XIX. PHOSPHORUS AND THE OTHER ELEMENTS OF THE FIFTH GROUP

XIX. PHOSPHORUS AND THE OTHER ELEMENTS OF THE FIFTH GROUP

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 R_{2}O_{5}, and a hydrogen compound of the form RH_{3}. Phosphorus, arsenic, bismuth, and antimony belong to the uneven series of this group. 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 a 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 10,000 parts. This 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 undoubtedly 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 an imperfect growth, even if all the other conditions (light, heat, water, air) are normal. The phosphoric compounds of the soil accumulated by plants pass into the organism of animals, in which these substances are assimilated in many instances in large quantities. Thus the chief component part of bones is calcium phosphate, Ca_{3}P_{2}O_{8}, and it is on this that their hardness depends.[1]

[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, CaR_{2},3Ca_{3}(PO_{4}){2}, where R = F 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, Fe{3}(PO_{4}){2},8H{2}O. Phosphates of copper are frequently found in copper mines; for example, tagilite, Cu_{3}(PO_{4}){2},Cu(OH){2},2H_{2}O. Lead and aluminium form similar salts. They are nearly all insoluble in water. The turquoise, for instance, is hydrated phosphate of alumina, (Al_{2}O_{3}){2},P{2}O_{5}5H_{2}O, coloured with a salt of copper. 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, 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.

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_{4}(PO_{4}){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{3}(PO_{4})_{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 acid calcium phosphate soluble in water, in which reaction naturally the chief part of the sulphuric acid is converted into calcium sulphate:

Ca_{3}(PO_{4}){2} + 2H{2}SO_{4} = 2CaSO_{4} + CaH_{4}(PO_{4}){2}. Ca{3}(PO_{4}){2} + 4HCl = 2CaCl{2} + CaH_{4}(PO_{4})_{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_{4}(PO_{4}){2}, first parts with water, and forms the metaphosphate, Ca(PO{3}){2}, which for the sake of simplicity may be regarded, like the acid salt, as composed of pyrophosphate and phosphoric anhydride, 2Ca(PO{3}){2} = Ca{2}P_{2}O_{7} + P_{2}O_{5}. The latter, with charcoal, gives phosphorus and carbonic oxide, P_{2}O_{5} + 5C = P_{2} + 5CO. So that in reality a somewhat complicated process takes place here, yielding ultimately products according to the following equation:

2CaH_{4}(PO_{4}){2} + 5C = 4H{2}O + Ca_{2}P_{2}O_{7} + P_{2} + 5CO.

After the steam has come over, phosphorus and carbonic oxide distil over from the retort and calcium pyrophosphate remains behind.[1 bis]

[1 bis] 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 to extract all the phosphorus. 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. A great many methods have been proposed for facilitating the extraction of phosphorus, but none of them differ essentially 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 sand, 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.

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, yellowish, waxy 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[2] (in contradistinction to the other varieties) is easily soluble in carbon bisulphide, and also partially in other oily liquids. In this it recalls common sulphur. Its specific gravity is 1·84. 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.[2 bis] It emits light in the air owing to its slow[3] oxidation, 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 and takes up the oxygen from many substances.[3 bis] Phosphorus 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.

[2] Vernon (1891) observed that ordinary (yellow) phosphorus is dimorphous. If it be melted and by careful cooling be brought in a liquid form to as low a temperature as possible, it gives a variety which melts at 45°·3 (the ordinary variety fuses at 44°·3), sp. gr. 1·827 (that of the ordinary variety is 1·818) at 13°, crystallises in rhombic prisms (instead of in forms belonging to the cubical system). This is similar to the relation between octahedral and prismatic sulphur (Chapter XX.).

[2 bis] According to Herr Irinyi (an Hungarian student), the first phosphorus matches were made in Austria at Roemer's works in 1835.

[3] The absorption of the oxygen of the atmosphere at a constant ordinary temperature by a large surface of phosphorus proceeds so uniformly, regularly, and rapidly, that it may serve, as Ikeda (Tokio, 1893) has shown, for demonstrating the law of the velocity (rate) of reaction, which is considered in theoretical chemistry, and shows that the rate of reaction is proportional to the active mass of a substance--i.e. dx/dt = k(A - x) where t is the time, A the initial mass of the reacting substance--in this case oxygen--x the amount of it which has entered into reaction, and k the coefficient of proportionality. Ikeda took a test-tube (diameter about 10 mm.), and covered its outer surface with a coating of phosphorus (by melting it in a test-tube of large diameter, inserting the smaller test-tube, and, when the phosphorus had solidified, breaking away the outer test-tube), and introduced it into a definite volume of air, contained in a Woulfe's bottle (immersed in a water bath to maintain a constant temperature), one of whose orifices was connected with a mercury manometer showing the fall of pressure, x. Knowing that the initial pressure of the oxygen (in air nearly 750 × ·0209) was about 155 mm. = A, the coefficient of the rate of reaction k is given, by the law of the variation of the rate of reaction with the mass of the reacting substance, by the equation: k = (1/t)log(A/(A - x)), where t is the time, counting from the commencement, of the experiment in minutes. When the surface of the phosphorus was about 11 sq. cm., the following results were actually obtained.

t = 10 20 30 40 50 60 minutes x = 10·5 21·5 31·1 40·7 49·1 57·3 mm 10,000 k = 32 32 32 33 33 33

The constancy of k is well shown in this case. The determination takes a comparatively short time, so that it may serve as a lecture experiment, and demonstrates one of the most important laws of chemical mechanics.

[3 bis] Not only do oxidising agents like nitric, chromic, and similar acids act upon phosphorus, but even the alkalis are attacked--that is, phosphorus acts as a reducing agent. In fact it reduces many substances, for instance, copper from its salts. When phosphorus is heated with sodium carbonate, the latter is partially reduced to carbon. If phosphorus be placed under water slightly warmed, and a stream of oxygen be passed over it, it will burn under the water.

Besides this, there is a red variety of phosphorus, which differs considerably from the above. Red phosphorus (sometimes wrongly called amorphous phosphorus) is partially formed when ordinary phosphorus remains exposed to the action of light for a long time. 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 2·14. It does not combine so energetically with oxygen and other substances as yellow phosphorus, and evolves less heat in combining with them.[4] Common phosphorus easily oxidises in the air; red phosphorus 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, like yellow phosphorus, fuse at 44°. After being converted into vapour at 290° or 300°, it again passes into the ordinary variety when slowly cooled. 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, and is used in many cases for which the ordinary phosphorus is unsuitable or dangerous; for example, in the manufacture of matches, which are then not poisonous or inflammable by accidental friction, and therefore the red variety has now replaced the ordinary phosphorus.[4 bis]

[4] 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. To determine the heat of combustion of yellow phosphorus, Thomsen oxidised it in a calorimeter with 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 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 P_{2} with O_{5} in the proportion for the formation of P_{2}O_{5}. 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:--P_{2} + O_{5} = 370; P_{2} + O_{3} + 3H_{2}O = 400; P_{2} + O_{5} + a mass of water = 405. Hence we see that P_{2}O_{5} + 3H_{2}O = 30; 2PH_{3}O_{4} + an excess of water = 5. Experiment further showed that crystallised PH_{3}O_{4}, in dissolving in water, evolves 2·7 thousand calories, and that fused (39°) PH_{3}O_{4} evolves 5·2 thousand calories, whence the heat of fusion of H_{3}PO_{4} = 2·5 thousand calories. For phosphorous acid, H_{3}PO_{3}, Thomsen obtained P_{2} + O_{3} + 3H_{2}O = 250, and the solution of crystallised H_{3}PO_{3} in water = -0·13, and of fused H_{3}PO_{3} = +2·9. For hypophosphorous acid, H_{3}PO_{2}, the heats of solution are nearly the same (-0·17 and +2·1), and the heat of formation P_{2} + O + 3H_{2}O = 75; hence its conversion into 2H_{3}PO_{3} evolves 175 thousand calories, and the conversion of 2H_{3}PO_{3} into 2H_{3}PO_{4} = 150 thousand calories. For the sake of comparison we will take the combination of chlorine with phosphorus, also according to Thomsen, per 2 atoms of phosphorus, P_{2} + 3Cl_{2} = 151, P_{2} + 5Cl_{2} = 210 thousand calories. In their reaction on an excess of water (with the formation of a solution), 2PCl_{3} = 130, 2PCl_{5} = 247, and 2POCl_{3} = 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 + H_{3} = 4·3, HI + PH_{3} = 24, PH_{3} + HBr = 22 thousand calories.

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 decrease of pressure.

[4 bis] 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. In order to illustrate the easy inflammability of common (yellow) phosphorus, its solution in carbon bisulphide may be poured over paper; this solvent quickly evaporates, and the free phosphorus spread over a large surface takes fire spontaneously, notwithstanding the cooling effect produced by the evaporation of the bisulphide. 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 to cause the phosphorus to take fire and burn at the expense of the oxygen of the other ingredients.

The heads of the 'safety' matches do not contain any phosphorus, but only substances capable of burning and of supporting combustion. Red phosphorus is spread over a surface on the box, and it is the friction against this phosphorus which ignites the matches. There is no danger of the matches taking fire accidentally, nor are they poisonous.[5] This red phosphorus is prepared by heating the ordinary phosphorus at 230° to 270°; it is evident that this must be done in an atmosphere incapable of supporting combustion--for example, in nitrogen, carbonic anhydride, steam, &c. On a large scale, ordinary phosphorus is placed in closed iron vessels,[5 bis] and immersed in a bath of different proportions of tin and lead, by which means the temperature of 250° necessary for the conversion is easily attained. It is kept at this temperature for some time. The temperature is at first cautiously raised, and the air is thus partially expelled by the heat, and also by the evolution of steam (the phosphorus is damp when put in), whilst the remaining oxygen is also partially absorbed by the phosphorus, so that an atmosphere of nitrogen is produced in the iron vessel. Red phosphorus enters into all the reactions proper to yellow phosphorus, only with greater difficulty and more slowly;[6] and, as its vapour tension (volatility) is less than that of the yellow variety, it may be supposed that a polymerisation takes place in the passage of the yellow into the red modification, just as in the passage of cyanogen into paracyanogen, or of cyanic acid into cyanuric acid (Chapter IX. Notes 39 bis and 48).

[5] In the so-called 'safety' or Swedish matches (which are not poisonous, and do not take fire from accidental friction) a mixture of red phosphorus and glass forms the surface on which the matches are struck, and the matches themselves do not contain any phosphorus at all, but a mixture of antimonious sulphide, Sb_{2}S_{3} (or similar combustible substances) and potassium chlorate (or other oxidising agents). The combustion, when once started by contact with the red phosphorus, proceeds by itself at the expense of the inflammatory and combustible elements contained in the tip of the match. The mixture applied on the match itself must not be liable to take fire from a blow or friction. The mixture forming the heads of the 'safety' matches has the following approximate composition: 55-60 parts of chlorate of potassium, 5-10 parts of peroxide of manganese (or of K_{2}Cr_{2}O_{7}), about 1 part of sulphur or charcoal, about 1 part of pentasulphide of antimony, Sb_{2}S_{5}, and 30-40 parts of rouge and powdered glass. This mixture is stirred up in gum or glue, and the matches are dipped into it. The paper on which the matches are struck is coated with a mixture of red phosphorus and trisulphide of antimony, Sb_{2}S_{3}, stirred up in dextrine.

[5 bis] Phosphorus only acts on iron at a red heat. The boiler is provided with a safety valve and gas-conducting tube, which is immersed in mercury or other liquid to prevent the admission of air into the boiler.

[6] The specific heat of the yellow variety is 0·189--that is, greater than that of the red variety, which is 0·170. The sp. gr. of the yellow is 1·84, and of the red prepared at 260° 2·15, and of that prepared at 580° and above (i.e. 'metallic' phosphorus, see below) = 2·34. At 230° the pressure of the vapour of ordinary phosphorus = 514 millimetres of mercury, and of the red = 0--that is to say, the red phosphorus does not form any vapour at this temperature; at 447° the vapour tension of ordinary phosphorus is at first = 5500 mm., but it gradually diminishes, whilst that of red phosphorus is equal to 1636 mm.

Hittorf, by heating the lower portion of a closed tube containing red phosphorus to 530° and the upper portion to 447°, obtained crystals of the so-called 'metallic' phosphorus at the upper extremity. As the vapour tensions (according to Hittorf, at 530° the vapour tension of yellow phosphorus = 8040 mm., of red = 6139 mm., and of metallic = 4130 mm.) and reactions are different, metallic phosphorus may be regarded as a distinct variety. It is still less energetic in its chemical reaction than red phosphorus, and it is denser than the two preceding varieties: sp. gr. = 2·34. It does not oxidise in the air; is crystalline, and has a metallic lustre. It is obtained when ordinary phosphorus is heated with lead for several hours at 400° in a closed vessel, from which the air has been exhausted. The resultant mass is then treated with dilute nitric acid, which first dissolves the lead (phosphorus is electro-negative to lead, and does not, therefore, act on the nitric acid at first) and leaves brilliant rhombohedral crystals of phosphorus of a dark violet colour with a slight metallic lustre, which conduct an electric current incomparably better than the yellow variety; this also is characteristic of the metallic state of phosphorus.

The researches of Lemoine partially explain the passage of yellow (ordinary) phosphorus into its other varieties. He heated a closed glass globe containing either ordinary or red phosphorus, in the vapour of sulphur (440°), and then determined the amount of the red and yellow varieties after various periods of time, by treating the mixture with carbon bisulphide. It appeared that after the lapse of a certain time a mixture of definite and equal composition is obtained from both--that is, between the red and yellow varieties a state of equilibrium sets in like that of dissociation, or that observed in double decompositions. But at the same time, the progress of the transformation appeared to be dependent on the relative quantity of phosphorus taken per volume of the globe (i.e. upon the pressure). Neglecting the latter, we will cite as an example the amounts of the red phosphorus transformed into the ordinary, and of the ordinary not converted into red, per 30 grams of red or yellow taken per litre capacity of the globe, heated to 440°. When red phosphorus was taken, 4·75 grams of yellow phosphorus were formed after two hours, four grams after eight hours, three grams after twenty-four hours, and the last limit remained constant on further heating. When thirty grams of yellow phosphorus were taken, five grams remained unaltered after two hours, four grams after eight hours, and after twenty-four hours and more three grams as before. Troost and Hautefeuille showed that liquid phosphorus in general changes more easily into the red than does phosphorus vapour, which, however, is able, although slowly, to deposit red phosphorus.

The question presents itself as to whether phosphorus in a state of vapour is the ordinary or some other variety? Hittorf (1865) collected many data for the solution of this problem, which leave no doubt that (as experimental figures show) the density of the vapour of phosphorus is always the same, although the vapour tension of the different varieties and their mixtures is very variable. This shows that the different varieties of phosphorus only occur in a liquid and solid state, as indeed is implied in the idea of polymerisation. Strictly speaking, the vapour of phosphorus is a particular state of this substance, and the molecular formula P_{4} refers only to it, and not to any other definite state of phosphorus. But Raoult's solution method showed that in a benzene solution the fall of the freezing point indicates for ordinary phosphorus a molecule P_{4}, judging by the determinations of Paterno and Nasini (1888), Hirtz (1890), and Beckmann (1891), who obtained for sulphur by the same method a molecular weight = S_{6}, in conformity with the vapour density. Further research in this direction will perhaps show the possibility of finding the molecular weight of red phosphorus, if a means be discovered for dissolving it without converting it into the yellow variety.

I think it will not be out of place here to draw the reader's attention to the fact that red phosphorus, which we must recognise as polymeric with the yellow, stands nearer to nitrogen, whose molecule is N_{2}, in its small inclination towards chemical reactions, although judging by its small vapour tension it must be more complex than ordinary (yellow and white) phosphorus.

The vapour of phosphorus is colourless; its density remains constant between 300° and 1000° (Dumas, 1833; Mitscherlich, Deville, and Troost, 1859, and others). The density with respect to air has been determined as from 4·3 to 4·5. Hence, referred to hydrogen, it is 4·4 × 14·4 = 63, corresponding with a molecular weight 124, i.e. the molecule of phosphorus in a state of vapour contains P_{4}. The reader will remember that the molecule of nitrogen contains N_{2}, of sulphur S_{6} or S_{2}, and of oxygen O_{2} or O_{3}.

The chemical energy of phosphorus in a free state more nearly approaches that of sulphur than nitrogen. Phosphorus is combustible and inflames at 60°; but having in the act of combination parted with a portion of its energy in the form of heat it becomes analogous to nitrogen, so long as there is no question of its reduction back again into phosphorus. Nitric acid is easily reduced to nitrogen, whilst phosphoric acid is reduced with very much greater difficulty. All the compounds of phosphorus are less volatile than those of nitrogen. Nitric acid, HNO_{3}, is easily distilled; metaphosphoric acid, HPO_{3}, is generally said to be non-volatile; triethylamine, N(C_{2}H_{5}){3}, boils at 90°, and triethylphosphine, P(C{2}H_{5})_{3}, at 127°.

Phosphorus not only combines easily and directly with oxygen, but also with chlorine, bromine, iodine, sulphur, and with certain metals, and red phosphorus when heated combines with hydrogen also.[6 bis] So, for instance, when fused with sodium under naphtha, phosphorus gives the compound Na_{3}P_{2}. Zinc, absorbing the vapour of phosphorus, gives the phosphide Zn_{3}P_{2} (sp. gr. 4·76); tin, SnP; copper, Cu_{2}P; even platinum combines with phosphorus (PtP_{2}, sp. gr. 8·77).[6 tri] Iron, when combined even with a small quantity of phosphorus, becomes brittle.[7] Some of these compounds of phosphorus are obtained by the action of phosphorus on the solutions of metallic salts, and by the ignition of metallic oxides in the vapour of phosphorus, or by heating mixtures of phosphates with charcoal and metals. Phosphides do not exhibit the external properties of salts, which are so clearly seen in the chlorides and still distinctly observable in the sulphides. The phosphides of the metals of the alkalis and of the alkaline earths are even immediately and very easily decomposed by water, whereas this is found to be the case with only a very few sulphides, and still more rarely and indistinctly with the chlorides. We may take calcium phosphide as an example.[7 bis] Phosphorus is laid in a deep crucible, and covered with a clay plug, over which lime is strewn. At a red heat the vapours of phosphorus combine with the oxygen of the lime and form phosphoric anhydride, which forms a salt with another portion of the lime, whilst the liberated calcium combines with the phosphorus and forms calcium phosphide. Its composition is not quite certain; it may be CaP (corresponding with liquid phosphuretted hydrogen). This substance is remarkable for the following reaction: if we take water--or, better still, a dilute solution of hydrochloric acid--and throw calcium phosphide into it, bubbles of gas are evolved, which take fire spontaneously in the air and form white rings. This is owing to the fact that the liquid hydrogen phosphide, PH_{2}, is first formed, thus, CaP + 2HCl = CaCl_{2} + PH_{2}, which, owing to its instability, very easily splits up into the solid phosphide, P_{2}H, and gaseous phosphide, PH_{3}; 5PH_{2} = P_{2}H + 3PH_{3}; the latter corresponds with ammonia. The mixture of the gaseous and liquid phosphides takes fire spontaneously in the air, forming phosphoric acid. The same hydrogen phosphides are formed when water acts on sodium phosphide (P_{2}Na_{3}). A similar mixture of gaseous liquid and solid phosphuretted hydrogen (Retgers 1894) is formed by heating (in a glass tube) red phosphorus in a stream of dry hydrogen. Hence we see that there are three compounds of phosphorus with hydrogen. (1) The first or solid yellow phosphide, P_{2}H (more probably P_{4}H_{2}), is obtained by the action of strong hydrochloric acid on sodium phosphide; it takes fire when struck or at 175°. (2) The liquid, PH_{2}, or more correctly expressed as the molecule, P_{2}H_{4}, is a colourless liquid which takes fire spontaneously in the air, boils at 30°, is very unstable, and is easily decomposed (by light or hydrochloric acid) into the two other phosphides of hydrogen. It is prepared by passing the gases evolved by the action of water on calcium phosphide through a freezing mixture.[8] And, lastly, (3), gaseous hydrogen phosphide, phosphine, PH_{3}, which is distinguished as being the most stable. It is a colourless gas, which does not take fire in the air. It has an odour of garlic, and is very poisonous. It resembles ammonia in many of its properties.[8 bis] It is easily decomposed by heat, like ammonia, forming phosphorus and hydrogen; but it is very slightly soluble in water, and does not saturate acids, although it forms compounds with some of them which resemble ammonium salts in their form and properties. Among them the compound with hydriodic acid, PH_{4}I, analogous to ammonium iodide, is remarkable. This compound crystallises on sublimation in well-formed cubes, like sal-ammoniac, which it resembles in many respects. However, this compound does not enter into those reactions of double decomposition which are proper to sal-ammoniac, because its saline properties are very feebly developed. Phosphuretted hydrogen also combines, like ammonia, with certain chloranhydrides; but they are decomposed by water, with the evolution of phosphine. Ogier (1880) showed that hydrochloric acid also combines with phosphine under a pressure of 20 atmospheres at +18°, and under the ordinary pressure at -35°, forming the crystalline phosphonium chloride PH_{4}Cl, corresponding to sal-ammoniac. Hydrobromic acid does the same with greater ease, and hydriodic acid with still greater facility, forming phosphonium iodide, PH_{4}I.[9]

[6 bis] Retgers (see further on) showed this in 1894, and observed that As when heated also combines with hydrogen.

