass of the earth as separate minerals; for example, the apatites contain this salt in a crystalline form, combined with calcium chloride or fluoride, CaR2,3Ca3(PO4)2, where RF or Cl, sometimes in a state of isomorphous mixture. This mineral often crystallises in fine hexagonal prisms; sp. gr., 3'17 to 3:22. Vivianite is a hydrated ferrous phosphate, Fe3(PO4)2,8H,O. Phosphates of copper are frequently found in copper mines; for example, tagilite, Cuz(PO4)2,Cu(OH)2,2H2O. Lead and aluminium form similar salts. They are nearly all insoluble in water. Sea and other waters always contain a small amount of phosphates. The ash of sea-plants, as well as of land-plants, always contains phosphates. Deposits of calcium phosphate are often met with; they are termed phosphorites and osteolites, and are composed of the fossil remains of the bones of animals; they are used for manure. Of the same nature are the so-called guano deposits from Baker's Island, and entire strata in Spain, France, and in the Governments of Orloff and Kursk in Russia. It is evident that if a soil destined for cultivation contain very little phosphoric acid, then the fertilisation by means of these minerals will be beneficial, but, naturally, only if the other elements necessary to plants be present in the soil.

Not unfrequently an untrue explanation of the conclusions made by Liebig with respect to the nourishment of plants by the component parts of the soil has led to a fervent and superfluous preaching in favour of manuring with phosphoric compounds. They are indispensable; but at the same time they are best and most economically applied together with other manures, whilst manuring with phosphoric compounds alone is not advantageous, and is therefore only necessary in very few cases-it cannot serve as a universal medicine for agriculture. It is not wise to introduce phosphoric manures in a given locality and under given conditions without making preliminary experi

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acid calcium phosphate soluble in water, in which reaction naturally the chief part of the sulphuric acid is converted into calcium sulphate: Ca3(PO4)2+2H,SO,=2CaSO,+CaH1(PO4)2. Ca3(PO4)2+4HCl =2CaCl2 + CaH1(PO4)2.

On evaporating the solution, crystallisable acid calcium phosphate is obtained. The extraction of the phosphorus from this salt consists in heating it with charcoal to a white heat. When heated, the acid phosphate, CaH,(PO4)2, first parts with water, and forms the metaphosphate, Ca(PO3)2, which, for the sake of simplicity, may be looked on, like the acid salt, as composed of pyrophosphate and phosphoric anhydride, 2Ca(PO3)2=Ca2P2O,+P205. The latter, with charcoal, gives phosphorus and carbonic oxide, P2O;+5C=P2+5CO. So that, in reality, a somewhat complicated process takes place here, which, in ultimate products, will be as follows:

2CaH (PO4)2+5C=4H2O+Ca2P20, +P2+5CO.

After the steam has come over, phosphorus and carbonic oxide distil over from the retort, and calcium pyrophosphate remains behind.2

As phosphorus melts at about 40°, it condenses at the bottom of the receiver in a molten liquid mass, which is cast under water in tubes, and is sold in the form of sticks. This is common or yellow phosphorus. It is a transparent, waxy, yellowish substance, which is not brittle, almost insoluble in water, and easily undergoes change in its external appearance and properties, under the action of light, heat, and of various substances. It crystallises (by sublimation or from its solution in carbon bisulphide) in the regular system, and is especially. characterised (in contradistinction to the other varieties) by its easy solubility in carbon bisulphide, and also partially in other oily liquids. In this it recalls common sulphur. Its specific gravity is 184. It fuses at 44°, and passes into vapour at 290°; it is easily inflammable, and must therefore be handled with great caution; careless rubbing is enough to cause phosphorus to ignite. Its application in the manufacture of matches is based on this. It emits light in the air owing to its oxidising, and is therefore kept under water (such water is phosphorescent in the dark, like phosphorus itself). It is also very easily oxidised by various oxidising agents, takes up the oxygen from many substances, and enters into direct combination with many metals and with sulphur, chlorine, &c., with development of a considerable amount of heat. It is very poisonous, although not soluble in water.

By subjecting the pyrophosphate to the action of sulphuric or hydrochloric acid it is possible to obtain a fresh quantity of the acid salt from the residue, and in this manner

Besides this, there is a red variety of phosphorus, which differs considerably from the above. Red phosphorus, also called amorphous phosphorus, owing to its showing no sign of crystallisation, is partially formed when ordinary phosphorus remains long exposed to the action of light. It is also formed in many reactions; for example, when ordinary phosphorus combines with chlorine, bromine, iodine, or oxygen, a portion of it is converted into red phosphorus. Schrötter, in Vienna, investigated this variety of phosphorus, and pointed out by what methods it may be produced in considerable quantities. Red phosphorus is a powdery red-brown opaque substance of specific gravity 214. It does not combine so energetically with oxygen and other substances as yellow phosphorus, and evolves less heat in combining with them. Common phosphorus easily oxidises in the air; red phos

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to extract all the phosphorus from the original normal salt Cas(PO4)2. It is usual to take burnt bones, but mineral phosphorites, osteolites, and apatites may also be employed as materials for the extraction of phosphorus. Its extraction for the manufacture of matches is everywhere extending, and in Russia, in the Urals, in the Government of Perm, it has attained such proportions that the district is able to supply other countries with phosphorus. great many methods have been proposed for facilitating the extraction of phosphorus, but they none of them essentially differ from the usual one, because the problem is dependent on the liberation of phosphoric acid by the action of acids, and on its ultimate reduction by charcoal. Thus the calcium phosphate may be mixed directly with charcoal and quartz (silica), and phosphorus will be liberated on heating the mixture, because the silica displaces the phosphoric anhydride, which gives carbonic oxide and phosphorus with the charcoal. It has also been proposed to pass hydrochloric acid over an incandescent mixture of calcium phosphate and charcoal; the acid then acts just as the silica does, liberating phosphoric anhydride, which is reduced by the charcoal. It is necessary to prevent the access of air in the condensation of the vapours of phosphorus, because they take fire very easily; hence they are condensed under water by causing the gaseous products to pass through a vessel full of water. For this purpose the condenser shown in fig. 83 is usually employed.

