CHAPTER XX

THE PHOSPHORUS FAMILY

================================================== ATOMIC MELTING SYMBOL WEIGHT DENSITY POINT —+—+—+—+— Phosphorus P 31.0 1.8 43.3 deg. Arsenic As 75.0 5.73 — Antimony Sb 120.2 6.7 432 deg. Bismuth Bi 208.5 9.8 270 deg. ==================================================

~The family.~ The elements constituting this family belong in the same group with nitrogen and therefore resemble it in a general way. They exhibit a regular gradation of physical properties, as is shown in the above table. The same general gradation is also found in their chemical properties, phosphorus being an acid-forming element, while bismuth is essentially a metal. The other two elements are intermediate in properties.

~Compounds.~ In general the elements of the family form compounds having similar composition, as is shown in the following table:

PH{3} PCl{3} PCl{5} P{2}O{3} P{2}O{5} AsH{3} AsCl{3} AsCl{5} As{2}O{3} As{2}O{5} SbH{3} SbCl{3} SbCl{5} Sb{2}O{3} Sb{2}O{5} .... BiCl{3} BiCl{5} Bi{2}O{3} Bi{2}O{5}

In the case of phosphorus, arsenic, and antimony the oxides are acid anhydrides. Salts of at least four acids of each of these three elements are known, the free acid in some instances being unstable. The relation of these acids to the corresponding anhydrides may be illustrated as follows, phosphorus being taken as an example:

P{2}O{3} + 3H{2}O = 2H{3}PO{3} (phosphorous acid).

P{2}O{5} + 3H{2}O = 2H{3}PO{4} (phosphoric acid).

P{2}O{5} + 2H{2}O = H{4}P{2}O{7} (pyrophosphoric acid).

P{2}O{5} + H{2}O = 2HPO{3} (metaphosphoric acid).

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.

~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:

Na{2}CO{3} + SiO{2} = Na{2}SiO{3} + CO{2}.

Silicon dioxide is a less volatile anhydride than phosphoric anhydride (P{2}O{5}), and when strongly heated with a phosphate the phosphoric anhydride is driven out, thus:

Ca{3}(PO{4}){2} + 3SiO{2} = 3CaSiO{3} + P{2}O{5}.

If carbon is added before the heat is applied, the P{2}O{5} is reduced to phosphorus at the same time, according to the equation

P{2}O{5} + 5C = 2P + 5CO.

~Physical properties.~ The purified phosphorus is a pale yellowish, translucent, waxy solid which melts at 43.3 deg. and boils at 269 deg.. 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 deg.-300 deg.). 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 (PH{3}) analogous to ammonia (NH{3}).

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

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 (P{2}O{5}) 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 (PH{4}Cl) analogous to ammonium chloride (NH{4}Cl). The phosphonium salts are of but little importance.

~Oxides of phosphorus.~ Phosphorus forms two well-known oxides,—the trioxide (P{2}O{3}) and the pentoxide (P{2}O{5}), 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:

H{3}PO{3} phosphorous acid. H{3}PO{4} phosphoric acid. H{4}P{2}O{7} pyrophosphoric acid. HPO{3} 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 (H{3}PO{3}). 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:

PCl{3} + 3H{2}O = H{3}PO{3} + 3HCl.

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

2. Orthophosphoric acid (phosphoric acid) (H{3}PO{4}). This acid can be obtained by dissolving phosphorus pentoxide in boiling water, as represented in the equation

P{2}O{5} + 3H{2}O = 2H{3}PO{4}.

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:

NaH{2}PO{4} monosodium hydrogen phosphate. Na{2}HPO{4} disodium hydrogen phosphate. Na{3}PO{4} 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(NH{4})HPO{4}.

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 (H{4}P{2}O{7}). On heating orthophosphoric acid to about 225 deg. pyrophosphoric acid is formed in accordance with the following equation:

2H{3}PO{4} = H{4}P{2}O{7} + H{2}O.

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

2Na{2}HPO{4} = Na{4}P{2}O{7} + H{2}O.

4. Metaphosphoric acid (glacial phosphoric acid) (HPO{3}). This acid is formed when orthophosphoric acid is heated above 400 deg.:

H{3}PO{4} = HPO{3} + H{2}O.