[6 tri] The capacity of mercury (Chapter XVI., Note 25 bis) to give unstable compounds with nitrogen gives rise to the supposition that similar compounds exist with phosphorus also. Such a compound was obtained by Granger (1892) by heating mercury with iodide of phosphorus in a closed tube at 275°-300°. After removing the iodide of mercury formed, there remain fine rhombic crystals having a metallic lustre, and composition Hg_{3}P_{2}. This compound is stable, does not alter at the ordinary temperature and only decomposes at a red heat; when heated in air it burns with a flame. Nitric and hydrochloric acids do not act upon it, but it is easily decomposed by aqua regia. A phosphide of copper, Cu_{2}P_{2}, was obtained by Granger (1893) by heating a mixture of water, finely divided copper and red phosphorus in a sealed tube to 130°. The excess of copper was afterwards washed away by a solution of NH_{3} in the presence of air.

[7] The metallic compounds of phosphorus possess a great chemical interest, because they show a transition from metallic alloys (for instance, of Sb, As) to the sulphides, halogen salts, and oxides, and on the other hand to the nitrides. Although there are already many fragmentary data on the subject, the interesting province of the metallic phosphides cannot yet be regarded as in any way generalised. The varied applications (phosphor-iron, phosphor-bronze, &c.), which the phosphides have recently acquired should give a strong incentive to the complete and detailed study of this subject, which would, in my opinion, help to the explanation of chemical relations beginning with alloys (solutions) and ending with salts and the compounds of hydrogen (hydrides), because the phosphor-metals, as is proved by direct experiment, stand in the same relation to phosphuretted hydrogen as the sulphides do towards sulphuretted hydrogen, or as the metallic chlorides to hydrochloric acid.

[7 bis] Many other compounds of phosphorus are also capable of forming phosphuretted hydrogen. Thus BP also gives PH_{3} (see Chapter XVII., Note 12). According to Lüpke (1890) phosphuretted hydrogen is formed by phosphide of tin. The latter is prepared by treating molten tin covered with a layer of carbonate of ammonium, with red phosphorus; 200-300 c.c. of water are then poured into a flask, 3-5 grams of this phosphide of tin dropped in, and after driving out the air by a stream of carbonic acid, hydrochloric acid (sp. gr. 1·104) is poured in. The disengagement of phosphuretted hydrogen takes place on heating the flask in a water bath. The following is another easy method for preparing PH_{3}. A mixture of 1 part of zinc dust (fume) and 2 parts of red phosphorus are heated in an atmosphere of hydrogen (the mixture burns in air). Combination takes place accompanied by a flash, and a grey mass of Zn_{3}P_{2} is formed which gives PH_{3} when treated with dilute H_{2}SO_{4}.

[8] The spontaneous inflammability of the hydride PH_{2} in air is very remarkable, and it is particularly interesting that its analogues in composition, P(C_{2}H_{5}){2} (the formula must be doubled) and Zn(C{2}H_{5})_{2}, also take fire spontaneously in air.

[8 bis] The analogy between PH_{3} and NH_{3} is particularly clear in the hydrocarbon derivatives. Just as NH_{2}R, NHR_{2}, and NR_{3}, where R is CH_{3}, and other hydrocarbon radicles, correspond to NH_{3}, so there are actually similar compounds corresponding to PH_{3}. These compounds form a branch of organic chemistry.

[9] The periodic law and direct experiment (the molecular weight) show that PH_{3} is the normal compound of P and H and that it is more simple than PH_{2} or P_{2}H_{4}, just as methane, CH_{4}, is more simple than ethane, C_{2}H_{6}, whose empirical composition is CH_{3}. The formation of liquid phosphuretted hydrogen may be understood from the law of substitution. The univalent radicle of PH_{3} is PH_{2}, and if it is combined with H in PH_{3} it replaces H in liquid phosphuretted hydrogen, which thus gives P_{2}H_{4}. This substance corresponds with free amidogen (hydrazine), N_{2}H_{4} (Chapter VI.) Probably P_{2}H_{4} is able to combine with HI, and perhaps also with 2HI, or other molecules--that is, to give a substance corresponding to phosphonium iodide.

Phosphonium iodide, PH_{4}I, may be prepared, according to Baeyer, in large quantities in the following manner:--100 parts of phosphorus are dissolved in dry carbon bisulphide in a tubulated retort: when the mixture has cooled, 175 parts of iodide are added little by little, and the carbon bisulphide is then distilled off, this being done towards the end of the operation in a current of dry carbonic anhydride at a moderate temperature. The neck of the retort is then connected with a wide glass tube, and the tubulure with a funnel furnished with a stopcock, and containing 50 parts of water. This water is added drop by drop to the phosphorous iodide, and a violent reaction takes place, with the evolution of hydriodic acid and phosphonium iodide. The latter collects as crystals in the glass tube and the retort itself. It is purified by further distillations; more than 100 parts may be obtained. Baeyer expresses the reaction by the equation P_{2}I + 2H_{2}O = PH_{4}I + PO_{2}; and the compound PO_{2} may be represented as phosphorous phosphoric anhydride: P_{2}O_{5} + P_{2}O_{3} = 4PO_{2}. As a better proportion we may take 400 grams of phosphorus, 680 grams of iodine, and 240 grams of water, and express the formation thus: 13P + 9I + 21H_{2}O = 3H_{4}P_{2}O_{7} + 7PH_{4}I + 2HI (Chapter XI., Note 77).

Phosphonium iodide and even phosphine act as reducing agents in solutions of many metallic salts. Cavazzi showed that with a solution of sulphurous anhydride phosphine gives sulphur and phosphoric acid.

Phosphuretted hydrogen, or phosphine, PH_{3}, is generally prepared by the action of caustic potash on phosphorus.[10] Small pieces of phosphorus are dropped into a flask containing a strong solution of caustic potash and heated. Potassium hypophosphite, H_{2}KPO_{2}, is then obtained in solution; gaseous phosphuretted hydrogen is evolved:

P_{4} + 3KHO + 3H_{2}O = 3(KH_{2}PO_{2}) + PH_{3}.

Liquid phosphuretted hydrogen (and free hydrogen) is also formed, together with the phosphine, so that the gaseous product, on escaping from the water into the air, takes fire spontaneously, forming beautiful white rings of phosphoric acid. In this experiment, as in that with calcium phosphide, it is the liquid, P_{2}H_{4}, that takes fire; but the phosphine set light to by it also burns, PH_{3} + O_{4} = PH_{3}O_{4}. The same phosphuretted hydrogen, PH_{3}, may be obtained pure, and not spontaneously combustible, by igniting the hydrates of phosphorous acid (4PH_{3}O_{3} = PH_{3} + 3PH_{3}O_{4}) and hypophosphorous acid (2PH_{3}O_{2} = PH_{3} + PH_{3}O_{4}); or, more simply, by the decomposition of calcium phosphide by hydrochloric acid, because then all the liquid phosphide, P_{2}H_{4}, is decomposed into non-volatile P_{2}H and gaseous PH_{3}. Pure phosphine liquefies when cooled to -90°, boils at -85°, and solidifies at -135° (Olszewski). When phosphorus burns in an excess[10 bis] of dry oxygen, then only phosphoric anhydride, P_{2}O_{5} is formed. It is prepared by dropping pieces of phosphorus through a wide tube, fixed into the upper neck of a large glass globe, on to a cup suspended in the centre of the globe. These lumps are set alight by touching them with a hot wire, and the phosphorus burns into P_{2}O_{5}. The dry air necessary for its combustion is forced into the globe through a lateral neck, and the white flakes of phosphoric anhydride formed are carried by the current of air through a second lateral neck into a series of Woulfe's bottles, where they settle as friable white flakes. Phosphoric anhydride may also be formed by passing dry air through a solution of phosphorus in carbon bisulphide. All the materials for the preparation of this substance must be carefully dried, because it combines with great eagerness with water, at the same time developing a large amount of heat and forming metaphosphoric acid, HPO_{3}, from which the water cannot be separated by heat. Phosphoric anhydride is a colourless snow-like substance, which attracts moisture from the air with the utmost avidity. It fuses at a red heat, and then volatilises. Its affinity for water is so great that it takes it up from many substances. Thus it converts sulphuric acid into sulphuric anhydride, and carbohydrates (wood, paper) are carbonised, and give up the elements of water when brought into contact with it.

[10] The air must first be expelled from the flask by hydrogen, or some other gas which will not support combustion, as otherwise an explosion might take place owing to the spontaneous inflammability of the phosphuretted hydrogen.

The combustion of phosphuretted hydrogen in oxygen also takes place under water when the bubbles of both gases meet, and it is very brilliant. The phosphuretted hydrogen obtained by the action of phosphorus on caustic potash always contains free hydrogen, and often even the greater part of the gas evolved consists of hydrogen.

Pure phosphuretted hydrogen (not containing hydrogen or liquid or solid phosphides) is obtained by the action of a solution of potash on phosphonium iodide: PH_{4}I + KHO = PH_{3} + KI + H_{2}O (in just the same way as ammonia is liberated from ammonium chloride). The reaction proceeds easily, and the purity of the gas is seen from the fact that it is entirely absorbed by bleaching powder and is not spontaneously inflammable. Its mixture with oxygen explodes when the pressure is diminished (Chapter XVIII., Note 8). The vapours of bromine, nitric acid, &c., cause it to again acquire the property of inflaming in the air; that is, they partially decompose it, forming the liquid hydride, P_{2}H_{4}. Oppenheim showed that when red phosphorus is heated at 200° with hydrochloric acid in a closed tube it forms the compound PCl_{3}(H_{3}PO_{3}), together with phosphine.

[10 bis] If there be a deficiency of oxygen, phosphorous anhydride P_{2}O_{3} is formed. It was obtained by Thorpe and Tutton (1890) and is easily volatilised, melts at 22°·5, boils without change (in an atmosphere of N_{2} or CO_{2}) at 173°, and is therefore easily separated from P_{2}O_{3}, which volatilises with difficulty. The vapour density shows that the molecular weight is double, i.e. P_{4}O_{6} (like As_{2}O_{3}). Although colourless, phosphorous anhydride (its density in a state of fusion at 24° = 1·936) turns yellow and reddens in sun-light (possibly red phosphorus separates out ?), and decomposes at 400° forming hypophosphorous anhydride P_{2}O_{4} (Note 11) and phosphorus. It passes into P_{2}O_{5} in air and oxygen, and when slightly heated in oxygen becomes luminous, and ultimately takes fire. Cold water slowly transforms P_{2}O_{3} into phosphoric acid, but hot water gives an explosion and leads to the formation of PH_{3}, (P_{4}O_{6} + 6H_{2}O = PH_{3} + 3PH_{3}O_{4}). Alkalis act in the same manner. It takes fire in chlorine and forms POCl_{3} and PO_{2}Cl, and combines with sulphur at 160°, forming P_{2}S_{2}O_{3} (the molecular formula is double this) a substance which volatilises in vacuo and is decomposed by water into H_{2}S and phosphoric acid, i.e. it may be regarded as P_{2}O_{5}, in which O_{2} has been replaced by two atoms of sulphur. Judging from the above, the mixture of P_{2}O_{3} and P_{2}O_{5} formed in the combustion of phosphorus in air is transformed into P_{2}O_{5} in an excess of oxygen.

When moist phosphorus slowly oxidises in the air, it not only forms phosphorous and phosphoric acids, but also hypophosphoric acid, H_{4}P_{2}O_{6}, which when in a dry state easily splits up at 60° into phosphorous and metaphosphoric acids (H_{4}P_{2}O_{6} = H_{3}PO_{3} + HPO_{3}), but differs from a mixture of these acids in that it forms well-characterised salts, of which the sodium salt, H_{2}Na_{2}P_{2}O_{6}, is but slightly soluble in water (the sodium salts of phosphoric and phosphorous acids are easily soluble), and that it does not act as a reducing agent, like mixtures containing phosphorous acid.[11]

[11] Salzer proved the existence of hypophosphoric acid (it is also called subphosphoric acid), in which many chemists did not believe. Drawe (1888) and Rammelsberg (1892) investigated its salts. It may be obtained in a free state by the following method. The solution of acid produced by the slow oxidation of moist phosphorus is mixed with a solution (25 p.c.) of sodium acetate. A salt, Na_{2}H_{2}P_{2}O_{6},6H_{2}O, crystallises out on cooling; it is soluble in 45 parts of water, and gives a precipitate of Pb_{2}P_{2}O_{6} with lead salts (Ag_{4}P_{2}O_{6} with salts of silver). The lead salt is decomposed by a current of hydrogen sulphide, when lead sulphide is precipitated, while the solution, evaporated under the receiver of an air-pump, gives crystals of H_{4}P_{2}O_{6},2H_{2}O, which easily lose water and give H_{4}P_{2}O_{6}. The salts in which the H_{4} is replaced by Ni_{2}, or NiNa_{2}, or CdNa_{2}, &c., are insoluble in water.

In order to see the relation between phosphoric acid and hypophosphoric acid which does not contain the elements of phosphorous acid (because it does not reduce either gold or mercury from their solutions), but which nevertheless is capable of being oxidised (for example, by potassium permanganate) into phosphoric acid, it is simplest to apply the law of substitution. This clearly indicates the relation between oxalic acid, (COOH){2}, and carbonic acid, OH(COOH). The relation between the above acids is exactly the same if we express phosphoric acid as OH(POO{2}H_{2}), because in this case P_{2}H_{4}O_{6}, or (POO_{2}H_{2}){3}, will correspond with it just as oxalic does with carbonic acid. A similar relationship exists between hyposulphuric or dithionic acid, (SO{2}OH){2}, and sulphuric acid, OH(SO{2}OH), as we shall find in the following chapter. Dithionic acid corresponds with the anhydride S_{2}O_{5}, intermediate between SO_{2} and SO_{3}; oxalic acid with C_{2}O_{3}, intermediate between CO and CO_{2}; hypophosphoric acid corresponds with the anhydride P_{2}O_{4}, intermediate between P_{2}O_{3} and P_{2}O_{5}, and the analogue of N_{2}O_{4}.

Judging by the general law of the formation of acids (Chapter XV.), the series of phosphorus compounds should include the following ortho-acids and their corresponding anhydrides, answering to phosphuretted hydrogen, H_{3}P:--

H_{3}PO_{4}, phosphoric acid, and P_{2}O_{5}, anhydride, H_{3}PO_{3}, phosphorous acid, and P_{2}O_{3}, anhydride, H_{3}PO_{2}, hypophosphorous acid, and P_{2}O, anhydride.[12]

The last of these (the analogue of N_{2}O) is almost unknown. Phosphoric anhydride (P_{2}O_{5}) with a small quantity of water does not at first give orthophosphoric acid, PH_{3}O_{4}, but a compound P_{2}O_{5},H_{2}O, or PHO_{3}, whose composition corresponds with that of nitric acid; this is metaphosphoric acid. Even with an excess of water, combining with phosphoric anhydride, this metaphosphoric acid, and not the ortho-, passes at first into solution. Metaphosphoric acid in solution only passes into orthophosphoric acid when the solution is heated or after a lapse of time.

[12] Besides the hydrates enumerated, a compound, PH_{3}O, should correspond with PH_{3}. This hydrate, which is analogous to hydroxylamine, is not known in a free state, but it is known as triethylphosphine oxide, P(C_{2}H_{5}){3}O, which is obtained by the oxidation of triethylphosphine, P(C{2}H_{5}){3}. It must be observed that there may also be lower oxides of phosphorus corresponding with PH{3}, like N_{2}O and NO, and there are even indications of the formation of such compounds, but the data concerning them cannot be considered as firmly established.

Orthophosphoric acid[13] is obtained by oxidising phosphorus with nitric acid until the phosphorus entirely passes into solution and the lower oxides of nitrogen cease to be evolved. The reaction takes place best with dilute nitric acid, and when aided by heat. The resultant solution is evaporated to a syrup. If a weighed quantity of phosphorus (dried in a current of dry carbonic anhydride) be taken, a crystalline mass of the acid can be obtained by evaporating the solution until it consists only of the quantity[14] of phosphoric acid corresponding with the amount of phosphorus taken (from 31 parts of P, 98 parts of solution). The acid fuses at +39°; specific gravity of the liquid 1·88. Phosphorus pentachloride, PCl_{5}, and oxychloride, POCl_{3} (see further on), give orthophosphoric acid and hydrochloric acid with water. The two other varieties of phosphoric acid, with which we shall presently become acquainted, give the same ortho-acid when under the influence of acids, with particular ease when boiled and more slowly in the cold. By itself orthophosphoric acid (either in solution or when dry) does not pass into the other varieties; it does not oxidise, and therefore presents the limiting and stable form. When heated to 300°, it loses water and passes into pyrophosphoric acid, 2H_{3}PO_{4} = H_{2}O + H_{4}P_{2}O_{7}, whilst at a red heat it loses twice as much water and is converted into metaphosphoric acid, H_{3}PO_{4} = H_{2}O + HPO_{3}. In aqueous solution orthophosphoric acid differs clearly from pyro- or metaphosphoric acids, because the solutions of these latter acids give different reactions: thus orthophosphoric acid does not precipitate albumin, does not give a precipitate with barium chloride, and forms a yellow precipitate of silver orthophosphate, Ag_{3}PO_{4}, with silver nitrate (in the presence of alkalis, but not otherwise); whilst a solution of pyrophosphoric acid, H_{4}P_{2}O_{7}, although it does not precipitate albumin or barium chloride, gives a white precipitate of silver pyrophosphate, Ag_{4}P_{2}O_{7}, with silver nitrate; and a solution of metaphosphoric acid, HPO_{3}, precipitates both albumin and barium chloride, and gives a white precipitate of silver metaphosphate, AgPO_{3}, with silver nitrate. These points of distinction were studied by Graham, and are exceedingly instructive. They show that the solution of a substance does not determine the maximum of chemical combination with water, that solutions may contain various degrees of combination with water, and that there is a clear difference between the water serving for solution and that entering into chemical combination. Graham's experiments also showed that the water whose removal or combination determines the conversion of ortho- into meta- and pyrophosphoric acids differs distinctly from water of crystallisation, for he obtained the salts of ortho-, meta-, and pyrophosphoric acids with water of crystallisation, and they differed in their reactions, like the acids themselves. This water of crystallisation was expelled with greater ease than the water of constitution of the hydrates in question.[14 bis]

[13] Phosphoric acid, being a soluble and almost non-volatile substance, cannot be prepared like hydrochloric and nitric acids by the action of sulphuric acid on the alkali phosphates, although it is partially liberated in the process. For this purpose the salts of barium or lead may be taken, because they give insoluble salts, thus Ba_{3}(PO_{4}){2} + 3H{2}SO_{4} = 3BaSO_{4} + 2H_{3}PO_{4}. Bone ash contains, besides calcium phosphate, sodium and magnesium phosphates, and fluorides and other salts, so that it cannot give directly a pure phosphoric acid.

