PHOSPHORUS

History. The element phosphorus was discovered by the alchemist Brand, of Hamburg, in 1669, while searching for the philosopher's stone. Owing to its peculiar properties and the secrecy which was maintained about its preparation, it remained a very rare and costly substance until the demand for it in the manufacture of matches brought about its production on a large scale.

Occurrence. Owing to its great chemical activity phosphorus never occurs free in nature. In the form of phosphates it is very abundant and widely distributed. Phosphorite and sombrerite are mineral forms of calcium phosphate, while apatite consists of calcium phosphate together with calcium fluoride or chloride. These minerals form very large deposits and are extensively mined for use as fertilizers. Calcium phosphate is a constituent of all fertile soil, having been supplied to the soil by the disintegration of rocks containing it. It is the chief mineral constituent of bones of animals, and bone ash is therefore nearly pure calcium phosphate.

Preparation. Phosphorus is now manufactured from bone ash or a pure mineral phosphate by heating the phosphate with sand and carbon in an electric furnace. The materials are fed in at M (Fig. 70) by the feed screw F. The phosphorus vapor escapes at P and is condensed under water, while the calcium silicate is tapped off as a liquid at S. The phosphorus obtained in this way is quite impure, and is purified by distillation.

Fig. 70

Explanation of the reaction. To understand the reaction which occurs, it must be remembered that a volatile acid anhydride is expelled from its salts when heated with an anhydride which is not volatile. Thus, when sodium carbonate and silicon dioxide are heated together the following reaction takes place:

Na₂CO₃ + SiO₂ = Na₂SiO₃ + CO₂.

Silicon dioxide is a less volatile anhydride than phosphoric anhydride (P₂O₅), and when strongly heated with a phosphate the phosphoric anhydride is driven out, thus:

Ca₃(PO₄)₂ + 3SiO₂ = 3CaSiO₃ + P₂O₅.

If carbon is added before the heat is applied, the P₂O₅ is reduced to phosphorus at the same time, according to the equation

P₂O₅ + 5C = 2P + 5CO.

Physical properties. The purified phosphorus is a pale yellowish, translucent, waxy solid which melts at 43.3° and boils at 269°. It can therefore be cast into any convenient form under warm water, and is usually sold in the market in the form of sticks. It is quite soft and can be easily cut with a knife, but this must always be done while the element is covered with water, since it is extremely inflammable, and the friction of the knife blade is almost sure to set it on fire if cut in the air. It is not soluble in water, but is freely soluble in some other liquids, notably in carbon disulphide. Its density is 1.8.

Chemical properties. Exposed to the air phosphorus slowly combines with oxygen, and in so doing emits a pale light, or phosphorescence, which can be seen only in a dark place. The heat of the room may easily raise the temperature to the kindling point of phosphorus, when it burns with a sputtering flame, giving off dense fumes of oxide of phosphorus. It burns with dazzling brilliancy in oxygen, and combines directly with many other elements, especially with sulphur and the halogens. On account of its great affinity for oxygen it is always preserved under water.

Phosphorus is very poisonous, from 0.2 to 0.3 gram being a fatal dose. Ground up with flour and water or similar substances, it is often used as a poison for rats and other vermin.

Precaution. The heat of the body is sufficient to raise phosphorus above its kindling temperature, and for this reason it should always be handled with forceps and never with the bare fingers. Burns occasioned by it are very painful and slow in healing.

Red phosphorus. On standing, yellow phosphorus gradually undergoes a remarkable change, being converted into a dark red powder which has a density of 2.1. It no longer takes fire easily, neither does it dissolve in carbon disulphide. It is not poisonous and, in fact, seems to be an entirely different substance. The velocity of this change increases with rise in temperature, and the red phosphorus is therefore prepared by heating the yellow just below the boiling point (250°-300°). When distilled and quickly condensed the red form changes back to the yellow. This is in accordance with the general rule that when a substance capable of existing in several allotropic forms is condensed from a gas or crystallized from the liquid state, the more unstable variety forms first, and this then passes into the more stable forms.

Matches. The chief use of phosphorus is in the manufacture of matches. Common matches are made by first dipping the match sticks into some inflammable substance, such as melted paraffin, and afterward into a paste consisting of (1) phosphorus, (2) some oxidizing substance, such as manganese dioxide or potassium chlorate, and (3) a binding material, usually some kind of glue. On friction the phosphorus is ignited, the combustion being sustained by the oxidizing agent and communicated to the wood by the burning paraffin. In sulphur matches the paraffin is replaced by sulphur.

In safety matches red phosphorus, an oxidizing agent, and some gritty material such as emery is placed on the side of the box, while the match tip is provided as before with an oxidizing agent and an easily oxidized substance, usually antimony sulphide. The match cannot be ignited easily by friction, save on the prepared surface.

Compounds of phosphorus with hydrogen. Phosphorus forms several compounds with hydrogen, the best known of which is phosphine (PH₃) analogous to ammonia (NH₃).

Preparation of phosphine. Phosphine is usually made by heating phosphorus with a strong solution of potassium hydroxide, the reaction being a complicated one.

Fig. 71

The experiment can be conveniently made in the apparatus shown in Fig. 71. A strong solution of potassium hydroxide together with several small bits of phosphorus are placed in the flask A, and a current of coal gas is passed into the flask through the tube B until all the air has been displaced. The gas is then turned off and the flask is heated. Phosphine is formed in small quantities and escapes through the delivery tube, the exit of which is just covered by the water in the vessel C. Each bubble of the gas as it escapes into the air takes fire, and the product of combustion (P₂O₅) forms beautiful small rings, which float unbroken for a considerable time in quiet air. The pure phosphine does not take fire spontaneously. When prepared as directed above, impurities are present which impart this property.