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

P{2}O{5} + H{2}O = 2HPO{3}.

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:

2Na{2}HPO{4} = Na{4}P{3}O{7} + H{2}O,

NaH{2}PO{4} = NaPO{3} + H{2}O.

~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(H{2}PO{4}){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:

Ca{3}(PO{4}){2} + 2H{2}SO{4} = Ca(H{2}PO{4}){2} + 2CaSO{4}.

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

ARSENIC

~Occurrence.~ Arsenic occurs in considerable quantities in nature as the native element, as the sulphides realgar (As{2}S{2}) and orpiment (As{2}S{3}), as oxide (As{2}O{3}), and as a constituent of many metallic sulphides, such as arsenopyrite (FeAsS).

~Preparation.~ The element is prepared by purifying the native arsenic, or by heating the arsenopyrite in iron tubes, out of contact with air, when the reaction expressed by the following equation occurs:

FeAsS = FeS + As.

The arsenic, being volatile, condenses in chambers connected with the heated tubes. It is also made from the oxide by reduction with carbon:

2As{2}O{3} + 3C = 4As + 3CO{2}.

~Properties.~ Arsenic is a steel-gray, metallic-looking substance of density 5.73. Though resembling metals in appearance, it is quite brittle, being easily powdered in a mortar. When strongly heated it sublimes, that is, it passes into a vapor without melting, and condenses again to a crystalline solid when the vapor is cooled. Like phosphorus it can be obtained in several allotropic forms. It alloys readily with some of the metals, and finds its chief use as an alloy with lead, which is used for making shot, the alloy being harder than pure lead. When heated on charcoal with the blowpipe it is converted into an oxide which volatilizes, leaving the charcoal unstained by any oxide coating. It burns readily in chlorine gas, forming arsenic trichloride,—

As + 3Cl = AsCl{3}.

Unlike most of its compounds, the element itself is not poisonous.

~Arsine~ (AsH{3}). When any compound containing arsenic is brought into the presence of nascent hydrogen, arsine (AsH{3}), corresponding to phosphine and ammonia, is formed. The reaction when oxide of arsenic is so treated is

As{2}O{3} + 12H = 2AsH{3} + 3H{2}O.

Arsine is a gas with a peculiar garlic-like odor, and is intensely poisonous. A single bubble of pure gas has been known to prove fatal. It is an unstable compound, decomposing into its elements when heated to a moderate temperature. It is combustible, burning with a pale bluish-white flame to form arsenic trioxide and water when air is in excess:

2AsH{3} + 6O = As{2}O{3} + 3H{2}O.

When the supply of air is deficient water and metallic arsenic are formed:

2AsH{3} + 3O = 3H{2}O + 2As.

These reactions make the detection of even minute quantities of arsenic a very easy problem.

~Marsh's test for arsenic.~ The method devised by Marsh for detecting arsenic is most frequently used, the apparatus being shown in Fig. 72. Hydrogen is generated in the flask A by the action of dilute sulphuric acid on zinc, is dried by passing over calcium chloride in the tube B, and after passing through the hard-glass tube C is ignited at the jet D. If a substance containing arsenic is now introduced into the generator A, the arsenic is converted into arsine by the action of the nascent hydrogen, and passes to the jet along with the hydrogen. If the tube C is strongly heated at some point near the middle, the arsine is decomposed while passing this point and the arsenic is deposited just beyond the heated point in the form of a shining, brownish-black mirror. If the tube is not heated, the arsine burns along with the hydrogen at the jet. Under these conditions a small porcelain dish crowded down into the flame is blackened by a spot of metallic arsenic, for the arsine is decomposed by the heat of the flame, and the arsenic, cooled below its kindling temperature by the cold porcelain, deposits upon it as a black spot. Antimony conducts itself in the same way as arsenic, but the antimony deposit is more sooty in appearance. The two can also be distinguished by the fact that sodium hypochlorite (NaClO) dissolves the arsenic deposit, but not that formed by antimony.