[14] If this is not done the orthophosphoric acid, PH_{3}O_{4}, loses a portion of its water, and then, as with an excess of water, it does not crystallise.

[14 bis] The difference between the reactions of ortho-, meta- and pyrophosphoric acids, established by Graham (see p. 163), is of such importance for the theory of hydrates and for explaining the nature of solutions, that in my opinion its influence upon chemical thought has been far from exhausted. At the present time many such instances are known both in organic (for instance, the difference between the reactions of the solutions of certain anhydrides and hydrates of acids), and inorganic chemistry (for example, the difference between the rose and purple cobalt compounds, Chapter XXII. &c.) They essentially recall the long known and generalised difference between C_{2}H_{4} (ethylene), C_{2}H_{6}O (ethyl alcohol = ethylene + water), and C_{4}H_{10}O (ethyl ether = 2 ethylene + water = 2 alcohol - water); but to the present day the numerous analogous phenomena existing among inorganic substances are only considered as a simple difference in degrees of affinity, distinguishing the water of constitution (hydration), crystallisation, and solution without penetrating into the difference of the structure or distribution of the elements, which exists here and gives rise to a distinct isomerism of solutions. In my opinion the progress of chemistry, especially with regard to solutions, should make rapid strides when the cause of the isomerism of solutions, for instance, of ortho- and pyrophosphoric acids, has become as clear to us as the cause of many well-studied instances of the isomerism, polymerism, and metamerism of organic compounds. Here it forms one of those many important problems which remain for the chemistry of the future in a state of only indistinct presentiments and in the form of facts empirically known but insufficiently comprehended.

Orthophosphoric acid has a pleasant acid taste and a distinctly acid reaction; it is used as a medicine, and is not poisonous (phosphorous acid is poisonous). Alkalis, like sodium, potassium, and ammonium hydroxides, saturate the acid properties of phosphoric acid when taken in the ratio 2NaHO : H_{3}PO_{4}--that is, when salts of the composition HNa_{2}PO_{4} are formed. When taken in the ratio NaHO : H_{3}PO_{4}, a solution having an acid reaction is obtained, and when 3NaHO : H_{3}PO_{4}--that is, when the salt Na_{3}PO_{4} is formed--an alkaline reaction is obtained. Hence many chemists (Berzelius) even regarded the salts of composition R_{2}HPO_{4} as normal, and considered phosphoric acid to be bibasic. But the salt Na_{2}HPO_{4} also shows a feeble alkaline reaction, so that it is impossible to judge the characteristic peculiarities of acids by the reactions on litmus paper, as we already know from many examples. Orthophosphoric acid is tribasic, because it contains three equivalents of hydrogen replaceable by metals, forming salts, such as NaH_{2}PO_{4}, Na_{2}HPO_{4}, and Na_{3}PO_{4}. It is also tribasic, because with silver nitrate its soluble salts always give Ag_{3}PO_{4},[15] a salt with three equivalents of silver, and because by double decomposition with barium chloride it forms a salt of the composition Ba_{3}(PO_{4}){2}, and silver and barium hardly ever give basic salts. With the metals of the alkalis, phosphoric acid forms soluble salts, but the normal salts of the metals of the alkaline earths, R{3}(PO_{4}){2} and even R{2}H_{2}(PO_{4}), are insoluble in water, but dissolve in feeble acids, such as phosphoric and acetic, because they then form soluble acid salts, especially RH_{4}(PO_{4})_{2}.[16]

[15] Silver orthophosphate, Ag_{3}PO_{4}, is yellow, sp. gr. 7·32, and insoluble in water. When heated it fuses like silver chloride, and if kept fused for some length of time it gives a white pyrophosphate (the decomposition which causes this is not known). It is soluble in aqueous solutions of phosphoric, nitric, and even acetic acids, of ammonia, and many of its salts. If silver nitrate acts on a dimetallic orthophosphate--for instance, Na_{2}HPO_{4}--it still gives Ag_{3}PO_{4}, nitric acid being disengaged: Na_{2}HPO_{4} + 3AgNO_{3} = Ag_{3}PO_{4} + 2NaNO_{3} + HNO_{3}. When alcohol is added to silver orthophosphate, Ag_{3}PO_{4}, dissolved in syrupy phosphoric acid, it precipitates a white salt (the alcohol takes up the free phosphoric acid) having the composition Ag_{2}HPO_{4}, which is immediately decomposed by water into the normal salt and phosphoric acid.

[16] The researches of Thomsen showed that in very dilute aqueous solutions the majority of monobasic acids--nitric, acetic, hydrochloric, &c. (but hydrofluoric acid more and hydrocyanic less)--HX evolve the following amounts of heat (in thousands of calories) with caustic soda: NaHO + 2HX = 14; NaHO + HX = 14; 2NaHO + HX = 14; that is, if n be a whole number _n_NaHO + HX = 14 and NaHO + _n_HX = 14. Hence reaction here only takes place between one molecule of NaHO and one molecule of acid, and the remaining quantity of acid or alkali does not enter into the reaction. In the case of bibasic acids, H_{2}R´´ (sulphuric, dithionic, oxalic, sulphuretted hydrogen, &c.), NaHO + 2H_{2}R´´ = 14; NaHO + H_{2}R´´ = 14; 2NaHO + H_{2}R´´ = 28; _n_NaHO + H_{2}R´´ = 28; that is, with an excess of acid (NaHO + 2H´{2}R´´) 14 thousand units of heat are developed, and with an excess of alkali 28. When phosphoric acid is taken (but not all tribasic acids--for instance, not citric) the general character of the phenomenon is similar to the preceding, namely, NaHO + 2H{3}PO_{4} = 14·7; NaHO + H_{3}PO_{4} = 14·8; 2NaHO + H_{3}PO_{4} = 27·1; 3NaHO + H_{3}PO_{4} = 34·0; 6NaHO + H_{3}PO_{4} = 35·3; or, in general terms, NaHO + _n_H_{3}PO_{4} = 14 (approximately) and _n_NaHO + H_{3}PO_{4} = 35 and not 42, which shows a peculiarity of phosphoric acid. In the case of energetic acids, when one equivalent (23 grams) of sodium (in the form of hydroxide) replaces one equivalent (1 gram) of hydrogen (with the formation of water and in dilute solutions), 14,000 heat units are evolved; and this is true for phosphoric acid when in H_{3}PO_{4}, Na or Na_{2} replaces H or H_{2}, but when Na_{3} replaces H_{3} less heat is developed. This will be seen from the following scheme based on the preceding figures: H_{3}PO_{4} + NaHO = 14·8; NaH_{2}PO_{4} + NaHO = 12·3; Na_{2}HPO_{4} + NaHO = 5·9; with Na_{3}PO_{4} + NaHO, a very small amount of heat is evolved, as may be judged from the fact that Na_{3}PO_{4} + 3NaHO = 1·3, but still heat is evolved. It must be supposed that in acting on phosphoric acid in the presence of a large quantity of water, a certain portion of the sodium hydroxide remains as alkali uncombined with the acid. Thus, on increasing the mass of the alkali, heat is still evolved, and a fresh interchange between Na and H takes place. Hence water shows a decomposing action on the alkali phosphates. The same decomposing action of water is seen, but to a less extent, with Na_{2}HPO_{4}, as may be judged both from the reactions of this salt and from the amount of heat developed by NaH_{2}PO_{4} with NaHO. Such an explanation is in accordance with many facts concerning the decomposition of salts by water already known to us. Recent researches made by Berthelot and Louguinine have confirmed the above deductions made by me in the first edition (1871) of this work. At the present time views of this nature are somewhat generally accepted, although they are not sufficiently strictly applied in other cases. As regards PH_{3}O_{4} it may be said that: on the substitution of the first hydrogen this acid acts as a powerful acid (like HCl, HNO_{3}, H_{2}SO_{4}); on the substitution of the second hydrogen as a weaker acid (like an organic acid); and on the substitution of the third, as an alcohol, for instance phenol, having the properties of a feeble acid.

Phosphoric anhydride, or any of its hydrates, when ignited with an excess of sodium hydroxide, carbonate, &c., forms normal or trisodium orthophosphate, Na_{3}PO_{4}, but when a solution of sodium carbonate is decomposed by orthophosphoric acid, only the salt Na_{2}HPO_{4} is formed; and when an excess of sodium chloride is ignited with orthophosphoric acid, hydrochloric acid is evolved, and the acid salt H_{2}NaPO_{4} alone is formed. These facts clearly indicate the small energy of phosphoric acid with respect to the formation of the tri-metallic salt, which is seen further from the fact that the salt Na_{3}PO_{4} has an alkaline reaction, decomposes in the presence of water and carbonic acid, forming Na_{2}HPO_{4}, corrodes glass vessels in which it is boiled or evaporated, just like solutions of the alkalis, disengages, like them, ammonia from ammonium chloride, and crystallises from solutions, as Na_{3}PO_{4},12H_{2}O, only in the presence of an excess of alkali. At 15° the crystals of this salt require five parts of water for solution; they fuse at 77°.

Disodium orthophosphate, or common sodium phosphate, Na_{2}HPO_{4}, is more stable both in solution and in the solid state. As it is used in medicine and in dyeing, it is prepared in considerable quantities, most frequently from the impure phosphoric acid obtained by the action of sulphuric acid on bone ash. The solution thus formed--which contains, besides phosphoric and sulphuric acids, salts of sodium, calcium, and magnesium--is heated, and sodium carbonate added so long as carbonic anhydride is disengaged. A precipitate is formed containing the insoluble salts of magnesium and calcium, whilst the solution contains sodium phosphate, Na_{2}HPO_{4}, with a small quantity of other salts, from which it may be easily purified by crystallisation. At the ordinary temperature its solutions, especially in the presence of a small amount of sodium carbonate, give finely-formed inclined prismatic crystals, Na_{2}HPO_{4},12H_{2}O; when the crystallisation takes place above 30° they only contain 7H_{2}O. The former crystals even lose a portion of their water of crystallisation at the ordinary temperature (the salt effloresces), and form the second salt with 7H_{2}O; whilst under the receiver of an air-pump and over sulphuric acid they also part with this water.[17] When ignited they lose the last molecule of water of constitution, and give sodium pyrophosphate, Na_{4}P_{2}O_{7}.

[17] Na_{2}HPO_{4},12H_{2}O has a sp. gr. 1·53. Poggiale determined the solubility in 100 parts of water (1) of the anhydrous ortho-salt Na_{2}HPO_{4}, and (2) of the corresponding pyro-salt Na_{4}P_{2}O_{7}:--

0° 20° 40° 80° 100° I. 1·5 11·1 30·9 81 108 II. 3·2 6·2 13·5 30 40

At temperatures of 20° to 100° the ortho-salt is so very much less soluble that this difference alone already indicates the deeply-seated alteration in constitution which takes place in the passage from the ortho- to the pyro-salts.

Monosodium orthophosphate, NaH_{2}PO_{4}, crystallises with one equivalent of water; its solution has an acid reaction. At 100° the salt only loses this water of crystallisation, and at about 200° it parts with all its water, forming the metaphosphate NaPO_{3}. It is prepared from ordinary sodium phosphate by adding phosphoric acid until the solution does not give a precipitate with barium chloride, and then evaporating and crystallising the solution. The solution of this salt does not absorb carbonic anhydride, and does not give a precipitate with salts of calcium, barium, &c.[18]

[18] The ammonium orthophosphates resemble the sodium salts in many respects, but the instability of the di- and tri-metallic salts is seen in them still more clearly than in the sodium salts; thus (NH_{4}){3}PO{4}, and even (NH_{4}){2}HPO{4}, lose ammonia in the air (especially when heated, even in solutions); NH_{4}H_{2}PO_{4} alone does not disengage ammonia and has an acid reaction. The crystals of the first salt contain 3H_{2}O, and are only formed in the presence of an excess of ammonia; both the others are anhydrous, and may be obtained like the sodium salts. When ignited these salts leave metaphosphoric acid behind; for example, (NH_{4}){2}HPO{4} = 2NH_{3 + H_{2}O + HPO_{3}. Ammonia also enters into the composition of many double phosphates. Ammonium sodium orthophosphate, or simply phosphate, NH_{4}NaHPO_{4},4H_{2}O, crystallises in large transparent crystals from a mixture of the solutions of disodium phosphate and ammonium chloride (in which case sodium chloride is obtained in the mother liquid), or, better still, from a solution of monosodium phosphate saturated with ammonia. It is also formed from the phosphates in urine when it ferments. This salt is frequently used in testing metallic compounds by the blow-pipe, because when ignited it leaves a vitreous metaphosphate, NaPO_{3}, which, like borax, dissolves metallic oxides, forming characteristic tinted glasses.

When a solution of trisodium phosphate is added to a solution of a magnesium salt it gives a white precipitate of the normal orthophosphate Mg_{2}(PO_{4}){2},7H{2}O. If the trisodium salt be replaced by the ordinary salt, Na_{2}HPO_{4}, a precipitate is also formed, and MgHPO_{4},7H_{2}O is obtained. It might be thought that the normal salt Mg_{3}(PO_{4})_{2} would be precipitated if disodium phosphate was added to ammonia and a salt of magnesium, but in reality ammonium magnesium orthophosphate, MgNH_{4}PO_{4},6H_{2}O, is precipitated as a crystalline powder, which loses ammonia and water when ignited, and gives a pyrophosphate, Mg_{2}P_{2}O_{7}. This salt occurs in nature as the mineral struvite, and in various products of the changes of animal matter. If we consider that the above salt parts with ammonia with difficulty, and that the corresponding salt of sodium is not formed under the same conditions (MgNaPO_{4},9H_{2}O is obtained by the action of magnesia on disodium phosphate), if we turn our attention to the fact that the salts of calcium and barium do not form double salts as easily as magnesium, and remember that the salts of magnesium in general easily form double ammonium salts, we are led to think that this salt is not really a normal, but an acid salt, corresponding with Na_{2}HPO_{4}, in which Na_{2} is replaced by the equivalent group NH_{3}Mg.

The common normal calcium phosphate, Ca_{3}(PO_{4}){2}, occurs in minerals, in animals, especially in bones, and also probably in plants, although the ash of many portions of plants, as a rule, contains less lime than the formation of the normal salt requires. Thus 100 parts of the ash (from 5,000 parts of grain) of rye grain contain 47·5 of phosphoric anhydride and only 2·7 of lime, and even the ash of the whole of the rye (including the straw) contains twice as much phosphoric anhydride as lime, and the normal salt contains almost equal weights of these substances. Only the ash of grasses, and especially of clover, and of trees, contains in the majority of cases more lime than is required for the formation of Ca{3}P_{2}O_{8}. This salt, which is insoluble in water, dissolves even in such feeble acids as acetic and sulphurous, and even in water containing carbonic acid. The latter fact is of immense importance in nature, since by reason of it rain water is able to transfer the calcium phosphates in the soil into solutions which are absorbed by plants. The solubility of the normal salt in acids takes place by virtue of the formation of an acid salt, which is evident from the quantity of acid required for its solution, and more especially from the fact that the acid solutions when evaporated give crystalline scales of the acid calcium phosphate, CaH_{4}(PO_{4}){2}, soluble in water. This solubility of the acid salt forms the basis of the treatment by acids of bones, phosphorites, guano, and other natural products containing the normal salt and employed for fertilising the soil. The perfect decomposition requires at least 2H{2}SO_{4} to Ca_{3}(PO_{4})_{2}, but in reality less is taken, so that only a portion of the normal salt is converted into the acid salt. Hydrochloric acid is sometimes used. (In practice such mixtures are known as superphosphates). Certain experiments, however, show that a thorough grinding, the presence of organic, and especially of nitrogenous, substances, and the porous structure of some calcium phosphates (for example, in burnt bones), render the treatment of phosphoric manures by acids superfluous--that is, the crop is not improved by it.

As a hydrate, orthophosphoric acid should be expressed, after the fashion of other hydrates, as containing three water residues (hydroxyl groups), i.e. as PO(OH){3}. This method of expression indicates that the type PX{5}, seen in PH_{4}I, is here preserved, with the substitution of X_{2} by oxygen and X_{3} by three hydroxyl groups. The same type appears in POCl_{3}, PCl_{5}, PF_{5}, &c. And if we recognise phosphoric acid as PO(OH){3}, we should expect to find three anhydrides corresponding with it: (1) [PO(OH){2}]{2}O, in which two of the three hydroxyls are preserved; this is pyrophosphoric acid, H{4}P_{2}O_{7}. (2) PO(OH)O, where only one hydroxyl is preserved. This is metaphosphoric acid. (3) (PO){2}O{3} or P_{2}O_{5}, that is, perfect phosphoric anhydride. Therefore, pyro- and metaphosphoric acids are imperfect anhydrides (or anhydro-acids) of orthophosphoric acid.[19]

[19] In this sense the ortho-acid itself might be regarded as an anhydro-acid, counting P(HO){5} as the perfect hydrate, if PH{5} existed; but as in general the normal hydrates correspond with the existing hydrogen compounds with the addition of up to 4 atoms of oxygen, therefore PH_{3}O_{4} is the normal acid, just as SH_{2}O_{4} and ClHO_{4}; while NHO_{3}, CH_{2}O_{3} are meta-acids, or higher normal acids (NH_{3}O_{4} and CH_{4}O_{4}) with the loss of a molecule of water.

In order to see the relation between the ortho-, pyro-, and metaphosphoric acids, the first thing to remark in them is that the anhydride P_{2}O_{5} is combined with 3, 2, and 1 molecules of water. In the absence of data for the molecular weight of ortho- and pyrophosphoric acids it is necessary to mention that all existing data for metaphosphoric acid indicate (Note 21) that its molecule is much more complex and contains at least H_{3}P_{3}O_{9} or H_{6}P_{6}O_{18}. The explanation of the problems which here present themselves can, it seems to me, be only looked for after a detailed study of the phenomena of the polymerisations of mineral substances, and of those complex acids, such as phosphomolybdic, which we shall hereafter describe (Chapter XXI.) A similar instance is exhibited in the solubility of hydrate of silica (produced by the action of silicon fluoride on water) in fused metaphosphoric acid, with the formation, on cooling, of an octahedral compound (sp. gr., 3·1) containing SiO_{2},P_{2}O_{5}. A certain indication (but no proof) that ordinary orthophosphoric acid is polymerised is given by Staudenmaier (1893), who obtained a salt, K_{5}H_{4}P_{3}O_{12}, by the action of a solution of KH_{2}PO_{4} upon K_{2}CO_{3}; and a compound, KH_{3}P_{2}O_{8}, corresponding to the doubled molecule of H_{3}PO_{4}, by the action of KH_{2}PO_{4} upon H_{3}PO_{4} itself.

Pyrophosphoric acid, H_{4}P_{2}O_{7}, is formed by heating orthophosphoric acid to 250° when it loses water.[19 bis] Its normal salts are formed by igniting the dimetallic salts of orthophosphoric acid of the types HM_{2}PO_{4}. Thus from the disodium salt we obtain sodium pyrophosphate, Na_{4}P_{2}O_{7} (it crystallises from water with 10H_{2}O, is very stable, fuses when heated, has an alkaline reaction, and does not form ortho-salts when its solution is boiled): and from the monosodium salt NaH_{2}PO_{4} the acid salt Na_{2}H_{2}P_{2}O_{7} (easily soluble in water) is formed; this has an acid reaction, and when ignited further gives the meta-salt.[20]

[19 bis] According to Watson (1893) the ortho-acid is partially transformed into the pyro-acid at 230°, whilst at 260° the latter begins to volatilise. At 300° the meta-acid only is formed.