C.

FIG. 83.-Preparation of phosphorus. The mixture is calcined in the retort The vapours of phosphorus pass through a into water without coming into contact with air. The p 10sphorus condenses in the water, and the gases accompanying it escape through i.

3 The thermochemical determinations for phosphorus and its compounds date from the last century, when Lavoisier and Laplace burnt phosphorus in oxygen in an ice calorimeter. Andrews, Despretz, Favre, and others have studied the same subject. The most accurate and complete data are due to Thomsen. In order to give an idea of the indirect and complex methods by which the figures cited hereafter are obtained, it will be enough to point out that Thomsen employed the following method to determine the heat of combustion of yellow phosphorus. Phosphorus was oxidised in a calorimeter by iodic acid in the presence of water, and a mixture of phosphorous and phosphoric acids was thus formed (was not any hypophosphoric acid formed?-Salzer), and the iodic acid converted into hydriodic acid. It was first necessary to introduce two corrections into the calorimetric result obtained, one for the oxidation of the phosphorous into phosphoric acid, knowing their relative amounts by analysis, and the other for the deoxidation of

phorus does not oxidise at all at the ordinary temperature; hence it does not phosphoresce in the air, and may be very conveniently kept in the form of powder. It does not fuse at 44°. Having been converted into vapour at 290° or 300°, it, when slowly cooled, again passes into the ordinary variety. Red phosphorus is not soluble in carbon bisulphide and other oily liquids, which permits of its being freed from any admixture of the ordinary phosphorus. It is not poisonous. It is used in many cases for which the ordinary phosphorus is unsuitable or dangerous; for example, in the manufacture of matches, which are then no longer poisonous or inflammable by accidental friction, and therefore the red variety has now replaced the ordinary phosphorus.

the iodic acid. The result then obtained expresses the conversion of phosphorous into hydrated phosphoric acid. This must be corrected for the heat of solution of the hydrate in water, and for the heat of combination of the anhydride with water, before we can obtain the heat evolved in the reaction of P2 with O5 in the proportion for the formation of P2O5. It is natural that with so complex a method there is a possibility of many small errors, and the resultant figures will only present a certain degree of accuracy after repeated corrections by various methods. Of such a kind are the following figures determined by Thomsen, which we express in thousands of calories:-P2+05=370; P2+05+3H2O=400; P2+O5+ a mass of water = 405. Hence we see that P2O3+3H2O =30; 2PH 04+ a mass of water = 5. Experiment further showed that crystallised PH3O4, in dissolving in water, evolves 27 thousand calories, and fused (39°) PH30 52 thousand calories, hence the heat of fusion of H3PO4=25 thousand calories. For phosphorous acid, H-PO3, Thomsen obtained P2+O3+3H2O=250, and the solution of crystallised H3PO3 in water = −0·13, and of fused H3PO3+29. For hypophosphorous acid, H-PO2, the heats of solution are nearly the same (-0·17 and +21), and the heat of formation P2+O+3H2O=75; hence its conversion into 2H3PO, evolves 175 thousand calories, and the conversion of 2H3PO; into 2H-PO4=150 thousand calories. For the sake of convenience we will express the combination of chlorine with phosphorus, also according to Thomsen, per 2 atoms of phosphorus, P2+3Cl2 = 151, P2+5Cl2 = 210 thousand calories. In their reaction on a inass of water (with the formation of a solution), 2PC13=130, 2PC15=247, and 2POCI;=142 thousand calories.

Besides which we will cite the following data given by various observers: heat of fusion for P (that is, for 31 parts of phosphorus by weight) -0·15 thousand calories; the conversion of yellow into red phosphorus for P, from +19 to +27 thousand calories; P+H3=43, HI+PH5=24, PH3+ HBr=22 thousand calories.

4 At the ordinary temperature (20° C.) phosphorus is not oxidised by pure oxygen; oxidation only takes place with a slight rise of temperature, or the dilution of the oxygen with other gases (especially nitrogen or hydrogen), or a direct decrease of pressure. Ordinary phosphorus takes fire at a temperature (60°) at which no other known substance will burn. Its application to the manufacture of matches is based on this property. The majority of phosphorus matches are composed of common phosphorus mixed with some oxidising substance which easily gives up oxygen, such as lead dioxide, potassium chlorate, nitre, &c. For this purpose common phosphorus is carefully triturated under warm water containing a little gum; lead dioxide, and potassium nitrate are then added to the resultant emulsion, and the match ends, previously coated with sulphur or paraffin, are dipped into this preparation. After this the matches are dipped into a solution of gum and shellac, in order to preserve the phosphorus from the action of the air. When such a match containing particles of yellow phosphorus is rubbed over a rough surface, it becomes (especially at the point of rupture of the brittle gummy coating) slightly heated, and this is sufficient for the phosphorus to take fire and burn at the