~Oxides of arsenic.~ Arsenic forms two oxides, As{2}O{3} and As{2}O{5}, corresponding to those of phosphorus. Of these arsenious oxide, or arsenic trioxide (As{2}O{3}), is much better known, and is the substance usually called white arsenic, or merely arsenic. It is found as a mineral, but is usually obtained as a by-product in burning pyrite in the sulphuric-acid industry. The pyrite has a small amount of arsenopyrite in it, and when this is burned arsenious oxide is formed as a vapor together with sulphur dioxide:

2FeAsS + 10O = Fe{2}O{3} + As{2}O{3} + 2SO{2}.

The arsenious oxide is condensed in appropriate chambers. It is a rather heavy substance, obtained either as a crystalline powder or as large, vitreous lumps, resembling lumps of porcelain in appearance. It is very poisonous, from 0.2 to 0.3 g. being a fatal dose. It is frequently given as a poison, since it is nearly tasteless and does not act very rapidly. This slow action is due to the fact that it is not very soluble, and hence is absorbed slowly by the system. Arsenious oxide is also used as a chemical reagent in glass making and in the dye industry.

~Acids of arsenic.~ Like the corresponding oxides of phosphorus, the oxides of arsenic are acid anhydrides. In solution they combine with bases to form salts, corresponding to the salts of the acids of phosphorus. Thus we have salts of the following acids:

H{3}AsO{3} arsenious acid.

H{3}AsO{4} orthoarsenic acid.

H{4}As{2}O{3} pyroarsenic acid.

HAsO{3} metarsenic acid.

Several other acids of arsenic are also known. Not all of these can be obtained as free acids, since they tend to lose water and form the oxides. Thus, instead of obtaining arsenious acid (H{3}AsO{3}), the oxide As{2}O{3} is obtained:

2H{3}AsO{3} = As{2}O{3} + 3H{2}O.

Salts of all the acids are known, however, and some of them have commercial value. Most of them are insoluble, and some of the copper salts, which are green, are used as pigments. Paris green, which has a complicated formula, is a well-known insecticide.

~Antidote for arsenical poisoning.~ The most efficient antidote for arsenic poisoning is ferric hydroxide. It is prepared as needed, according to the equation

Fe{2}(SO{4}){3} + 3Mg(OH){2} = 2Fe(OH){3} + 3MgSO{4}.

~Sulphides of arsenic.~ When hydrogen sulphide is passed into an acidified solution containing an arsenic compound the arsenic is precipitated as a bright yellow sulphide, thus:

2H{3}AsO{3} + 3H{2}S = As{2}S{3} + 6H{2}O,

2H{3}AsO{4} + 5H{2}S = As{2}S{5} + 8H{2}O.

In this respect arsenic resembles the metallic elements, many of which produce sulphides under similar conditions. The sulphides of arsenic, both those produced artificially and those found in nature, are used as yellow pigments.

ANTIMONY

~Occurrence.~ Antimony occurs in nature chiefly as the sulphide (Sb{2}S{3}), called stibnite, though it is also found as oxide and as a constituent of many complex minerals.

~Preparation.~ Antimony is prepared from the sulphide in a very simple manner. The sulphide is melted with scrap iron in a furnace, when the iron combines with the sulphur to form a slag, or liquid layer of melted iron sulphide, while the heavier liquid, antimony, settles to the bottom and is drawn off from time to time. The reaction involved is represented by the equation

Sb{2}S{3} + 3Fe = 2Sb + 3FeS.

~Physical properties.~ Antimony is a bluish-white, metallic-looking substance whose density is 6.7. It is highly crystalline, hard, and very brittle. It has a rather low melting point (432 deg.) and expands very noticeably on solidifying.

~Chemical properties.~ In chemical properties antimony resembles arsenic in many particulars. It forms the oxides Sb{2}O{3} and Sb{2}O{5}, and in addition Sb{2}O{4}. It combines with the halogen elements with great energy, burning brilliantly in chlorine to form antimony trichloride (SbCl{3}). When heated on charcoal with the blowpipe it is oxidized and forms a coating of antimony oxide on the charcoal which has a characteristic bluish-white color.

~Stibine~ (SbH{3}). The gas stibine (SbH{3}) is formed under conditions which are very similar to those which produce arsine, and it closely resembles the latter compound, though it is still less stable. It is very poisonous.