[20] The method of preparation of the acid itself consists in converting the sodium salt, Na_{4}P_{2}O_{7}, by double decomposition with water and a salt of lead, into insoluble lead pyrophosphate, Pb_{2}P_{2}O_{7}, which is then suspended in water and decomposed by sulphuretted hydrogen; lead sulphide is thus precipitated, and pyrophosphoric acid remains in solution. This solution cannot be heated, or the pyro-acid will pass into the ortho-, but must be evaporated under the receiver of an air-pump. It concentrates to a syrup and crystallises, and when ignited in this form loses water, and forms metaphosphoric acid. It resembles orthophosphoric acid in many respects; its salts with the alkalis are also soluble, and the others insoluble in water but soluble in acids. When heated in solution with acid it gives orthophosphoric acid, as well as when fused with an excess of alkali.

Witt heated ammonium chloride with phosphoric acid (hydrochloric acid was evolved), ignited the residue to drive off ammonia, and obtained pyrophosphoric acid in the residue.

Metaphosphoric acid, HPO_{3} (the analogue of nitric acid), is formed by the ignition of the pyro- and ortho-acids (or, better, of their ammonium salts), as a vitreous, hygroscopic, fused mass (glacial phosphoric acid, acidum phosphoricum glaciale), soluble in water and volatilising without decomposition. It is also formed in the first slow action of cold water on the anhydride, but metaphosphoric acid gradually changes into the ortho-acid when its solution is boiled, or when it is kept for any length of time, especially in the presence of acids.[21]

[21] As when using phenolphthalein as an indicator in neutralising by an alkali metaphosphoric acid is monobasic, and orthophosphoric acid is bibasic, it is possible by means of this difference to follow the transition of meta- into orthophosphoric acid. Sabatier (1888) carried on an investigation of this nature, and found that the rate of transformation is dependent on the temperature, and is subject to the general laws of the rate of chemical transformations which belongs to physical chemistry.

Metaphosphoric acid has a particular interest in respect to the variations to which its salts are subject. The metaphosphates are formed by the ignition of the acid orthophosphates, MH_{2}PO_{4}, or MNH_{4}HPO_{4}, or of the acid pyrophosphates, M_{2}H_{2}P_{2}O_{7}, or M_{2}(NH_{4}){2}P{2}O_{7}, water and ammonia being given off in the process. The properties of the metaphosphates, which have a similar composition to nitrates--for instance, NaPO_{3}, or Ba(PO_{3}){2}--vary according to the duration of the ignition to which the ortho-, or pyrophosphates from which they are prepared have been subjected. When the salts NaH{2}PO_{4} or NH_{4}NaHPO_{4} are strongly ignited, a salt NaPO_{3} is formed, which deliquesces in the air, and gives a gelatinous precipitate with salts of the alkaline earths. But, as Graham (in 1830-40), and many others, especially Fleitmann and Henneberg (in 1840-50), and Tamman (in the nineties), observed, under other conditions the salts of the same composition acquire other properties. The above chemists recognise five polymeric forms of metaphosphates, (HPO_{3})_{n}. We will follow the nomenclature and researches of Fleitmann.

Monometaphosphoric acid. The salts are distinguished for their insolubility in water; even the salts NaPO_{3}, KPO_{3}, are insoluble. They are obtained by igniting the monometallic orthophosphates--for example, RH_{2}PO_{4}--up to the temperature at which all water is evolved (316°), but not to fusion. No double salts are known.

Dimetaphosphoric acid, on the contrary, easily forms double salts--for example, KNaP_{2}O_{6}, and also the copper potassium salt, &c. The copper salt is obtained by evaporating a solution of copper oxide in orthophosphoric acid. A blue ortho-salt, CuRHO_{4}, first separates from the solution, then a light-blue pyro-salt, Cu_{2}P_{2}O_{7}; and above 350°, when metaphosphoric acid itself begins to volatilise, the dimetaphosphate, CuP_{2}O_{6}, is formed. The residue is washed with water, and decomposed with a hot solution of sodium sulphide, when the sodium salt, Na_{2}P_{2}O_{6}, is obtained in solution. This salt, when evaporated with alcohol, gives crystals containing 2 mol. H_{2}O, which, however, retain their solubility (in 7 parts of water) after the water is driven off at 100°. When fused, these crystals give a deliquescent salt (hexa-metaphosphate). The solution of the salt has a neutral reaction, which only after prolonged boiling becomes acid, owing to the formation of orthophosphate, NaH_{2}PO_{4}. The soluble salts of dimetaphosphoric acid give the insoluble silver salt, Ag_{2}P_{2}O_{6}, with silver nitrate, and a precipitate of BaP_{2}O_{6}2H_{2}O with barium chloride.

Trimetaphosphoric acid is obtained as the sodium salt Na_{3}P_{3}O_{9} when any other metaphosphate of sodium is fused and slowly cooled, then dissolved in a slight excess of warm water, and the resultant solution evaporated. The crystals contain 6 mol. H_{2}O, and dissolve in four parts of water. An acid reaction is only obtained, as with the preceding salt, after prolonged boiling with water. The acid is a true analogue of nitric acid, because all its metallic salts are soluble.

Hexametaphosphoric acid. Fleitmann so named the ordinary metaphosphoric acid (glacial) which attracts moisture. The deliquescent sodium salt is obtained, like the trimetaphosphate, only by rapid cooling. It is also formed by fusing silver oxide with an excess of phosphoric acid. The sodium salt is soluble in water, and gives viscous, elastic precipitates with salts of Ba, Ca, and Mg. Lubert (1893) obtained salts of Ag, Pb, &c.

Jawein and Thillot (1889), who investigated the sodium salts of metaphosphoric acid by Raoult's method, came to the conclusion that the salts of di- and tri-metaphosphoric acid behave in such a manner that their molecule must be represented as non-polymerised NaPO_{3}, whilst those of hexametaphosphoric acid behave as (NaPO_{3})_{4}. At all events, the series of salts which Fleitmann and Henneberg regard as monometaphosphates--i.e. as non-polymerised--are most probably the most polymerised, because they are insoluble.

According to Tamman's researches, vitreous metaphosphoric acid contains a mixture consisting chiefly of two varieties, differing in the solubility and degree of stability of their salts. The least stable corresponds to Fleitmann's hexa-acid, and gives three isomeric salts. Tamman came to the conclusion that there exist polymers also in the form of penta-, ortho-, and deca-metaphosphoric acids. Without going into details upon this subject, I do not think it superfluous to point out that the undoubted capability of metaphosphoric acid to polymerise should be connected with its faculty of combining with water, whilst the degree of polymerisation and the number of polymeric forms cannot yet be considered as sufficiently explained.

In order to see the relation between phosphoric acid and the lower acids of phosphorus, it is simplest to imagine the substitution of hydroxyl in H_{3}PO_{4} or PO(OH){3} by hydrogen. Then from orthophosphoric acid, PO(OH){3}, we shall obtain phosphorous acid, POH(OH){2}, and hypophosphorous acid, POH(OH); and, furthermore, phosphorous acid should be bibasic if orthophosphoric acid was tribasic, and hypophosphorous acid should be monobasic. This conclusion[21 bis] is, in fact, true, and hence all the acids of phosphorus may be referred to one common type, PX{5}, whose representatives are PH_{4}I and PCl_{5}, POCl_{3}, PCl_{2}F_{3}, &c.

[21 bis] The bibasity of H_{3}PO_{3}, established by Würtz, has been proved by many direct experiments (see, for instance, Note 22), among which we may mention that Amat (1892) took a mixture of the aqueous solutions of Na_{2}HPO_{3} and NaHO and added absolute alcohol to it. Two layers were formed; the upper, alcoholic, contained all the excess of NaHO, whilst the lower only contained the salt Na_{2}HPO_{3}, which was therefore unable to react with the excess of NaHO. Amat also obtained NaH_{2}PO_{3} by saturating H_{3}PO_{3} with soda until he obtained a neutral reaction with methyl-orange. The replacement of one atom of H by sodium here, as in phosphoric acid (Note 16), gives more heat than the replacement of the second atom. For the third atom there is no formation of a salt, and therefore no evolution of heat. The monometallic salts--for example, NaH_{2}PO_{3}--or the ammonia salts, when heated to 160°, give, as Amat had previously shown, a salt of bibasic pyrophosphorous acid, Na_{2}H_{2}P_{2}O_{5}.

Phosphorous acid, PH_{3}O_{3}, is generally obtained from phosphorus trichloride, PCl_{3}, by the action of water: PCl_{3} + 3H_{2}O = 3HCl + PH_{3}O_{3}. Both acids formed are soluble in water, but are easily separated, because hydrochloric acid is volatile whilst phosphorous acid volatilises with difficulty, and if a small amount of water be originally taken the hydrochloric acid nearly all passes off directly. Concentrated solutions of phosphorous acid give crystals of H_{3}PO_{3}, which fuse at 70°, attract moisture from the air, and deliquesce when ignited, giving phosphine and phosphoric acid,[22] and are oxidised into orthophosphoric acid by many oxidising agents. In its salts only two hydrogen atoms are replaced by metals (Würtz); the salts of the alkaline metals are soluble, and give precipitates with salts of the majority of other metals.

[22] Phosphorous acid, when subjected to the action of nascent hydrogen (zinc and sulphuric acid), evolves phosphine, and when boiled with an excess of alkali it evolves hydrogen (PH_{3}O_{3} + 3KHO = PK_{3}O_{4} + 2H_{2}O + H_{2}); owing to its liability to oxidation, it is a reducing agent--for instance, it reduces cupric chloride to cuprous chloride, and precipitates silver from the nitrate and mercury from its salts.

These reactions are perhaps connected with the fact that in this acid one atom of hydrogen should be considered as in the same condition as in phosphuretted hydrogen, which is expressed by the formula PHO(OH){2}, if we represent it as PH{4}X, with the substitution of two of the hydrogen atoms by oxygen and of HX by two of hydroxyl. The direct passage of phosphorous chloride into phosphorous acid would, however, indicate that all the three atoms of hydrogen in it occur in the form of hydroxyl, because no difference is known between the three atoms of chlorine in PCl_{3}--they all react alike, as a rule. However, Menschutkin, by acting on alcohol, C_{2}H_{5}OH, with phosphorous chloride, obtained hydrochloric acid and a substance P(C_{2}H_{5}O)Cl_{2}, and from it by the action of bromine he obtained ethyl bromide, C_{2}H_{5}Br, and a compound PBrOCl_{2}, which proves, to a certain extent, the existence of a difference between the three atoms of chlorine in phosphorous chloride. If we turn our attention to the formation of phosphine by the ignition of phosphorous acid, we see that 4PH_{3}O_{3} only evolve 3H in the form of PH_{3}, and therefore the residue--that is, 3PH_{3}O_{4}--will still contain one hydrogen of the same nature as in phosphine, because in 4PH_{3}O_{3} we should recognise four such hydrogens as in phosphine. We arrive at the same conclusion by examining the decomposition of hypophosphorous acid, 2PH_{3}O_{2} = PH_{3} + PH_{3}O_{4}. In the two molecules of the monobasic hypophosphorous acid taken, there are only two atoms of hydrogen replaceable by metals, whilst in the molecule of the resultant phosphoric acid there are three. Perhaps relations of this nature determine the relative stability of the dimetallic salts of orthophosphoric acid.

The monobasic hypophosphorous acid, PH_{3}O_{2}, gives salts PH_{2}O_{2}Na, (PH_{2}O_{2}){2}Ba, &c.; the two remaining atoms of hydrogen (which exist in the same form as in phosphine, PH{3}) are not replaceable by metals, and this determines the property of these salts of evolving phosphuretted hydrogen when heated (especially with alkalis). In acting on substances liable to reduction it is this hydrogen which acts, and, for example, reduces gold and mercury from the solutions of their salts, or converts cupric into cuprous salts. In all these instances the hypophosphorous acid is converted into phosphoric acid. Under the action of zinc and sulphuric acid it gives phosphine, PH_{3}. Nevertheless, neither hypophosphorous acid nor its dry salts absorb oxygen from the air. The salts of hypophosphorous acid are more soluble than those of the preceding acids of phosphorus. Thus the sodium salt PNaH_{2}O_{2} does not give a precipitate with barium chloride, and the salts of calcium, barium, and many other metals are soluble.[23] The hypophosphites are prepared by boiling an alkali with phosphorus so long as phosphuretted hydrogen is evolved. The acid itself is obtained from barium hypophosphite (prepared in the same manner by boiling phosphorus in baryta water), by decomposing its solution with sulphuric acid. By concentration of the solution of hypophosphorous acid (it must not be heated above 130°, at which temperature it decomposes) a syrup is formed which is able to crystallise. In the solid state hypophosphorous acid fuses at +17°, and has the properties of a clearly defined acid.

[23] Calcium hypophosphite is used in medicine. According to Cavazzi, a mixture of sodium hypophosphite, NaH_{2}PO_{2}, and sodium nitrate explodes violently.

The types PX_{3} and PX_{5}, which are evident for the hydrogen and oxygen compounds of phosphorus, are most clearly seen in its halogen compounds,[24] to the consideration of which we will proceed, fixing our attention more especially on the chlorine compounds, as being the most important from the historical, theoretical, and practical point of view.

[24] Fluorine and bromine give PX_{3} and PX_{5}, like chlorine. With respect to iodine PI_{5} is, in a chemical sense, a very unstable substance, and generally phosphorus tri-iodide only is formed (from yellow or red phosphorus and iodine in the requisite proportions). It is a red crystalline substance, fuses at 55°, is easily decomposed by water, forming phosphorous and hydriodic acids, and when heated it evolves iodine vapours and forms phosphorus di-iodide, PI_{2}. This substance may be obtained in the same manner as the preceding by taking a smaller proportion of iodine (8 parts of iodine to 1 part of phosphorus, whilst the tri-iodide requires 12·3); it also forms red crystals, which melt at 110°. When decomposed by water it not only gives phosphorous and hydriodic acids, but also phosphine and a yellow substance (a lower oxide of phosphorus). In its composition di-iodide of phosphorus corresponds with liquid phosphuretted hydrogen, PH_{2}, and probably its molecular weight is much higher: P_{2}I_{4} or P_{3}I_{6}, &c. As the iodine compounds of phosphorus give hydriodic and phosphorous acids with water, and as both these substances are reducing agents in the presence of water (and hydrates), iodide of phosphorus also acts as a reducing agent.

Phosphorus burns in chlorine, forming phosphorous chloride, PCl_{3}, and with an excess of chlorine, phosphoric chloride, PCl_{5}. The oxychloride, POCl_{3}, as the simplest chloranhydride according to the type PX_{5}, and also phosphoric chloride, correspond with orthophosphoric acid, PO(OH){3}, while phosphorous chloride, PCl{3}, corresponds with phosphorous acid and the type PX_{3}. Phosphoric oxychloride, POCl_{3}, is a colourless liquid, boiling at 110°. Phosphorus trichloride is also a colourless liquid, boiling at 76°,[25] whilst phosphoric chloride is a solid yellowish substance, which volatilises without melting at about 168°. They are all heavier than water, and form types of the chloranhydrides or chlorine compounds of the non-metallic elements whose hydrates are acids, just as NaCl or BaCl_{2} are types of halogen metallic salts.

[25] In a liquid state the density of phosphorous chloride at 10° = 1·597, and therefore its molecular volume = 137·5/1·597 = 86·0, and that of phosphorus oxychloride is equal to 153·5/1·693 = 90·7; hence the addition of oxygen has produced considerable increase in volume, just as in the conversion of sulphur dichloride, SCl_{2}, into sulphuryl chloride, SOCl_{2}, the volume changes from 64 to 71. It is the same with the boiling-points; phosphorus trichloride boils at 70°, the oxychloride at 100°, sulphur dichloride at 64°, and sulphuryl chloride at 78°--that is, the addition of oxygen raises the boiling points.

The vapour density of phosphorus trichloride and oxychloride corresponds with their formulæ (Cahours, Würtz)--namely, is equal to half the molecular weight referred to hydrogen. But it is not so with phosphorus pentachloride. Cahours showed that the vapour density of phosphorus pentachloride referred to air = 3·65, to hydrogen = 52·6, whilst according to the formula PCl_{5} it should be = 104·2. Hence this formula corresponds with four, and not with two, molecules. This shows that the vapour of phosphoric chloride contains two and not one molecule, that in a state of vapour it splits up, like sal-ammoniac, sulphuric acid, &c. The products of disruption must here be phosphorous chloride, PCl_{3}, and chlorine, Cl_{2}, bodies which easily re-form phosphoric chloride, PCl_{5}, at a lower temperature. This decomposition of phosphoric chloride in its conversion into vapour is confirmed by the fact that the vapour of this almost colourless substance shows the greenish-yellow colour proper to chlorine. This dissociation of phosphoric chloride has been considered by some chemists as a sign that phosphorus, like nitrogen, does not give volatile compounds of the type PX_{5}, and that such substances are only obtained as unstable molecular compounds which break up when distilled; for example, PH_{3},HI, PCl_{3},Cl_{2}, NH_{3},HCl, &c. To prove that the molecule PCl_{5} actually exists, Würtz in 1870 observed that when mixed with the vapour of phosphorous chloride the vapour of phosphoric chloride distils over (from 160° to 190°) perfectly colourless, and has a density which is really near to the formula--namely, to 104--and the same density was determined for the pentachloride in an atmosphere of chlorine. Hence at low temperatures and in admixture with one of the products of dissociation, there is no longer that decomposition which occurs at higher temperatures--that is, we have here a case of dissociation proceeding at moderate temperatures.

An important proof in favour of the type PX_{5} is exhibited by phosphorus pentafluoride PF_{5}, obtained by Thorpe as a colourless gas which only corrodes glass after the lapse of time; it may be kept over mercury, and has a normal density. It is formed when liquid arsenic trifluoride, AsF_{3}, is added to phosphoric chloride surrounded by a freezing mixture: 3PCl_{5} + 5AsF_{3} = 3PF_{5} + 5AsCl_{3}.

In general, fluorine and phosphorus give stable compounds: PF_{3}, POF_{3}, and PF_{5}, as would be expected from the fact that in passing from Cl to I (i.e. as the atomic weight of the halogen increases) the stability of the compounds with P and the tendency to give PX_{5} (Note 24) decreases. Phosphorus trifluoride is obtained by heating a mixture of ZnF_{2} and PBr_{3}, by the action of AsF_{3} upon PCl_{3}, by heating phosphide of copper with PbF_{2}, &c. It is a strong-smelling gas, which liquefies at -10° under a pressure of 40 atmospheres, giving a colourless liquid. It dissolves easily in (is absorbed by, reacts with) water, and acts upon glass; when mixed with Cl_{2} it combines with it (Poulenc, 1891), forming PCl_{2}F_{3}, a colourless gas of normal density, which is transformed into a liquid at 8°, decomposes into PF_{3} + Cl_{2} at 250°, and, with a small amount of water, gives oxy-fluoride of phosphorus, POF_{3} (with a large amount of water it gives PH_{3}O_{4}), which Moissan (1891) obtained by the action of dry HF upon P_{2}O_{5}, and Thorpe and Tutton (1890) by heating a mixture of cryolite and P_{2}O_{5}. It is a gas of normal density, like PF_{3}, and was obtained by Moissan by the action of fluorine upon PF_{3} (PSF_{3}, see Chapter XX., Note 20). Thus the forms PX_{3} and PX_{5} not only exist in many solid and non-volatile substances, but also as vapours.

If a piece of phosphorus be dropped into a flask containing chlorine, it burns when touched with a red-hot wire, and combines with the chlorine. If the phosphorus be in excess, liquid phosphorus trichloride, PCl_{3}, is always formed, but if the chlorine be in excess the solid pentachloride is obtained. The trichloride is generally prepared in the following manner. Dry chlorine (passed through a series of Woulfe's bottles containing sulphuric acid) is led into a retort containing sand and phosphorus. The retort is heated, the phosphorus melts, spreads through the sand, and gradually forms the trichloride, which distils over into a receiver, where it condenses. Phosphoric chloride or phosphorus pentachloride, PCl_{5}, is prepared by passing dry chlorine into a vessel containing phosphorus trichloride (purified by distillation). Phosphorous chloride combines directly with oxygen, but more rapidly with ozone or with the oxygen of potassium chlorate (3PCl_{3} + KClO_{3} = 3POCl_{3} + KCl), forming phosphorus oxychloride, POCl_{3} (Brodie). This compound is also formed by the first action of water on phosphoric chloride; for example, if two vessels, one containing phosphoric chloride and the other water, are placed under a bell jar, after a certain time the crystals of the chloride disappear and hydrochloric acid passes into the water. The aqueous vapour acts on the pentachloride, and the following reaction occurs: PCl_{5} + H_{2}O = POCl_{3} + 2HCl, the result being that liquid phosphorus oxychloride is found in one vessel, and a solution of hydrochloric acid in the other. However, an excess of water directly transforms phosphoric chloride into orthophosphoric acid, PCl_{5} + 4H_{2}O = PH_{3}O_{4} + 5HCl,[26] since POCl_{3} reacts with water (3H_{2}O), forming 3HCl and phosphoric acid PO(OH)_{3}.