Properties. Phosphine is a gas of unpleasant odor and is exceedingly poisonous. Like ammonia it forms salts with the halogen acids. Thus we have phosphonium chloride (PH₄Cl) analogous to ammonium chloride (NH₄Cl). The phosphonium salts are of but little importance.

Oxides of phosphorus. Phosphorus forms two well-known oxides,—the trioxide (P₂O₃) and the pentoxide (P₂O₅), sometimes called phosphoric anhydride. When phosphorus burns in an insufficient supply of air the product is partially the trioxide; in oxygen or an excess of air the pentoxide is formed. The pentoxide is much the better known of the two. It is a snow-white, voluminous powder whose most marked property is its great attraction for water. It has no chemical action upon most gases, so that they can be very thoroughly dried by allowing them to pass through properly arranged vessels containing phosphorus pentoxide.

Acids of phosphorus. The important acids of phosphorus are the following:

H₃PO₃	phosphorous acid.
H₃PO₄	phosphoric acid.
H₄P₂O₇	pyrophosphoric acid.
HPO₃	metaphosphoric acid.

These may be regarded as combinations of the oxides of phosphorus with water according to the equations given in the discussion of the characteristics of the family.

1. Phosphorous acid (H₃PO₃). Neither the acid nor its salts are at all frequently met with in chemical operations. It can be easily obtained, however, in the form of transparent crystals when phosphorus trichloride is treated with water and the resulting solution is evaporated:

PCl₃ + 3H₂O = H₃PO₃ + 3HCl.

Its most interesting property is its tendency to take up oxygen and pass over into phosphoric acid.

2. Orthophosphoric acid (phosphoric acid) (H₃PO₄). This acid can be obtained by dissolving phosphorus pentoxide in boiling water, as represented in the equation

P₂O₅ + 3H₂O = 2H₃PO₄.

It is usually made by treating calcium phosphate with concentrated sulphuric acid. The calcium sulphate produced in the reaction is nearly insoluble, and can be filtered off, leaving the phosphoric acid in solution. Very pure acid is made by oxidizing phosphorus with nitric acid. It forms large colorless crystals which are exceedingly soluble in water. Being a tribasic acid, it forms acid as well as normal salts. Thus the following compounds of sodium are known:

NaH₂PO₄	monosodium hydrogen phosphate.
Na₂HPO₄	disodium hydrogen phosphate.
Na₃PO₄	normal sodium phosphate.

These salts are sometimes called respectively primary, secondary, and tertiary phosphates. They may be prepared by bringing together phosphoric acid and appropriate quantities of sodium hydroxide. Phosphoric acid also forms mixed salts, that is, salts containing two different metals. The most familiar compound of this kind is microcosmic salt, which has the formula Na(NH₄)HPO₄.

Orthophosphates. The orthophosphates form an important class of salts. The normal salts are nearly all insoluble and many of them occur in nature. The secondary phosphates are as a rule insoluble, while most of the primary salts are soluble.

3. Pyrophosphoric acid (H₄P₂O₇). On heating orthophosphoric acid to about 225° pyrophosphoric acid is formed in accordance with the following equation:

2H₃PO₄ = H₄P₂O₇ + H₂O.

It is a white crystalline solid. Its salts can be prepared by heating a secondary phosphate:

2Na₂HPO₄ = Na₄P₂O₇ + H₂O.

4. Metaphosphoric acid (glacial phosphoric acid) (HPO₃). This acid is formed when orthophosphoric acid is heated above 400°:

H₃PO₄ = HPO₃ + H₂O.

It is also formed when phosphorus pentoxide is treated with cold water:

P₂O₅ + H₂O = 2HPO₃.

It is a white crystalline solid, and is so stable towards heat that it can be fused and even volatilized without decomposition. On cooling from the fused state it forms a glassy solid, and on this account is often called glacial phosphoric acid. It possesses the property of dissolving small quantities of metallic oxides, with the formation of compounds which, in the case of certain metals, have characteristic colors. It is therefore used in the detection of these metals.

While the secondary phosphates, on heating, give salts of pyrophosphoric acid, the primary phosphates yield salts of metaphosphoric acid. The equations representing these reactions are as follows:

2Na₂HPO₄ = Na₄P₃O₇ + H₂O,

NaH₂PO₄ = NaPO₃ + H₂O.

Fertilizers. When crops are produced year after year on the same field certain constituents of the soil essential to plant growth are removed, and the soil becomes impoverished and unproductive. To make the land once more fertile these constituents must be replaced. The calcium phosphate of the mineral deposits or of bone ash serves well as a material for restoring phosphorus to soils exhausted of that essential element; but a more soluble substance, which the plants can more readily assimilate, is desirable. It is better, therefore, to convert the insoluble calcium phosphate into the soluble primary phosphate before it is applied as fertilizer. It will be seen by reference to the formulas for the orthophosphates (see page 244) that in a primary phosphate only one hydrogen atom of phosphoric acid is replaced by a metal. Since the calcium atom always replaces two hydrogen atoms, it might be thought that there could be no primary calcium phosphate; but if the calcium atom replaces one hydrogen atom from each of two molecules of phosphoric acid, the salt Ca(H₂PO₄)₂ will result, and this is a primary phosphate. It can be made by treatment of the normal phosphate with the necessary amount of sulphuric acid, calcium sulphate being formed at the same time, thus:

Ca₃(PO₄)₂ + 2H₂SO₄ = Ca(H₂PO₄)₂ + 2CaSO₄.

The resulting mixture is a powder, which is sold as a fertilizer under the name of "superphosphate of lime."