~Acids of antimony.~ The oxides Sb{2}O{3} and Sb{2}O{5} are weak acid anhydrides and are capable of forming two series of acids corresponding in formulas to the acids of phosphorus and arsenic. They are much weaker, however, and are of little practical importance.

~Sulphides of antimony.~ Antimony resembles arsenic in that hydrogen sulphide precipitates it as a sulphide when conducted into an acidified solution containing an antimony compound:

2SbCl{3} + 3H{2}S = Sb{2}S{3} + 6HCl,

2SbCl{5} + 5H{2}S = Sb{2}S{5} + 10HCl.

The two sulphides of antimony are called the trisulphide and the pentasulphide respectively. When prepared in this way they are orange-colored substances, though the mineral stibnite is black.

~Metallic properties of antimony.~ The physical properties of the element are those of a metal, and the fact that its sulphide is precipitated by hydrogen sulphide shows that it acts like a metal in a chemical way. Many other reactions show that antimony has more of the properties of a metal than of a non-metal. The compound Sb(OH){3}, corresponding to arsenious acid, while able to act as a weak acid is also able to act as a weak base with strong acids. For example, when treated with concentrated hydrochloric acid antimony chloride is formed:

Sb(OH){3} + 3HCl = SbCl{3} + 3H{2}O.

A number of elements act in this same way, their hydroxides under some conditions being weak acids and under others weak bases.

ALLOYS

Some metals when melted together thoroughly intermix, and on cooling form a homogeneous, metallic-appearing substance called an alloy. Not all metals will mix in this way, and in some cases definite chemical compounds are formed and separate out as the mixture solidifies, thus destroying the uniform quality of the alloy. In general the melting point of the alloy is below the average of the melting points of its constituents, and it is often lower than any one of them.

Antimony forms alloys with many of the metals, and its chief commercial use is for such purposes. It imparts to its alloys high density, rather low melting point, and the property of expanding on solidification. Such an alloy is especially useful in type founding, where fine lines are to be reproduced on a cast. Type metal consists of antimony, lead, and tin. Babbitt metal, used for journal bearings in machinery, contains the same metals in a different proportion together with a small percentage of copper.

BISMUTH

~Occurrence.~ Bismuth is usually found in the uncombined form in nature. It also occurs as oxide and sulphide. Most of the bismuth of commerce comes from Saxony, and from Mexico and Colorado, but it is not an abundant element.

~Preparation.~ It is prepared by merely heating the ore containing the native bismuth and allowing the melted metal to run out into suitable vessels. Other ores are converted into oxides and reduced by heating with carbon.

~Physical properties.~ Bismuth is a heavy, crystalline, brittle metal nearly the color of silver, but with a slightly rosy tint which distinguishes it from other metals. It melts at a low temperature (270 deg.) and has a density of 9.8. It is not acted upon by the air at ordinary temperatures.

~Chemical properties.~ When heated with the blowpipe on charcoal, bismuth gives a coating of the oxide Bi{2}O{3}. This has a yellowish-brown color which easily distinguishes it from the oxides formed by other metals. It combines very readily with the halogen elements, powdered bismuth burning readily in chlorine. It is not very easily acted upon by hydrochloric acid, but nitric and sulphuric acids act upon it in the same way that they do upon copper.

~Uses.~ Bismuth finds its chief use as a constituent of alloys, particularly in those of low melting point. Some of these melt in hot water. For example, Wood's metal, consisting of bismuth, lead, tin, and cadmium, melts at 60.5 deg..

~Compounds of bismuth.~ Unlike the other elements of this group, bismuth has almost no acid properties. Its chief oxide, Bi{2}O{3}, is basic in its properties. It dissolves in strong acids and forms salts of bismuth:

Bi{2}O{3} + 6HCl = 2BiCl{3} + 3H{2}O,

Bi{2}O{3} + 6HNO{3} = 2Bi(NO{3}){3} + 3H{2}O.

The nitrate and chloride of bismuth can be obtained as well-formed colorless crystals. When treated with water the salts are decomposed in the manner explained in the following paragraph.