[26] Phosphorus oxychloride is obtained by the action of phosphoric chloride on hydrates of acids (because alkalis decompose phosphorus oxychloride), according to the equation PCl_{5} + RHO = POCl_{3} + RCl + HCl, where RHO is an acid. The reaction only proceeds according to this equation with monobasic acids, but then RCl is volatile, and therefore a mixture is obtained of two volatile substances, the acid chloride and phosphorus oxychloride, which are sometimes difficult to separate; whilst if the hydrate be polybasic the reaction frequently proceeds so that an anhydride is formed: RH_{2}O_{2} + PCl_{5} = RO + POCl_{3} + 2HCl. If the anhydride be non-volatile (like boric), or easily decomposed (like oxalic), it is easy to obtain pure oxychloride. Thus phosphorus oxychloride is often prepared by acting on boric or oxalic acid with phosphoric chloride. It is also formed when the vapour of phosphoric chloride is passed over phosphoric anhydride, P_{2}O_{5} + 3PCl_{5} = 5POCl_{3}. This forms an excellent example in proof of the fact that the formation of one substance from two does not necessarily show that the resultant compound contains the molecules of these substances in its molecule. But other oxychlorides of phosphorus are also formed by the interaction of phosphoric anhydride and chloride; thus at 200° the chloranhydride, PO_{2}Cl, or chloranhydride of metaphosphoric acid, is formed (Gustavson). The chloranhydride of pyrophosphoric acid, P_{2}O_{3}Cl_{4}, was obtained (Hayter and Michaelis), together with NOCl, &c., by the action of NO upon cold PCl_{3}, as a fuming liquid boiling at 210°.

The above chlorine compounds serve not only as a type of the chloranhydrides, but also as a means for the preparation of other acid chloranhydrides. Thus the conversion of acids XHO into chloranhydrides, XCl, is generally accomplished by means of phosphorus pentachloride. This fact was discovered by Chancel, and adopted by Gerhardt as an important method for studying organic acids. By this means organic acids, containing, as we know, RCOOH (where R is a hydrocarbon group, and where carboxyl may repeat itself several times by replacing the hydrogen of hydrocarbon compounds), are converted into their chloranhydrides, RCOCl. With water they again form the acid, and resemble the chloranhydrides of mineral acids in their general properties.

Since carbonic acid, CO(OH){2}, contains two hydroxyl groups, its perfect chloranhydride, COCl{2}, carbonic oxychloride, carbonyl chloride or phosgene gas, contains two atoms of chlorine, and differs from the chloranhydrides of organic acids in that in them one atom of chlorine is replaced by the hydrocarbon radicle RCOCl, if R be a monatomic radicle giving a hydrocarbon RH. It is evident, on the one hand, that in RCOCl the hydrogen is replaced by the radicle COCl, which is also able to replace several atoms of hydrogen (for example, C_{2}H_{4}(COCl){2} corresponds with the bibasic succinic acid); and, on the other hand, that the reactions of the chloranhydrides of organic acids will answer to the reactions of carbonyl chloride, as the reactions of the acids themselves answer to those of carbonic acid. Carbonyl chloride is obtained directly from dry carbon monoxide and chlorine[27] exposed to the action of light, and forms a colourless gas, which easily condenses into a liquid, boiling at +8°, specific gravity 1·43, and having the suffocating odour belonging to all chloranhydrides. Like all chloranhydrides, it is immediately decomposed by water, forming carbonic anhydride, according to the equation COCl{2} + H_{2}O = CO_{2} + 2HCl, and thus expresses the type proper to all chloranhydrides of both mineral and organic acids.[28]

[27] The direct action of the sun's rays, or of magnesium light, is necessary to start the reaction between carbonic oxide and chlorine, but when once started it will proceed rapidly in diffused light. An excess of chlorine (which gives its coloration to the colourless phosgene) aids the completion of the reaction, and may afterwards be removed by metallic antimony. Porous substances, like charcoal, aid the reaction. Phosgene may be prepared by passing a mixture of carbonic anhydride and chlorine over incandescent charcoal. Lead or silver chloride, when heated in a current of carbonic oxide, also partially form phosgene gas. Carbon tetrachloride, CCl_{4}, also forms it when heated with carbonic anhydride (at 400°), with phosphoric anhydride (200°), and most easily of all with sulphuric anhydride (2SO_{3} + CCl_{4} = COCl_{2} + S_{2}O_{5}Cl_{2}, this is pyrosulphuryl chloride). Chloroform, CHCl_{3}, is converted into carbonyl chloride when heated with SO_{2}(OH)Cl (the first chloranhydride of sulphuric acid); CHCl_{3} + SO_{3}HCl = COCl_{2} + SO_{2} + 2HCl (Dewar), and when oxidised by chromic acid.

Among the reactions of phosgene we may mention the formation of urea with ammonia, and of carbonic oxide when heated with metals.

[28] We are already acquainted with some of the chloranhydrides of the inorganic acids--for instance, BCl_{3}, and SiCl_{4}--and here we shall describe those which correspond with sulphuric acid in the following chapter. It may be mentioned here that when hydrochloric acts on nitric acid (aqua regia, Vol. I. p. 467) there is formed, besides chlorine, the oxychlorides NOCl and NO_{2}Cl, which may be regarded as chloranhydrides of nitric and nitrous acids (nitrogen chloride, Vol. I. p. 476). The former boils at -5°, the latter at +5°, the specific gravity of the first at -12° = 1·416, and at -18° = 1·433 (Geuther), and of the second = 1·3; the first is obtained from nitric oxide and chlorine, the second from nitric peroxide and chlorine, and also by the action of phosphoric chloride on nitric acid. If the gases evolved by aqua regia be passed into cold and strong sulphuric acid, they form crystals of the composition NHSO_{3} (like chamber crystals), which melt at 86°, and with sodium chloride form acid sodium sulphate and the oxychloride NOCl. This chloranhydride of nitric acid is termed nitrosyl chloride.

Cyanogen chloride, CNCl, is the gaseous chloranhydride of cyanic acid; it is formed by the action of chlorine on aqueous mercury cyanide, Hg(CN){2} + 2Cl{2} = HgCl_{2} + 2CNCl. When chlorine acts on cyanic acid, it forms not only this cyanogen chloride, but also polymerides of it--a liquid, boiling at 18°, and a solid, boiling at 190°. The latter corresponds with cyanuric acid, and consequently contains C_{3}N_{3}Cl_{3}. Details concerning these substances must be looked for in works on organic chemistry.

In order to show the general method for the preparation of acid chloranhydrides, we will take that of acetic acid, CH_{3}·COOH, as an example. Phosphorus pentachloride is placed in a glass retort, and acetic acid poured over it; hydrochloric acid is then evolved, and the substance distilling over directly after is a very volatile liquid, boiling at 50°, and having all the properties of the chloranhydrides. With water it forms hydrochloric and acetic acids. The reaction here taking place may be explained thus: the substitution of the oxygen taken from the acetic acid (from its carboxyl) by two atoms of chlorine from the PCl_{5} should be as follows: CH_{3}·COOH + PCl_{5} = CH_{3}·COHCl_{2} + POCl_{3}. But the compound CH_{3}·COHCl_{2} does not exist in a free state (because it would indicate the possibility of the formation of compounds of the type CX_{6}, and carbon only gives those of the type CX_{4}); it therefore splits up into HCl and the chloranhydride CH_{3}·COCl. The general scheme for the reaction of phosphorus pentachloride with hydrates ROH is exactly the same as with water; namely, ROH with PCl_{5}, gives POCl_{3} + HCl + RCl--that is a chloranhydride.[28 bis]

[28 bis] This reaction indeed proceeds very easily and completely with a number of hydroxides, if they do not react on hydrochloric acid and phosphorus oxychloride, which is the case when they have alkaline properties. When the hydroxide is bibasic and is present in excess, it not unfrequently happens that the elements of water are taken up: R(OH){2} + PCl{5} = RO + 2HCl + POCl_{3}. The anhydride RO may then be converted into chloranhydride, RO + PCl_{5} = RCl_{2} + POCl_{3}--that is, phosphorus pentachloride brings about the substitution of O by Cl_{2}. Thus carbonyl chloride, COCl_{2}, boron chloride, 2BCl_{3}, and succinic chloride, C_{4}H_{4}O_{2}Cl_{2}, &c., are respectively obtained by the action of phosphoric chloride on carbonic, boric, and succinic anhydrides. Phosphorus pentachloride reacts in a similar manner on the aldehydes, RCHO, forming RCHCl_{2}, and on the chloranhydrides themselves--for example, with acetic chloride, CH_{3}.COCl (when heated in a closed tube), it forms a substance having the composition CH_{3}CCl_{3}.

Phosphorus trichloride and oxychloride act in a similar manner to phosphoric chloride. When phosphorus trichloride acts on an acid, 3RHO + PCl_{3} = 3RCl + P(HO){3}. If a salt is taken, then by the action of phosphorus oxychloride a corresponding chloranhydride and salt of orthophosphoric acid are easily formed: 3R(KO) + POCl{3} = 3RCl + PO(KO){3}. The chloranhydride RCl is always more volatile than its corresponding acid, and distils over before the hydrate RHO. Thus acetic acid boils at 117°, and its chloranhydride at 50°. Phosphoric and phosphorous acids are very slightly volatile, whilst their chloranhydrides are comparatively easily converted into vapour. The faculty of the chloranhydrides to react at the expense of their own chlorine determines their great importance in chemistry. For instance, suppose we require to know the molecular formula of some hydrate which does not pass into a state of vapour and does not give a chloranhydride with hydrochloric acid--that is, which has not any basic or alkaline properties; we must then endeavour to obtain this chloranhydride by means of phosphoric chloride, and it frequently happens that the corresponding chloranhydride is volatile. The resultant chloranhydride is then converted into vapour, and its composition is determined; and if we know its composition we are able to decide that of its corresponding hydrate. Thus, for example, from the formula of silicon chloride, SiCl{4}, or of boron chloride, BCl_{3}, we can judge the composition of their corresponding hydrates, Si(HO){4}, B(HO){3}. Having obtained the chloranhydride RCl or RCl_{n}, it is possible by its means to obtain many other compounds of the same radicle R according to the equation MX + RCl = MCl + RX. M may be = H, K, Ag, or other metal. The reaction proceeds thus if M forms a stable compound with chlorine--for example, silver chloride, hydrochloric acid, and R, an unstable substance. Hence, a chloranhydride is frequently employed for the formation of other compounds of a given radicle; for instance, with ammonia they form amides RNH_{2}, and with salts ROK, with anhydrides R_{2}O, &c.

Containing, as they do, chlorine, which easily reacts with hydrogen, phosphorus pentachloride, trichloride, and oxychloride enter into reaction with ammonia, and give a series of amide and nitrile compounds of phosphorus. Thus, for example, when ammonia acts on the oxychloride we obtain sal-ammoniac (which is afterwards removed by water) and an orthophosphoric triamide, PO(NH_{2}){3}, as a white insoluble powder on which dilute acids and alkalis do not act, but which, when fused with potassium hydroxide, gives potassium phosphate and ammonia like other amides. When ignited, the triamide liberates ammonia and forms the nitrile PON, just as urea, CO(NH{2})_{2}, gives off ammonia and forms the nitrile CONH. This nitrile, called monophosphamide, PON, naturally corresponds with metaphosphoric acid, namely, with its ammonium salt. NH_{4}PO_{3} - H_{2}O = PO_{2}·NH_{2}, an as yet unknown amide, and PO_{2}·NH_{2} - H_{2}O gives the nitrile PON. This relation is confirmed by the fact that PON, moistened with water, gives metaphosphoric acid when ignited. It is the analogue of nitrous oxide, NON. It is a very stable compound, more so than the preceding.[29]

[29] The reaction of ammonia on phosphorus pentachloride is more complex than the preceding. This is readily understood: to the oxychloride, POCl_{3}, tere corresponds a hydrate PO(OH){3}, and a salt PO(NH{4}O){3}, and consequently also an amide PO(NH{2}){3}, whilst the pentachloride, PCl{5}, has no corresponding hydrate P(OH){5}, and therefore there is no amide P(NH{2}){5}. The reaction with ammonia will be of two kinds: either instead of 5 mol. NH{3}, only 3 mol. NH_{3} or still less will act; i.e. PCl_{2}(NH_{2}){3}, PCl{3}(NH_{2}){2}, &c. are formed; or else the pentachloride will act like a mixture of chlorine with the trichloride, and then as the result there will be obtained the products of the action of chlorine on those amides which are formed from phosphorus trichloride and ammonia. It would appear that both kinds of reaction proceed simultaneously, but both kinds of products are unstable, at all events complex, and in the result there is obtained a mixture containing sal-ammoniac, &c. The products of the first kind should react with water, and we should obtain, for example, PCl{3}(NH_{2}){2} + 2H{2}O = 3HCl and PO(HO)(NH_{2}){2}. This substance has not actually been obtained, but the compound PONH(NH{2}) derived from it by elimination of the elements of water is known, and is termed diphosphamide; it is, however, more probable that it is a nitrile than an amide, because only amides contain the group NH_{2}. It is a colourless, stable, insoluble powder, which possibly corresponds with pyrophosphoric acid, more especially since when heated it evolves ammonia and gives and leaves phosphoryl nitride, PON--that is, the nitrile of metaphosphoric acid. The amide corresponding with the pyrophosphate P_{2}O_{3}(NH_{4}O){4} should be P{2}O_{3}(NH_{2}){4}, and the nitriles corresponding to the latter would be P{2}O_{2}N(NH_{2}){3}, P{2}ON_{2}(NH_{2}){2}, and P{2}N_{3}(NH_{2}). The composition of the first is the same as that of the above diphosphamide. The third pyrophosphoric nitrile has a formula P_{2}N_{4}H_{2}, and this is the composition of the body known as phospham, PHN_{2} (in a certain sense this is the analogue of N_{3}H polymerised, Chapter VI.) Indeed, phospham has been obtained by heating the products of the action of ammonia on phosphoric chloride, as an insoluble and alkaline powder, which gives ammonia and phosphoric acid when subjected to the action of water. The same substance is obtained by the action of ammonium chloride on phosphoric chloride (PNCl_{2} is first formed, and reacts further with ammonia, forming phospham), and by igniting the mass which is formed by the action of ammonia on phosphorus trichloride. Formerly the composition of phospham was supposed to be PHN_{2}, now there is reason to think that its molecular weight is P_{3}H_{3}N_{6}.

The above compounds correspond with normal salts, but nitriles and amides corresponding to acid salts are also possible, and they will be acids. For example, the amide PO(HO){2}(NH{2}), and its nitrile, will be either PN(HO){2} or PO(HO)(NH), but at all events of the composition PNH{2}O_{2}, and having acid properties. The ammonium salt of this phosphonitrilic acid (it is called phosphamic acid), PNH(NH_{4})O_{2}, is obtained by the action of ammonia on phosphoric anhydride, P_{2}O_{5} + 4NH_3 = H_{2}O + 2PNH(NH_{4})O_{2}. A non-crystalline soluble mass is thus formed, which is dissolved in a dilute solution of ammonia and precipitated with barium chloride, and the resultant barium salt is then decomposed with sulphuric acid, and thus a solution of the acid of the above composition is obtained.

It is evident from the theory of the formation of amides and nitriles (Chapter IX.) that very many compounds of this kind can correspond with the acids of phosphorus; but as yet only a few are known. The easy transitions of the ortho-, meta-, and pyrophosphoric acids, by means of the hydrogen of ammonia, into the lower acids, and conversely, tend to complicate the study of this very large class of compounds, and it is rarely that the nature of a product thus obtained can be judged from its composition; and this all the more that instances of isomerism and polymerism, of mixture between water of crystallisation and of constitution, &c., are here possible. Many data are yet needed to enable us to form a true judgment as to the composition and structure of such compounds. As the best proof of this we will describe the very interesting and most fully investigated compound of this class, PNCl_{2}, called chlorophosphamide, or nitrogen chlorophosphorite. It is formed in small quantities when the vapour of phosphoric chloride is passed over ignited sal-ammoniac. Besson (1892) heated the compound PCl_{5}8NH_{3} (which is easily and directly formed from PCl_{5} and NH_{3}) under a pressure of about 50 mm. (of mercury) to 200°, and obtained brilliant crystals of PNCl_{2}, which melted at 106° (in the residue after the distillation of sal-ammoniacal phospham). The chlorine in it is very stable--quite different from that in phosphoric chloride. Indeed, the resultant substance is not only insoluble in water (though soluble in alcohol and ether), but it is not even moistened by it, and distils over, together with steam, without being decomposed. In a free state it easily crystallises in colourless prisms, fuses at 114°, boils at 250° (Gladstone, Wichelhaus), and when fused with potash gives potassium chloride and the amidonitrile of phosphoric acid. Judging from its formula and the simplicity of its composition and reactions, it might be thought that the molecular weight of this substance would be expressed by the formula PCl_{2}N, that it corresponds with PON and with PCl_{5} (like POCl_3), with the substitution of Cl_3 by N, just as in POCl_3 two atoms of chlorine are replaced by oxygen; but all these surmises are incorrect, because its vapour density (referred to hydrogen--Gladstone, Wichelhaus) = 182--that is, the molecular formula must be three times greater, P_{3}N_{3}Cl_{6}. The polymerisation (tripling) is here of exactly the same kind as with the nitriles.

The most important analogue of phosphorus is arsenic, the metallic aspect of which and the general character of its compounds of the types AsX_{3} and AsX_{5} at once recall the metals. The hydrate of its highest oxide, arsenic acid (ortho-arsenic acid), H_{3}AsO_{4}, is an oxidising agent, and gives up a portion of its oxygen to many other substances; but, nevertheless, it is very like phosphoric acid. Mitscherlich established the conception of isomorphism by comparing the salts of these acids.[30]

[30] It is necessary to remark that, although arsenic is so closely analogous to phosphorus (especially in the higher forms of combination, RX_{3} and RX_{5}), at the same time it exhibits a certain resemblance and even isomorphism with the corresponding compounds of sulphur (especially the metallic compounds of the type MAs, corresponding with MS). Thus compounds containing metals, arsenic, and sulphur are very frequently met with in nature. Sometimes the relative amounts of arsenic and sulphur vary, so that an isomorphous substitution between the arsenides and sulphides must be recognised. Besides FeS_{2} (ordinary pyrites), and FeAs_{2}, iron forms an arsenical pyrites containing both sulphur and arsenic, which from its composition, FeAsS or FeS_{2}FeAs_{2}, resembles the two preceding.

Arsenic occurs in nature, not only combined with metals, but also, although rarely, native and also in combination with sulphur in two minerals--one red, realgar, As_{2}S_{2}, and the other yellow, orpiment, As_{2}S_{3} (Chapter XX., Note 29). Arsenic occurs, but more rarely, in the form of salts of arsenic acid--for instance, the so-called cobalt and nickel blooms, two minerals which are found accompanying other cobalt ores, are the arsenates of these metals. Arsenic is also found in certain clays (ochres) and has been discovered in small quantities in some mineral springs, but it is in general of rarer occurrence in nature than phosphorus. Arsenic is most frequently extracted from arsenical pyrites, FeSAs, which, when roasted without access of air, evolves the vapour of arsenic, ferrous sulphide being left behind. It is also obtained by heating arsenious anhydride with charcoal, in which case carbonic oxide is evolved. In general, the oxides and other compounds are very easily reduced. Solid arsenic is a steel-grey brittle metal, having a bright lustre and scaly structure. Its specific gravity is 5·7. It is opaque and infusible, but volatilises as a yellow vapour which on cooling deposits rhombohedral crystals.[30 bis] The vapour density of arsenic is 150 times greater than that of hydrogen--that is, its molecule, like that of phosphorus, contains 4 atoms, As_{4}. When heated in the air, arsenic easily oxidises into white arsenious anhydride, As_{2}O_{3}, but even at the ordinary temperature it loses its lustre (becomes dull), owing to the formation of a coating of a lower oxide. The latter appears to be as volatile as arsenious anhydride, and it is probable that it is owing to the presence of this compound that the vapours of arsenious compounds, when heated with charcoal (for example, in the reducing flame of a blow-pipe), have the characteristic smell of garlic, because the vapour of arsenic itself has not this odour.

[30 bis] According to Retgers (1893) the arsenic mirror (see further on) is an unstable variety of metallic arsenic, whilst the brown product which is formed together with it in Marsh's apparatus is a lower hydride AsH. Schuller and McLeod (1894), however, recognise a peculiar yellow variety of arsenic.

Arsenic easily combines with bromine and chlorine;[31] nitric acid and aqua regia also oxidise it into the higher oxide, or rather its hydrate, arsenic acid.[32] As far as is known, it does not decompose steam, and it acts exceedingly slowly on those acids, like hydrochloric, which are not capable of oxidising.

[31] Hydrochloric acid dissolves arsenious anhydride in considerable quantities, and this is probably owing to the formation of unstable compounds in which the arsenious anhydride plays the part of a base. A compound called arsenious oxychloride, having the composition AsOCl, is even known. It is formed when arsenious anhydride is added little by little to boiling arsenic trichloride, As_{2}O_{3} + AsCl_{3} = 3AsOCl. It is a transparent substance, which fumes in air, and combines with water to form a crystalline mass having the composition As_{2}(OH){4}Cl{2}. When heated it decomposes into arsenious chloride and a fresh oxychloride of a more complex composition, As_{6}O_{8}Cl_{2}· Arsenic trichloride, when treated with a small quantity of water, forms the crystalline compound, As_{2}(HO){4}Cl{2}, mentioned above. These compounds resemble the basic salts of bismuth and aluminium. The existence of these compounds shows that arsenic is of a more metallic or basic character than phosphorus. Nevertheless arsenic trichloride, AsCl_{3}, resembles phosphorus trichloride in many respects. It is obtained by the direct action of chlorine on arsenic, or by distilling a mixture of common salt, sulphuric acid, and arsenious anhydride. The latter mode of preparation already indicates the basic properties of the oxide. Arsenious chloride is a colourless oily liquid, boiling at 130°, and having a sp. gr. of 2·20. It fumes in air like other chloranhydrides, but it is much more slowly and imperfectly decomposed by water than phosphorus trichloride. A considerable quantity of water is required for its complete decomposition into hydrochloric acid and arsenious anhydride. It forms an excellent example of the transition from true metallic chlorides to true chloranhydrides of the acids. It hardly combines with chlorine, i.e. if AsCl_{5} is formed it is very unstable. Arsenic tribromide, AsBr_{3}, is formed as a crystalline substance, fusing at 20° and boiling at 220°, by the direct action of metallic arsenic on a solution of bromine in carbon bisulphide, the latter being then evaporated. The specific gravity of arsenic tribromide is 3·36. Crystalline arsenic tri-iodide, AsI_{3}, having a sp. gr. 4·39, may be obtained in a like manner; it may be dissolved in water, and on evaporation separates out from the solution in an anhydrous state--that is, it is not decomposed--and consequently behaves like metallic salts. Arsenic trifluoride, AsF_{3}, is obtained by heating fluor spar and arsenious anhydride with sulphuric acid. It is a fuming, colourless, and very poisonous liquid, which boils at 63° and has a sp. gr. of 2·73. It is decomposed by water. It is very remarkable that fluorine forms a pentafluoride of arsenic also, although this compound has not yet been obtained in a separate state, but only in combination with potassium fluoride. This compound, K_{3}AsF_{8}, is formed as prismatic crystals when potassium arsenate, K_{3}AsO_{4}, is dissolved in hydrofluoric acid.

[32] Arsenic acid, H_{3}AsO_{4}, corresponding with orthophosphoric acid, is formed by oxidising arsenious anhydride with nitric acid, and evaporating the resultant solution until it attains a sp. gr. of 2·2; on cooling it separates in crystals having the above composition. This hydrate corresponds with the normal salts of arsenic acid; but on dissolving in water (without heating), and on cooling a strong solution, crystals containing a greater amount of water, namely, (AsH_{3}O_{4}){2},H{2}O, separate. This water, like water of crystallisation, is very easily expelled at 100°. At 120° crystals having a composition identical with that of pyrophosphoric acid, As_{2}H_{4}O_{7}, separate, but water, on dissolving this hydrate with the development of heat, forms a solution in no way differing from a solution of ordinary arsenic acid, so that it is not an independent pyroarsenic acid that is formed. Neither is there any true analogue of metaphosphoric acid, although the compound AsHO_{3} is formed at 200°, and on solidifying forms a mass having a pearly lustre and sparingly soluble in cold water; but on coming into contact with warm water it becomes very hot, and gives ordinary orthoarsenic acid in solution. Arsenic acid forms three series of salts, which are perfectly analogous to the three series of orthophosphates. Thus the normal salt, K_{3}AsO_{4}, is formed by fusing the other potassium arsenates with potassium carbonate; it is soluble in water and crystallises in needles which do not contain water. Di-potassium arsenate, K_{2}HAsO_{4}, is formed in solution by mixing potassium carbonate and arsenic acid until carbonic anhydride ceases to be evolved; it does not crystallise, and has an alkaline reaction; hence it corresponds perfectly with the sodium phosphate. As was mentioned above, arsenic acid itself acts as an oxidising agent; for example, it is used in the manufacture of aniline dyes for oxidising the aniline, and it is prepared in large quantities for this purpose. When sulphuretted hydrogen is passed through its solution, sulphuric acid and arsenious anhydride are obtained in solution. Arsenic acid is very easily soluble in water, and its solution has an exceedingly acid reaction, and when boiled with hydrochloric acid evolves chlorine, like selenic, chromic, manganic, and certain other higher metallic acids.

Arsenic anhydride, As_{2}O_{5}, is produced when arsenic acid is heated to redness. It must be carefully heated, as at a bright red heat it decomposes into oxygen and arsenious anhydride. Arsenic anhydride is an amorphous substance almost entirely insoluble in water, but it attracts moisture from the air, deliquesces, and passes into the acid. Hot water produces this transformation with great ease.

Arseniuretted hydrogen, arsine, AsH_{3}, resembles phosphuretted hydrogen in many respects. This colourless gas, which liquefies into a mobile liquid at -40°, has a disagreeable garlic-like odour, is only slightly soluble in water, and is exceedingly poisonous. Even in a small quantity it causes great suffering, and if present to any considerable amount in air it even causes death. The other compounds of arsenic are also poisonous, with the exception of the insoluble sulphur compound and some compounds of arsenic acid. Arseniuretted hydrogen, AsH_{3}, is obtained by the action of water on the alloy of arsenic and sodium, sodium hydroxide and arseniuretted hydrogen being formed. It is also formed by the action of sulphuric acid on the alloy of arsenic and zinc: Zn_{3}As_{2} + 3H_{2}SO_{4} = 2AsH_{3} + 3ZnSO_{4}.[33] The oxygen compounds of arsenic are very easily reduced by the action of hydrogen at the moment of its evolution from acids, and the reduced arsenic then combines with the hydrogen; hence, if a certain amount of an oxygen compound of arsenic be put into an apparatus containing zinc and sulphuric acid (and thus serving for the evolution of hydrogen), the hydrogen evolved will contain arseniuretted hydrogen. In this case it is diluted with a considerable amount of hydrogen. But its presence in the most minute quantities may be easily recognised from the fact that it is easily decomposed by heat (200° according to Brunn) into metallic arsenic and hydrogen, and therefore if such impure hydrogen he passed through a moderately-heated tube metallic arsenic will be deposited as a bright layer on the part of the tube which was heated (see Note 30 bis). This reaction is so sensitive that it enables the most minute traces of arsenic to be discovered; hence it is employed in medical jurisprudence, as a test in poisoning cases. It is easy to discover the presence of arsenic in common zinc, copper, sulphuric and hydrochloric acids, &c. by this method. It is obvious that in testing for poison by Marsh's apparatus it is necessary to take zinc and sulphuric acid quite free from arsenic. The arsenic deposited in the tube may be driven as a volatile metal from one place to another in the current of hydrogen evolved, owing to its volatility. This forms a distinction between arseniuretted and antimoniuretted hydrogen, which is decomposed by heat in just the same way as arseniuretted hydrogen, but the mirror given by Sb is not so volatile as that formed by As.

[33] The formation of arseniuretted hydrogen is accompanied by the absorption of 37,000 heat units, while phosphine evolves 18,000 (Ogier), and ammonia 27,000. Sodium (0·6 p.c.) amalgam, with a strong solution of As_{2}O_{3}, gives a gas containing 86 vols. of arsenic and 14 vols. of hydrogen (Cavazzi).

If hydrogen contains arseniuretted hydrogen, it also gives metallic arsenic when it burns, because in the reducing flame of hydrogen the oxygen attracted combines entirely with the hydrogen and not with the arsenic, so that if a cold object, such as a piece of china, be held in the hydrogen flame the arsenic will be deposited upon it as a metallic spot.[34]

[34] This spot, or the metallic ring which is deposited on the heated tube, may easily be tested as to whether it is really due to arsenic or proceeds from some other substance reduced in the hydrogen flame--for instance, carbon or antimony. The necessity for distinguishing arsenic from antimony is all the more frequently encountered in medical jurisprudence, from the fact that preparations of antimony are very frequently used as medicine, and antimony behaves in the hydrogen apparatus just like arsenic, and therefore in making an investigation for poisoning by arsenic it is easy to mistake it for antimony. The best method to distinguish between the metallic spots of arsenic and antimony is to test them with a solution of sodium hypochlorite, free from chlorine, because this will dissolve arsenic and not antimony. Such a solution is easily obtained by the double decomposition of solutions of sodium carbonate and bleaching powder. A solution of potassium chlorate acts in the same manner, only more slowly. Further particulars must be looked for in analytical works.

Arseniuretted hydrogen, like phosphuretted hydrogen, is only slightly soluble in water, has no alkaline properties--that is, it does not combine with acids--and acts as a reducing agent. When passed into a solution of silver nitrate it gives a blackish brown precipitate of metallic silver, the arsenic being oxidised. If acting on copper sulphate and similar salts, arseniuretted hydrogen sometimes forms arsenides--i.e. it reduces the metallic salt with its hydrogen, and is itself reduced to arsenic. Sulphuric, and even hydrochloric, acid reduces arseniuretted hydrogen to arsenic, and it is still more easily decomposed by arsenious chloride, and with phosphorous chloride it gives the compound PAs. Arseniuretted hydrogen gives metallic arsenic with an acid solution of arsenious anhydride (Tivoli).

The most common compound of arsenic is the solid and volatile arsenious anhydride, As_{2}O_{3}, which corresponds with phosphorous and nitrous anhydrides. This very poisonous, colourless, and sweet-tasting substance is generally known under the name of arsenic, or white arsenic. The corresponding hydrate is as yet unknown; its solutions, when evaporated, yield crystals of arsenious anhydride. It is chiefly prepared for the dyer, and is also used as a vermin killer, and sometimes in medicine; it is a product from which all other compounds of arsenic can be prepared. It is obtained as a by-product in roasting cobalt and other ores containing arsenic. Arsenical pyrites are sometimes purposely roasted for the extraction of arsenious anhydride. When arsenical ores are burnt in the air, the sulphur and arsenic are converted into the oxides As_{2}O_{3} and SO_{2}. The former is a solid at the ordinary temperature, and the latter gaseous, and therefore the arsenious anhydride is deposited as a sublimate in the cooler portion of the flues through which the vapours escape from the furnace. It collects in condensing chambers especially constructed in the flues. The deposit is collected, and after being distilled gives arsenious anhydride in the form of a vitreous non-crystalline mass. This is one of the varieties of arsenious anhydride, which is also known in two crystalline forms. When sublimed--i.e. when it rapidly passes from the state of vapour to the solid state--it appears in the regular system in the form of octahedra.[35] It is obtained in the same form when it is crystallised from acid solutions. The specific gravity of the crystals is 3·7. The other crystalline form (in prisms) belongs to the rhombohedral system, and is also formed by sublimation when the crystals are deposited on a heated surface, or when it is crystallised from alkaline solutions.[36]

[35] According to Mitscherlich's determination, the vapour density of arsenious anhydride is 199 (H = 1)--that is, it answers to the molecular formula As_{4}O_{6}. Probably this is connected with the fact that the molecule of free arsenic contains As_{4}. V. Meyer and Biltz, however, showed (1889) that at a temperature of about 1,700° the vapour density of arsenic corresponds with the molecule As_{2}, and not As_{4}, as at lower temperatures.

[36] Arsenious anhydride is obtained in an amorphous form after prolonged heating at a temperature near to that at which it volatilises, or, better still, by heating it in a closed vessel. It then fuses to a colourless liquid, which on cooling forms a transparent vitreous mass, whose specific gravity is only slightly less than that of the crystalline anhydride. On cooling, this vitreous mass undergoes an internal change, in which it crystallises and becomes opaque, and acquires the appearance of porcelain. The following difference between the vitreous and opaque varieties is very remarkable: when the vitreous variety is dissolved in strong and hot hydrochloric acid it gives crystals of the anhydride on cooling, and this crystallisation is accompanied by the emission of light (which is visible in the dark), and the entire liquid glows as the crystals begin to separate. The opaque variety does not emit light when the crystals separate from its hydrochloric acid solution. It is also remarkable that the vitreous variety passes into the opaque form when it is pounded--that is, under the action of a series of blows. Thus, several varieties of arsenious anhydride are known, but as yet they are not characterised by any special chemical distinctions, and even differ but little in their specific gravities, so that it cannot be said that the above differences are due to any isomeric transformation--that is, to an arrangement of the atoms in the molecule--but probably only depend on a difference in the distribution of the molecules, or, in other terms, are physical and not chemical variations. One part of the vitreous anhydride requires twelve parts of boiling water for its solution, or twenty-five parts at the ordinary temperature. The opaque variety is less soluble, and at the ordinary temperature requires about seventy parts of water for its solution.

Solutions of arsenious anhydride have a sweet metallic taste, and give a feeble acid reaction. Its solubility increases with the admixture of acids and alkalis. This shows the property of arsenious anhydride of forming salts with acids and alkalis. And in fact compounds of it with hydrochloric acid (Note 31), sulphuric anhydride (see further on), and with the alkali oxides are known.[37] If silver nitrate be added to a solution of arsenious anhydride, it does not give any precipitate unless a certain amount of the arsenious anhydride is saturated with an alkali--for instance, ammonia. It then gives a precipitate of silver arsenite, Ag_{3}AsO_{3}. This is yellow, soluble in an excess of ammonia, and anhydrous; it distinctly shows that arsenious acid is tribasic, and that it differs in this respect from phosphorous acid, in which only two atoms of hydrogen can be replaced by metals.[38] The feeble acid character of arsenious anhydride is confirmed by the formation of saline compounds with acids. In this respect the most remarkable example is the anhydrous compound with sulphuric acid, having the composition As_{2}O_{3},SO_{3}. It is formed in the roasting of arsenical pyrites in those spaces where the arsenious anhydride condenses, a portion of the sulphurous anhydride being converted into sulphuric anhydride, SO_{3}, at the expense of the oxygen of the air. The compound in question forms colourless tabular crystals, which are decomposed by water with formation of sulphuric acid and arsenious anhydride.[39]

[37] Arsenious anhydride does not oxidise in air, either in a dry state or in solution, but in the presence of alkalis it absorbs oxygen from the air, and acts as an excellent reducing agent. This probably is connected with the fact that arsenic acid is much more energetic than arsenious acid, and that it is arsenic acid which is formed by the oxidation of the latter in the presence of alkalis. Arsenious anhydride is easily reduced to arsenic by many metals, even by copper.

[38] The feebleness of the acid properties of arsenious anhydride is seen in the fact that if it be dissolved in ammonia water, and then a still stronger solution of ammonia be added, prismatic crystals separate having the composition of ammonium metarsenite, NH_{4}AsO_{3}. This ammonium salt deliquesces in air, and loses all its ammonia. The magnesium salt is tri-metallic, Mg_{3}(AsO_{3})_{2}; it is insoluble in water, and is formed by mixing an ammoniacal solution of arsenious anhydride with an ammoniacal solution of a magnesium salt. It is insoluble even in ammonia, although it dissolves in an excess of acids. Magnesium hydroxide gives the same salt with arsenious solutions, and hence magnesia is one of the best antidotes for arsenic poisoning. The arsenites of copper are much used in the manufacture of colours, more especially of pigments. They are distinguished by their insolubility in water and by their remarkably vivid green colour, but at the same time by their poisonous character. Not only do such pigments applied to wall papers or other materials easily dust off from them, but they give exhalations containing AsH_{3}. The cupric salts, CuX_{2}, when mixed with an alkaline solution of arsenious acid, give a green precipitate of a copper salt called Scheele's green. Its composition is probably CuHAsO_{3}. Ammonia dissolves it, and gives a colourless solution, containing cuprous arsenate--that is, the cupric compound is reduced and the arsenic subjected to a further oxidation. The so-called Schweinfurt green was still more used, especially in former times; it is an insoluble green cupric salt, which resembles the preceding in many respects, but has a different tint. It is prepared by mixing boiling solutions of arsenious acid and cupric acetate. Arsenious acid forms an insoluble compound with ferric hydroxide, resembling the phosphate; and this is the reason why freshly precipitated oxide of iron is employed as an antidote for arsenic. The freshly precipitated oxide of iron, taken immediately after poisoning by arsenic, converts the arsenious acid into an insoluble state, by forming a compound on which the acids of the stomach have no action, so that the poisoning cannot proceed. It is remarkable that the inhabitants of certain mountainous countries accustom themselves to taking arsenic, as a means which, according to their experience, helps to overcome the fatigue of mountain ascents. Arsenious anhydride and certain of its salts are also used in medicine, naturally only in small quantities. When taken internally arsenic passes into the blood, and is mainly excreted by the urine.

[39] Adie (1889) obtained compounds of As_{2}O_{3} with 1, 2, 4, and 8 SO_{3} by the direct action of ordinary and Nordhausen sulphuric acid upon As_{2}O_{3}. Weber had previously obtained As_{2}O_{3}SO_{3} (which disengages SO_{3} at 225°), and also other As_{2}O_{3}_n_SO_{3} (where n = 3, 6, and 8), by the action of the vapours of SO_{3} upon As_{2}O_{3} at a definite temperature. The compound As_{2}O_{3},8SO_{3} loses SO_{3} at 100°. Oxide of antimony, Sb_{2}O_{3}, gives similar compounds. Adie (1891) also obtained (by the action of SO_{3} upon H_{3}PO_{4}) a compound H_{3}PO_{4}3SO_{3} in the form of a viscous liquid decomposed by water.