HYDROLYSIS

Many salts such as those of antimony and bismuth form solutions which are somewhat acid in reaction, and must therefore contain hydrogen ions. This is accounted for by the same principle suggested to explain the fact that solutions of potassium cyanide are alkaline in reaction (p. 210). Water forms an appreciable number of hydrogen and hydroxyl ions, and very weak bases such as bismuth hydroxide are dissociated to but a very slight extent. When Bi+++ ions from bismuth chloride, which dissociates very readily, are brought in contact with the OH- ions from water, the two come to the equilibrium expressed in the equation

Bi+++ + 3OH- Bi(OH){3}.

For every hydroxyl ion removed from the solution in this way a hydrogen ion is left free, and the solution becomes acid in reaction.

Reactions of this kind and that described under potassium cyanide are called hydrolysis.

DEFINITION: Hydrolysis is the action of water upon a salt to form an acid and a base, one of which is very slightly dissociated.

~Conditions favoring hydrolysis.~ While hydrolysis is primarily due to the slight extent to which either the acid or the base formed is dissociated, several other factors have an influence upon the extent to which it will take place.

1. Influence of mass. Since hydrolysis is a reversible reaction, the relative masses of the reacting substances influence the point at which equilibrium will be reached. In the equilibrium

BiCl{3} + 3H{2}O Bi(OH){3} + 3HCl

the addition of more water will result in the formation of more bismuth hydroxide and hydrochloric acid. The addition of more hydrochloric acid will convert some of the bismuth hydroxide into bismuth chloride.

2. Formation of insoluble substances. When one of the products of hydrolysis is nearly insoluble in water the solution will become saturated with it as soon as a very little has been formed. All in excess of this will precipitate, and the reaction will go on until the acid set free increases sufficiently to bring about an equilibrium. Thus a considerable amount of bismuth and antimony hydroxides are precipitated when water is added to the chlorides of these elements. The greater the dilution the more hydroxide precipitates. The addition of hydrochloric acid in considerable quantity will, however, redissolve the precipitate.

~Partial hydrolysis.~ In many cases the hydrolysis of a salt is only partial, resulting in the formation of basic salts instead of the free base. Most of these basic salts are insoluble in water, which accounts for their ready formation. Thus bismuth chloride may hydrolyze by successive steps, as shown in the equations

BiCl{3} + H{2}O = Bi(OH)Cl{2} + HCl,

BiCl{3} + 2H{2}O = Bi(OH){2}Cl + 2HCl,

BiCl{3} + 3H{2}O = Bi(OH){3} + 3HCl.

The basic salt so formed may also lose water, as shown in the equation

Bi(OH){2}Cl = BiOCl + H{2}O.

The salt represented in the last equation is sometimes called bismuth oxychloride, or bismuthyl chloride. The corresponding nitrate, BiONO{3}, is largely used in medicine under the name of subnitrate of bismuth. In these two compounds the group of atoms, BiO, acts as a univalent metallic radical and is called bismuthyl. Similar basic salts are formed by the hydrolysis of antimony salts.

EXERCISES

1. Name all the elements so far studied which possess allotropic forms.

2. What compounds would you expect phosphorus to form with bromine and iodine? Write the equations showing the action of water on these compounds.

3. In the preparation of phosphine, why is coal gas passed into the flask? What other gases would serve the same purpose?

4. Give the formula for the salt which phosphine forms with hydriodic acid. Give the name of the compound.

5. Could phosphoric acid be substituted for sulphuric acid in the preparation of the common acids?

6. Write the equations for the preparation of the three sodium salts of orthophosphoric acid.

7. Why does a solution of disodium hydrogen phosphate react alkaline?

8. On the supposition that bone ash is pure calcium phosphate, what weight of it would be required in the preparation of 1 kg. of phosphorus?

9. If arsenopyrite is heated in a current of air, what products are formed?

10. (a) Write equations for the complete combustion of hydrosulphuric acid, methane, and arsine. (b) In what respects are the reactions similar?

11. Write the equations for all the reactions involved in Marsh's test for arsenic.

12. Write the names and formulas for the acids of antimony.

13. Write the equations showing the hydrolysis of antimony trichloride; of bismuth nitrate.

14. In what respects does nitrogen resemble the members of the phosphorus family?