Antimony (stibium), Sb = 120, is another analogue of phosphorus. In its external appearance and the properties of its compounds it resembles the metals still more closely than arsenic. In fact, antimony has the appearance, lustre, and many of the characteristic properties of the metals. Its oxide, Sb_{2}O_{3}, exhibits the earthy appearance of rust or of lime, and has distinctly basic properties, although it corresponds with nitrous and phosphorous anhydride, and is able, like them, to give saline compounds with bases. At the same time antimony presents, in the majority of its compounds, an entire analogy with phosphorus and arsenic. Its compounds belong to the type SbX_{3} and SbX_{5}. It is found in nature chiefly in the form of sulphide, Sb_{2}S_{3}. This substance sometimes occurs in large masses in mineral veins and is known in mineralogy under the name of antimony glance or stibnite, and commercially as antimony (Chapter XX., Note 29). The most abundant deposits of antimony ore occur in Portugal (near Oporto on the Douro). Besides which antimony partially or totally replaces arsenic in some minerals; thus, for example, a compound of antimony sulphide and arsenic sulphide with silver sulphide is found in red silver ore. But in every case antimony is a rather rare metal found in few localities. In Russia it is known to occur in Daghestan in the Caucasus. It is extracted chiefly for the preparation of alloys with lead and tin, which are used for casting printing type.[40] Some of its compounds are also used in medicine, the most important in this respect being antimony pentasulphide, Sb_{2}S_{5} (sulfur auratum antimonii), and tartar emetic, which is a double salt derived from tartaric acid and has the composition C_{4}H_{4}K(SbO)O_{6}. Even the native antimony sulphide is used in large quantities as a purgative for horses and dogs. Metallic antimony is extracted from the glance, Sb_{2}S_{2}, by roasting, when the sulphur burns away and the antimony oxidises, forming the oxide Sb_{2}O_{3}, which is then heated with charcoal, and thus reduced to a metallic state. The reduction may be carried on in the laboratory on a small scale by fusing the sulphide with iron which takes up the sulphur.[40 bis]

[40] Printers' type consists of an alloy known as 'type-metal,' containing usually about 15 parts of antimony to 85 parts of lead; sometimes (for example, for stereotypes) from 10 to 15 per cent. Bi or 8 per cent. Sn and even Cu is added. The hardness of the alloy, which is essential for printing, evidently depends upon the presence of antimony, but an excess must be avoided, since this renders the alloy brittle, and the type after a time loses its sharpness.

[40 bis] Antimony is prepared in a state of greater purity by heating with charcoal the oxide obtained by the action of nitric acid on the impure commercial metallic antimony. This is based on the fact that by the action of the acid, antimony forms the oxide Sb_{2}O_{3}, which is but slightly soluble in water. The arsenic, which is nearly always present, forms soluble arsenious and arsenic acids, and remains in solution. The purest antimony is easily obtained from tartar emetic, by heating it with a small quantity of nitre. Metallic antimony also occurs, although rarely, native; and as it is very easily obtained, it was known to the alchemists of the fifteenth century. Very pure metallic antimony may be deposited by the electric current from a solution of antimonious sulphide in sodium sulphide after the addition of sodium chloride to the solution.

Metallic antimony has a white colour and a brilliant lustre; it remains untarnished in the air, for the metal does not oxidise at the ordinary temperature. It crystallises in rhombohedra, and always shows a distinctly crystalline structure which gives it quite a different aspect from the majority of the metals yet known. It is most like tellurium in this respect. Antimony is brittle, so that it is very easily powdered; its specific gravity is 6·7, it melts at about 432°, but only volatilises at a bright red heat. When heated in the air--for instance, before the blow-pipe--it burns and gives white odourless fumes, consisting of the oxide. This oxide is termed antimonious oxide, although it might as well be termed antimonious anhydride. It is given the first name because in the majority of cases its compounds with acids are used, but it forms compounds with the alkalis just as easily.

Antimonious oxide, like arsenious anhydride, crystallises either in regular octahedra or in rhombic prisms; its specific gravity is 5·56; when heated it becomes yellow and then fuses, and when further heated in air it oxidises, forming an oxide of the composition Sb_{2}O_{4}. Antimonious oxide is insoluble in water and in nitric acid, but it easily dissolves in strong hydrochloric acid and in alkalis, as well as in tartaric acid or solutions of its acid salts. When dissolved in the latter it forms tartar emetic. It is precipitated from its solutions in alkalis and acids (by the action of acids on the former and alkalis on the latter). It occurs native but rarely. As a base it gives salts of the type SbOX (as if the basic salts = SbX_{3}, Sb_{2}O_{3}) and hardly ever forms salts, SbX_{3}. In the antimonyl salts, SbOX, the group SbO is univalent, like potassium or silver. The oxide itself is (SbO)_{2}O, the hydroxide, SbO(OH), &c.; tartar emetic is a salt in which one hydrogen of tartaric acid is replaced by potassium and the other by antimonyl, SbO. Antimonious oxide is very easily separated from its salts by any base, but it must be observed that this separation does not take place in the presence of tartaric acid, owing to the property of tartaric acid of forming a soluble double salt--i.e. tartar emetic.[41]

[41] As antimonious oxide answers to the type SbX_{3}, it is evident that compounds may exist in which antimony will replace three atoms of hydrogen; such compounds have been to some extent obtained, but they are easily converted by water into substances corresponding with the ordinary formulæ of the compounds of antimony. Thus tartar emetic, C_{4}H_{4}(SbO)KO_{6}, loses water when heated, and forms C_{4}H_{2}SbKO_{6}--that is, tartaric acid, C_{2}H_{6}O_{6}, in which one atom of hydrogen is replaced by potassium and three by antimony. But this substance is reconverted into tartar emetic by the action of water.

A similar compound is seen in that intermediate oxide of antimony which is formed when antimonious oxide is heated in air: its composition is SbO_{2} or Sb_{2}O_{4}. This oxide may be regarded as orthantimonic acid, SbO(HO)_{3}, in which three atoms of hydrogen are replaced by antimony in that state in which it occurs in oxide of antimony--i.e. SbO(SbO_{3}) = Sb_{2}O_{4}. Oxide of antimony is also formed when antimonic acid is ignited; it then loses water and oxygen, and gives this intermediate oxide as a white infusible powder, of sp. gr. 6·7. It is somewhat soluble in water, and gives a solution which turns litmus paper red.

If metallic antimony, or antimonious oxide, be oxidised by an excess of nitric acid and the resultant mass be carefully evaporated to dryness, metantimonic acid, SbHO_{3}, is formed. Its corresponding potassium salt, 2SbKO_{3},5H_{2}O, is prepared by fusing metallic antimony with one-fourth its weight of nitre and washing the resultant mass with cold water. This potassium salt is only slightly soluble in water (in 50 parts) and the sodium salt is still less so. An ortho-acid, SbH_{3}O_{4}, also appears to exist;[41 bis] it is obtained by the action of water on antimony pentachloride, but it is very unstable, like the pentachloride, SbCl_{5}, itself, which easily gives up Cl_{2}, leaving antimony trichloride, SbCl_{3}, and this is decomposed by water, forming an oxychloride--SbOCl, only slightly soluble in water. When antimonic acid is heated to an incipient red heat, it parts with water and forms the anhydride, Sb_{2}O_{5}, of a yellow colour and specific gravity 6·5.[42]

[41 bis] Beilstein and Blaese (1889), after preparing many salts of antimonic acid, came to the conclusion that it is monobasic, but all the salts still contain water, so that their general type is mostly: MSbO_{3}3H_{2}O, for example, M = Li, Hg (salts of the suboxide), 1/2 Pb, &c. The type of the ortho-salts, M_{2}SbO_{4}, is quite unknown, although it is reproduced in the thio-compounds, for instance, Schlippe's salt, Na_{2}SbS_{4}, but this salt also contains water of crystallisation, 9H_{2}O (Chapter XX., Note 29).

[42] Among the other compounds of antimony, antimoniuretted hydrogen, SbH_{3}, resembles arseniuretted hydrogen in its mode of formation and properties (it splits up at 150°, Brunn 1890; when liquified, it boils at -65° and solidifies at -92°), whilst the halogen compounds differ in many respects from those of arsenic. When chlorine is passed over an excess of antimony powder, it forms antimony trichloride, SbCl_{3}, but if the chlorine be in excess it forms the pentachloride, SbCl_{5}. The trichloride is a crystalline substance which melts at 72° and distils at 230°, whilst the pentachloride is a yellow liquid, which splits up into chlorine and the trichloride when heated; at 140° it begins to give off chlorine abundantly, carrying away the vapour of the trichloride with it; and at 200° the decomposition is complete, and pure antimonious chloride only passes over. This property of antimony pentachloride has caused it to be applied in many cases for the transference of chlorine; all the more that when it has given up its chlorine, it leaves the trichloride, which is able to absorb a fresh amount of chlorine; and therefore many substances which are unable to react directly with gaseous chlorine do so with antimony pentachloride, and in the presence of a small quantity of it chlorine will act on them, just as oxygen is able, in the presence of nitrogen oxides, to oxidise substances which could not be oxidised by means of free oxygen. Thus carbon bisulphide is not acted on by chlorine at low temperatures--this reaction requires a high temperature--but in the presence of antimony pentachloride its conversion into carbon tetrachloride takes place at low temperatures. Antimony tri- and pentachloride, having the character of chloranhydrides, fume in air, attract moisture, and are decomposed by water, forming antimonious and antimonic acids. But in the first action of water the trichloride does not evolve all its chlorine as hydrochloric acid, which is intelligible in view of the fact that antimonious anhydride is also a base, and is therefore able to react with acids; indeed antimony sulphide dissolved in an excess of hydrochloric acid (hydrogen sulphide is evolved) gives an aqueous solution of antimony trichloride, which, when carefully distilled, even gives the anhydrous compound. Antimony trichloride is only decomposed by an excess of water, and then not completely, for with a large quantity of water it forms powder of algaroth--i.e. antimony oxychloride. The first action of water consists in the formation of oxychloride, SbOCl--that is, a salt corresponding to oxide of antimony as a base. If antimony oxide or antimony chloride be dissolved in an excess of hydrochloric acid, and the solution diluted with a considerable amount of water, then this same powder of algaroth is precipitated. The composition varies with the relative amount of water; namely, between the limits SbOCl and Sb_{4}O_{5}Cl_{2}. The latter compound is, as it were, a basic salt of the former, because its composition = 2(SbOCl)Sb_{2}O_{3}.

With bromine and iodine, antimony forms compounds similar to those with chlorine. Antimonious bromide, SbBr_{3}, crystallises in colourless prisms, melts at 94°, and boils at 270°; antimonious iodide, SbI_{3}, forms red crystals of sp. gr. 5·0; antimony trifluoride, SbF_{3} separates from a solution of antimonious oxide in hydrofluoric acid, and SbF_{5} is formed by a similar treatment of antimonic acid. The latter gives easily-soluble double salts with the fluorides of the metals of the alkalis.

De Haën (1887) obtained very stable double soluble salts, SbF_{3},KCl (100 parts of water dissolve 57 parts of salt), SbF_{3},K_{2}SO_{4}, &c., which he proposed to make use of in the arts as very easily crystallisable and soluble salts of antimony.

Engel, by passing hydrochloric acid gas into a saturated solution of antimonious chloride at 0°, obtained a compound HCl,2SbCl_{3},2H_{2}O, and with the pentachloride a compound SbCl_{5},5HCl,10H_{2}O. Bismuth trichloride, BiCl_{3}, gives a similar compound.

Saunders (1892) obtained 5RbCl,3SbCl_{3} and RbCl,SbCl_{3}. Ditte and Metzner (1892) showed that Sb and Bi dissolve in hydrochloric acid only owing to the participation of the oxygen of the air or of that dissolved in the acid.

The heaviest analogue of nitrogen and phosphorus is bismuth, Bi = 208. Here, as in the other groups, the basic, metallic, properties increase with the atomic weight. Bismuth does not give any hydrogen compound and the highest oxide, Bi_{2}O_{5}, is a very feeble acid oxide. Bismuthous oxide, Bi_{2}O_{3}, is a base, and bismuth itself a perfect metal. To explain the other properties of bismuth it must further be remarked that in the eleventh series it follows mercury, thallium and lead, whose atomic weights are near to that of bismuth, and that therefore it resembles them and more especially its nearest neighbour, lead. Although PbO and PbO_{2}, represent types different from Bi_{2}O_{3} and Bi_{2}O_{5}, they resemble them in many respects, even in their external appearance, moreover the lower oxides both of Pb and Bi are basic and the higher acid, which easily evolve oxygen. But judging by the formula, Bi_{2}O_{3} is a more feeble base than PbO. They both easily give basic salts.

Bismuth forms compounds of two types, BiX_{3} and BiX_{5},[43] which entirely recall the two types we have already established for the compounds of lead. Just as in the case of lead, the type PbX_{2}, is basic, stable, easily formed, and passes with difficulty into the higher and lower types, which are unstable, so also in the case of bismuth the type of combination BiX_{3} is the usual basic form. The higher type of combination, BiX_{5},[44] in fact behaves toward this stable type, BiX_{3}, in exactly the same manner as lead dioxide does to the monoxide; and bismuthic acid is obtained by the action of chlorine on bismuth oxide suspended in water, in exactly the same way as lead dioxide is obtained from lead oxide. It is an oxidising agent like lead dioxide, and even the acid character in bismuthic acid is only slightly more developed than in lead dioxide. Here, as in the case of lead (minium), intermediate compounds are easily formed in which the bismuth of the lower oxide plays the part of a base combined with the acid which is formed by the higher form of the oxidation of bismuth.

[43] Metallic bismuth is very easily obtained when the compounds of the oxide are reduced by powerful reducing agents, but when less powerful reducing agents--for example, stannous oxide--are taken, bismuth suboxide is formed as a black crystalline powder. It is a compound of the type BiX_{2}, its composition being BiO; it is decomposed by acids into the metal and oxide, which passes into solution.

[44] The type BiX_{5} is represented by the pentoxide, Bi_{2}O_{5}, its metahydrate, Bi_{2}O_{5},H_{2}O, or BiHO_{3}, known as bismuthic acid, and the pyrohydrate, Bi_{2}H_{4}O_{7}. Bismuth pentoxide is obtained by the prolonged passage of chlorine through a boiling solution of potassium hydroxide (sp. gr. 1·38), containing bismuth oxide in suspension; the precipitate is washed with water, with boiling nitric acid (but not for long, as otherwise the bismuthic acid is decomposed), then again with water, and finally the resultant bright red powder of the hydrate BiHO_{3} is dried at 125°. The prolonged action of nitric acid on bismuthic anhydride, Bi_{2}O_{5}, results in the formation of the compound Bi_{2}O_{4},H_{2}O, which decomposes in moist air, forming Bi_{2}O_{3}. The density of bismuthic anhydride is 5·10, of the tetroxide, Bi_{2}O_{4}, 3·60, and of bismuthic acid, BiHO_{3}, 5·75. Pyrobismuthic acid, Bi_{2}H_{4}O_{7}, forms a brown powder, which loses a portion of its water at 150°, and decomposes on further heating, with the evolution of oxygen and water. It is obtained by the action of potassium cyanide on a solution of bismuth nitrate. The meta-salts of bismuthic acid are known, for example KBiO_{3}. They generally occur, however, in combinations with metabismuthic acid itself. Thus André (1891) took a solution of the double salt of BiBr_{3} and KBr, treated it with bromine after adding ammonia, and obtained a red-brown precipitate, which after being washed (for several weeks) had the composition KBiO_{3},HBiO_{3} When washed with dilute nitric acid this salt gave bismuthic acid.

In nature, bismuth occurs in only a few localities and in small quantities, most frequently in a native state, and more rarely as oxide and as a compound of bismuth sulphide with the sulphides of other metals, and sometimes in gold ores. It is extracted from its native ores by simple fusion in the furnace shown in fig. 85. This furnace contains an inclined iron retort, into the upper extremity of which the ore is charged, and the molten metal flows from the lower extremity. It is refined by re-melting, and the pure metal may be obtained by dissolving in nitric acid, decomposing the resultant salt with water, and reducing the precipitate by heating it with charcoal. Bismuth is a metal which crystallises very well from a molten state. Its specific gravity is 9·8; it melts at 269°, and if it be melted in a crucible, allowed to cool slowly, and the crust broken and the remaining molten liquid poured out, perfect rhombohedral crystals of bismuth are obtained on the sides of the crucible.[44 bis] It is brittle, has a grey-coloured fracture with a reddish lustre, is not hard, and is but very slightly ductile and malleable; it volatilises at a white heat and easily oxidises. It recalls antimony and lead in many of its properties. When oxidised in air, or when the nitrate is ignited, bismuth forms the oxide, Bi_{2}O_{3}, as a white powder which fuses when heated and resembles massicot. The addition of an excess of caustic potash to a solution of a bismuthous salt gives a white precipitate of the hydroxide, BiO(OH), which loses its water and gives the anhydrous oxide when boiled with a solution of caustic potash. Both the hydroxide and oxide easily dissolve in acids and form bismuthous salts.

[44 bis] Hérard (1889) obtained a peculiar variety of bismuth by heating pure crystalline bismuth to a bright red heat in a stream of nitrogen. A greenish vapour was deposited in the cooler portions of the apparatus in the form of a grey powder, which under the microscope had the appearance of minute globules. An atmosphere of nitrogen is necessary for this transformation, other gases such as hydrogen and carbonic oxide do not favour the transition. The resultant amorphous bismuth fuses at 410° (the crystalline variety at 269°), sp. gr. 9·483. (Does it not contain a nitride?)

Bismuthous oxide, Bi_{2}O_{3}, is a feeble and unenergetic base. The normal hydroxide of the oxide Bi_{2}O_{3} is Bi(OH){3}; it parts with water and forms a metahydroxide (bismuthyl hydroxide), BiO(OH). Both of these hydroxides have their corresponding saline compounds of the composition BiX{3} and BiOX. And the form BiOX is nothing else but the type of the basic salt, because 3ROX = RX + R_{2}O_{3}. It is evident that in the type BiX_{3} the bismuth replaces three atoms of hydrogen. And indeed with phosphoric acid solutions of the bismuthous salts give a precipitate of the composition BiPO_{4}. On the other hand, in the form of compounds BiOX or Bi(OH){2}X, the univalent group (BiO) or (BiH{2}O_{2}) is combined with X. Many bismuth salts are formed according to the type BiOX. For instance the carbonate, (BiO){2}CO{3}, which corresponds with the other carbonates M_{2}CO_{3}. It is obtained as a white precipitate when a solution of sodium carbonate is added to a solution of a bismuth salt.[45] The compound radicle BiO is not a special natural grouping, as it was formerly represented to be; it is simply a mode of expression for showing the relation between the compound in question and the compounds of other oxides.

[45] Basic bismuth carbonate is employed for whitening the skin (veloutine, &c.)

Three salts of nitric acid are known containing bismuthous oxide. If metallic bismuth or its oxide be dissolved in nitric acid, it forms a colourless transparent solution containing a salt which separates in large transparent crystals containing Bi(NO_{3}){3},5H{2}O. When heated at 80° these crystals melt in their water of crystallisation, and in so doing lose a portion of their nitric acid together with water, forming a salt whose empirical formula is Bi_{2}N_{2}H_{2}O_{9}. If the preceding salt belongs to the type BiX_{3}, this one should belong to the form BiOX, because it = BiO(NO_{3}) + Bi(H_{2}O_{2})(NO_{3}). This salt may be heated to 150° without change. When the first colourless crystalline salt dissolves in water it is decomposed. There is no decomposition if an excess of acid be added to the water--that is to say, the salt is able to exist in an acid solution without decomposing, without separation of the so-called basic salt--but by itself it cannot be kept in solution; water decomposes this salt, acting on it like an alkali. In other words the basic properties of bismuthic oxide are so feeble that even water acts by taking up a portion of the acid from it. Here we see one of the most striking facts, long since observed, confirming that action of water on salts about which we have spoken in Chapter X. and elsewhere. This action on water may be expressed thus:--BiX_{3} + 2H_{2}O = Bi(OH){2}X + 2XH. A salt of the type Bi(OH){2}X is obtained in the precipitate. But if the quantity of acid, HX, be increased, the salt BiX_{3} is again formed and passes into solution. The quantity of the salt BiOX which passes into solution on the addition of a given quantity of acid depends indisputably on the amount (mass) of water (Muir). The solution, which is perfectly transparent with a small amount of water, becomes cloudy and deposits the salt of the type BiOX, when diluted. The white flaky precipitate of Bi(OH){2}NO{3} formed from the normal salt Bi(NO_{3})_{3} by mixing it with five parts of water, and in general with a small amount of water, is used in medicine under the name of magistery of bismuth.[46]

[46] With an excess of water a further quantity of acid is separated and a still more basic salt formed. The ultimate product, on which an excess of water has apparently no action whatever, is a substance having the composition BiO(NO_{3}).BiO(OH). In the latter salt we see the limit of change, and this limit appears to show that the type of the saline compounds of bismuthic oxide is of the form Bi_{2}X_{6}, and not BiX_{3}; but it is very probable, on the basis of the examples which we considered in the case of lead, that this type should be still further polymerised in order to give a correct idea of the type of the bismuthous compounds. If we refer all the bismuthous compounds to this type, Bi_{2}X_{6}, we shall obtain the following expression for the composition of the nitrates: normal salt, Bi_{2}(NO_{3}){6}, first basic salt, Bi{2}O(OH){2}(NO{3}){2}, magistery of bismuth, Bi{2}(OH){4}(NO{3}){2}, and the limiting form Bi{2}O_{2}(OH)(NO_{3}).

The general character of bismuthous oxide in its compounds is well exemplified in the nitrate; bismuthous chloride, BiCl_{3}, which is obtained by heating bismuth in chlorine, or by dissolving it in aqua regia, and then distilling without access of air, is also decomposed by water in exactly the same manner, and forms basic salts--for instance, first, BiOCl, like the above salt of nitric acid. Bismuth chloride boils at 447° and probably its formula is BiCl_{3}. Polymerisation may take place in some compounds and not in others. A volatile compound of the composition Bi(C_{2}H_{5}){3} is also known as a liquid which is insoluble in water and decomposes with explosion when heated at 130°. Double salts containing chloride of bismuth are: 2(KCl)BiCl{3}2H_{2}O (from a solution of Bi_{2}O_{3} and KCl in hydrochloric acid) and KClBiCl_{3}H_{2}O. Bigham (1892) also obtained KBr(SO_{4})_{2} in tabular crystals by treating the above-named double salt with strong sulphuric acid. The composition of this salt recalls that of alum.

Metallic bismuth is used in the preparation of fusible alloys. The addition of bismuth to many metals renders them very hard, and at the same time generally lowers their melting point to a considerable extent. Thus Wood's metal, which contains one part of cadmium, one part of tin, two parts of lead, and four parts of bismuth, fuses at about 60°, and in general many alloys composed of bismuth, tin, lead, and antimony melt below or about the boiling point of water.[47]

[47] As the metals contained in alloys like the above (bismuth, lead, tin, cadmium) are difficultly volatile and their alloys are fusible, they may be employed in the place of mercury in many physical experiments conducted at or above 70°, and they offer the advantage that they do not give any vapour having an appreciable tension (mercury at 100°, 0·75 mm.) Bismuth expands in passing into a molten state, but it has a temperature of maximum density. According to Luedeking the mean coefficient of expansion of liquid bismuth is 0·0000442 (between 270° and 303°), and of solid bismuth 0·0000411.

Just as in group II., side by side with the elements zinc, cadmium, and mercury in the uneven series, we found calcium, strontium, and barium in the even series; and as in group IV., parallel to silicon, germanium, tin, and lead, we noticed thallium, zirconium, cerium, and thorium; so also in group V. we find, beside those elements of the uneven series just considered by us, a series of analogues in the even series, which, with a certain degree of similarity (mainly quantitative, or relative to the atomic weights), also present a series of particular (qualitative) independent points of distinction. In the even series are known vanadium, which stands between titanium and chromium, niobium, between zirconium and molybdenum, and tantalum, situated near tungsten (an element of group VI. like chromium and molybdenum). Just as bismuth is similar in many respects to its neighbour lead, so also do these neighbouring elements resemble each other, even in their external appearance, not to mention the quality of their compounds, naturally taking into account the differences of type corresponding with the different groups. The occurrence in group V. determines the type of the oxides, R_{2}O_{3} and R_{2}O_{5}, and the development of an acid character in the higher oxides. The occurrence in the even series determines the absence of volatile compounds, RH_{3}, for these metals, and a more basic character of the oxides of a given composition than in the uneven series, &c.[48] Vanadium, niobium, and tantalum belong to the category of rare metals, and are exceedingly difficult to obtain pure, more especially owing to their similarity to, and occurrence with, chromium, tungsten and other metals, and also in combination among themselves; therefore it is natural that they have been far from completely studied, although since 1860 chemists have devoted not a little time to their investigation. The researches carried out by Marignac, at Geneva, on niobium, and by Sir Henry Roscoe, at Manchester, on vanadium deserve special attention. The undoubted external resemblance of the compounds of chromium and vanadium, as well as the want of completeness in the knowledge of the compounds of vanadium, long caused its oxides to be considered analogous in atomic composition to those formed by chromium. The higher oxide of vanadium was therefore supposed to have the formula VO_{3}. But the fact of the matter is, that the chemical analogy of the elements does not hold in one direction only; vanadium is at one and the same time the analogue of chromium, and consequently of the elements like sulphur of group VI, and also the analogue of phosphorus, arsenic, and antimony; just as bismuth stands in respect to lead and antimony. Investigation has shown that the compounds of vanadium are always accompanied by those of phosphorus as well as of iron, and that it is even more difficult to separate it from the compounds of phosphorus than from those of iron and tungsten. We should have to extend our description considerably if we wished to give the complete history, even of vanadium alone, not to mention niobium and tantalum, all the more as questions would not unfrequently arise concerning the compounds of these elements which have not yet been fully elucidated. We shall therefore limit ourselves to pointing out the most important features in the history of these elements, the more so since the minerals themselves in which they occur are exceedingly rare and only accessible to a few investigators.

[48] Although, guided by Brauner, who showed that didymium gives a higher oxide, Di_{2}O_{5}, I place this element in the fifth group, still I am not certain as to its position, because I consider that the questions relating to this metal are still far from being definitely answered.

An important point in the history of the members of this group is the circumstance that they form volatile compounds with chlorine, similar to the compounds of the elements of the phosphorus group, namely, of the type RX_{5}. The vapour densities of the compounds of this kind were determined, and served as the most important basis for the explanation of the atomic composition of these molecules. In this we see the power of general and fundamental laws, like the law of Avogadro-Gerhardt. An oxychloride, VOCl_{3}, is known for vanadium, which is the perfect analogue of phosphorus oxychloride. It was formerly considered to be vanadium chloride, for just as in the case of uranium (Chapter XXI.), its lower oxide, VO, was considered to be the metal, because it is exceedingly difficultly reduced--even potassium does not remove all the oxygen, besides which it has a metallic appearance, and decomposes acids like a metal; in a word, it simulates a metal in every respect. Vanadium oxychloride is obtained by heating the trioxide, V_{2}O_{3}, mixed with charcoal, in a current of hydrogen; the lower oxide of vanadium is then formed, and this, when heated in a current of dry chlorine, gives the oxychloride VOCl_{3} as a reddish liquid which does not act on sodium and may be purified by distillation over this metal. It fumes in the air, giving reddish vapours; it reacts on water, forming hydrochloric and vanadic acids; hence, on the one hand it is very similar to phosphorus oxychloride, and on the other hand to chromium oxychloride, CrO_{2}Cl_{2} (Chapter XXI.). It is of a yellow colour, its specific gravity is 1·83, it boils at 120°, and its vapour density is 86 with respect to hydrogen; therefore the above formula expresses its molecular weight.[49]

[49] When the vapours of vanadium oxychloride are heated with zinc in a closed tube at 400°, they lose a portion of their chlorine and form a green crystalline mass of sp. gr. 2·88, which is deliquescent in air and has the composition VOCl_{2}. Only its vapour density is unknown, and it would be extremely important to determine whether its molecular composition is that given above, or whether it corresponds with the formula V_{2}O_{2}Cl_{4}. Another less volatile oxychloride, VOCl, is formed with it as a brown insoluble substance, which is, however, soluble in nitric acid like the preceding. Roscoe obtained a still less chlorinated substance, namely, (VO)_{2}Cl; but it may only consist of a mixture of VO and VOCl. At all events, we here find a graduated series such as is met with in the compounds of very few other elements.

Vanadic anhydride, V_{2}O_{5}, is obtained either in small quantities from certain clays where it accompanies the oxides of iron (hence some sorts of iron contain vanadium) and phosphoric acid, or from the rare minerals: volborthite, CuHVO_{4}, or basic vanadate of copper; vanadinite, PbCl_{2}3Pb_{3}(VO_{4}){2}; lead vanadate, Pb{3}(VO_{4}){2}, &c. The latter salts are carefully ignited for some time with one-third of their weight of nitre; the fused mass thus formed is powdered and boiled in water: the yellow solution obtained contains potassium vanadate. The solution is neutralised with acid, and barium chloride added; a meta-salt, Ba(VO{3}){2}, is then precipitated as an almost insoluble white powder, which gives a solution of vanadic acid when boiled with sulphuric acid. (The precipitate is at first yellow, as long as it remains amorphous, but it afterwards becomes crystalline and white.) The solution thus obtained is neutralised with ammonia, which thus forms ammonium (meta) vanadate, NH{4}VO_{3}, which, when evaporated, gives colourless crystals, insoluble in water containing sal-ammoniac; hence this salt is precipitated by adding solid sal-ammoniac to the solution. Ammonium vanadate, when ignited, leaves vanadic acid behind. In this it differs from the corresponding chromium salt, which is deoxidised into chromium oxide when ignited. In general, vanadic acid has but a small oxidising action. It is reduced with difficulty, like phosphoric or sulphuric acid, and in this differs from arsenic and chromic acids. Vanadic acid, like chromic acid, separates from its solution as the anhydride V_{2}O_{5}, and not in a hydrous state. Vanadic anhydride, V_{2}O_{5}, forms a reddish-brown mass, which easily fuses and re-solidifies into transparent crystals having a violet lustre (another point of resemblance to chromic acid); it dissolves in water, forming a yellow solution with a slightly acid reaction.[50]

[50] Strong acids and alkalis dissolve vanadic anhydride in considerable quantities, forming yellow solutions. When it is ignited, especially in a current of hydrogen, it evolves oxygen and forms the lower oxides; V_{2}O_{4} (acid solutions of a green colour, like the salts of chromic oxide), V_{2}O_{3}, and the lowest oxide, VO. The latter is the metallic powder which is obtained when the vanadium oxychloride is heated in an excess of hydrogen, and was formerly mistaken for metallic vanadium. When a solution of vanadic acid is treated with metallic zinc it forms a blue solution, which seems to contain this oxide. It acts as a reducing agent (and forms a close analogue to chromous oxide, CrO). Metallic vanadium can only be obtained from vanadium chloride which is quite free from oxygen. Moissan (1893) obtained it by reducing the oxide with carbon in the electric furnace, and considered it to be most infusible of the metals in the series Pt, Cr, Mo, U, W, and V (he also obtained a compound of vanadium and carbon). The specific gravity of this metal is 5·5. It is of a grey-white colour, is not decomposed by water, and is not oxidised in air, but burns when strongly heated, and can be fused in a current of hydrogen (forming perhaps a compound with hydrogen). It is insoluble in hydrochloric acid, but easily dissolves in nitric acid, and when fused with caustic soda it forms sodium vanadate.

As regards the salts of vanadic acid, three different classes are known; the first correspond with metavanadic acid, VMO_{3} = M_{2}OV_{2}O_{5}, the second correspond with the dichromates--that is, have the composition V_{4}M_{2}O_{11}, which is equal to M_{2}O + 2V_{2}O_{5}--and the third correspond with orthovanadic acid, VM_{3}O_{4} or 3M_{2}O + V_{2}O_{5}. The latter are formed when vanadic anhydride is fused with an excess of an alkaline carbonate.

Vanadic acid gives the so-called 'complex' acids (which are considered more fully in Chapter XXI. in speaking of Mo and W)--i.e. acids formed of two acids assimilated into one. Thus Friedheim (1890) obtained phosphor-vanadic acid, and Schmitz-Dumont (1890) a similar arseno-vanadic acid. The former is obtained by heating V_{2}O_{5} with sirupy phosphoric acid. The resultant golden-yellow tabular crystals have the composition H_{2}OV_{2}O_{5}P_{2}O_{5}9H_{2}O, and there are corresponding salts--for example, (NH_{4}){2}V{2}O_{5}P_{2}O_{5} with 3 and 7H_{2}O, &c. These salts cannot be separated by crystallisation, so that there are 'complexes' of these acids in a whole series of salts (and also in nature). It may be supposed (Friedheim) that V_{2}O_{5} here, as it were, plays the part of a base, or that those acids may be looked upon as double salts. Among the true double salts of vanadium (Nb and Ta) very many are known among the fluorides, such as VF_{3}2NH_{4}F, VOF_{2}2NH_{4}F, VO_{2}F,3NH_{4}F, &c. (Pettersson, Piccini, and Georgi, 1890-92).

Vanadium was discovered at the beginning of this century by Del-Rio, and afterwards investigated by Sefström, but it was only in 1868 that Roscoe established the above formulæ of the vanadic compounds.

Niobium and tantalum[51] occur as acids in rare minerals, and are mainly extracted from tantalite and columbite, which are found in Bavaria, Finland, North America, and in the Urals. These minerals are composed of the ferrous salts of niobic and tantalic acids; they contain about 15 per cent. of ferrous oxide in isomorphous mixture with manganous oxide, in combination with various proportions of tantalic and niobic anhydrides. These minerals are first fused with a considerable amount of potassium bisulphate, and the fused mass is boiled in water, which dissolves the ferrous and potassium salts and leaves an insoluble residue of impure niobic and tantalic acids. This raw product is then treated with ammonium sulphide, in order to extract the tin and tungsten, which pass into solution. The residue containing the acids (according to Marignac) is then treated with hydrofluoric acid, in which it entirely dissolves, and potassium fluoride is added to the resultant hot solution; on cooling, a sparingly soluble double fluoride of potassium and tantalum separates out in fine crystals, while the much more soluble niobium salt remains in solution. The difference in the solubility of these double salts in water acidified with hydrofluoric acid (in pure water the solution becomes cloudy after a certain time) is so great that the tantalum compound requires 150 parts of water for its solution, and the niobium compound only 13 parts. The Greenland columbite (specific gravity 5·36) only contains niobic acid, and that from Bodenmais, Bavaria (specific gravity 6·06) almost equal quantities of tantalic and niobic acids. Having isolated tantalic and niobic salts, Marignac found that the relation between the potassium and fluorine in them is very variable--that is, that there exist various double salts of fluoride of potassium, and of the fluorides of the metals of this group, but that with an excess of hydrofluoric acid both the tantalum and niobium compounds contain seven atoms of fluorine to two of potassium, whence it must be concluded that the simplest formula for these double salts will be K_{2}RF_{7} = RF_{5},2KF; that is, that the type of the higher compounds of niobium and tantalum is RX_{5}, and hence is similar to phosphoric acid. A chloride, TaCl_{5}, may be obtained from pure tantalic acid by heating it with charcoal in a current of chlorine. This is a yellow crystalline substance, which melts at 211°, and boils at 241°; its vapour density with respect to hydrogen is 180, as would follow from the formula TaCl_{5}. It is completely decomposed by water into tantalic and hydrochloric acids. Niobium pentachloride may be prepared in the same manner; it fuses at 194°, and boils at 240°. When treated with water this substance gives a solution containing niobic acid, which only separates out on boiling the solution. Delafontaine and Deville found its vapour density to be 9·3 (air = 1), as is shown by its formula NbCl_{5}.[52]

[51] The researches made by Roscoe were preceded by those of Marignac in 1865, on the compounds of niobium and tantalum, to which were also ascribed different formulæ from those now recognised. Tantalum was discovered simultaneously with vanadium by Hatchett and Ekeberg, and was afterwards studied by Rose, who in 1844 discovered niobium in it. Notwithstanding the numerous researches of Hermann (in Moscow), Kobell, Rose, and Marignac, still there is not yet any certainty as to the purity of, and the properties ascribed to, the compounds of these elements. They are difficult to separate from each other, and especially from the cerite metals and titanium, &c., which accompany them. Before the investigations of Rose the highest oxide of tantalum was supposed to belong to the type TaX_{6}--that is, its composition was taken as TaO_{3}, and to the lower oxide was ascribed a formula TaO_{2}. Rose gave the formula TaO_{2} to the higher oxide, and discovered a new element called niobium in the substance previously supposed to be the lower oxide. He even admitted the existence of a third element occurring together with tantalum and niobium, which he named pelopium, but he afterwards found that pelopic acid was only another oxide of niobium, and he considered it probable that the higher oxide of this element is NbO_{2}, and the lower Nb_{2}O_{3}. Hermann found that niobic acid which was considered pure contained a considerable quantity of tantalic acid, and besides this he admitted the existence of another special metallic acid, which he called ilmenic acid, after the locality (the Ilmen mountains of the Urals) of the mineral from which he obtained it. V. Kobell recognised still another acid, which he called dianic acid, and these diverse statements were only brought into agreement in the sixties by Marignac. He first of all indicated an accurate method for the separation of tantalic and niobic compounds, which are always obtained in admixture.Return to passage

[52] If niobic acid be mixed with a small quantity of charcoal and ignited in a stream of chlorine, a difficultly-fusible and difficultly-volatile oxychloride, NbOCl_{3} separates. The vapour density of this compound with respect to air is 7·5, and this vapour density perfectly confirms the accuracy of the formulæ given by Marignac, and indicates the quantitative analogy between the compounds of niobium and tantalum, and those of phosphorus and arsenic, and consequently also of vanadium. In their qualitative relations (as is evident also from the correspondence of the atomic weights), the compounds of tantalum and niobium exhibit a great analogy with the compounds of molybdenum and tungsten. Thus zinc, when acting on acid solutions of tantalic and niobic compounds, gives a blue coloration, exactly as it does with those of tungsten and molybdenum (also titanium). These acids form the same large number of salts as those of tungsten and molybdenum. The anhydrides of the acids are also insoluble in water, but as colloids are sometimes held in solution, just like those of titanic and molybdic acids. Furthermore, niobium is in every respect the nearest analogue of molybdenum, and tantalum of tungsten. Niobium is obtained by reducing the double fluoride of niobium and sodium, with sodium. It is difficult to obtain in a pure state. It is a metal on which hydrochloric acid acts with some energy, as also does hydrofluoric acid mixed with nitric acid, and also a boiling solution of caustic potash. Tantalum, which is obtained in exactly the same way, is a much heavier metal. It is infusible, and is only acted on by a mixture of hydrofluoric and nitric acids. Rose in 1868 showed that in the reduction of the double fluoride, NbF_{5},2KF, by sodium, a greyish powder is obtained after treating with water. The specific gravity of this powder is 6·8, and he considers it to be niobium hydride, NbH. Neither did he obtain metallic niobium when he reduced with magnesium and aluminium, but an alloy, Al_{3}Nb, having a sp. gr. of 4·5.Return to passage

Niobium, so far as is known, unites in three proportions with oxygen. NbO, which is formed when NbOF_{3},2KF is reduced by sodium; NbO_{2}, which is formed by igniting niobic acid in a stream of hydrogen, and niobic anhydride, Nb_{2}O_{5}, a white infusible substance, which is insoluble in acids, and has a specific gravity of 4·5. Tantalic anhydride closely resembles niobic anhydride, and has a specific gravity of 7·2. The tantalates and niobates present the type of ortho-salts--for example, Na_{2}HNbO_{4},6H_{2}O, and also of pyro-salts, such as K_{3}HNb_{2}O_{7},6H_{2}O, and of meta-salts--for example, KNbO_{3},2H_{2}O. And, besides these, they give salts of a more complex type, containing a larger amount of the elements of the anhydride; thus, for instance, when niobic anhydride is fused with caustic potash it forms a salt which is soluble in water, and crystallises in monoclinic prisms, having the composition K_{8}Nb_{6}O_{19},16H_{2}O. There is a perfectly similar isomorphous salt of tantalic acid. Tantalite is a salt of the type of metatantalic acid, Fe(TaO_{3})_{2}. The composition of Yttrotantalite appears to correspond with orthotantalic acid.