CHAPTER XI

THE HALOGENS: CHLORINE, BROMINE, IODINE, AND FLUORINE

Although hydrochloric acid, like water, is one of the most stable substances, it is nevertheless decomposed not only by the action of a galvanic current,[1] but also by a high temperature. Sainte-Claire Deville showed that decomposition already occurs at 1,300°, because a cold tube (as with CO, Chapter IX.) covered with an amalgam of silver absorbs chlorine from hydrochloric acid in a red-hot tube, and the escaping gas contains hydrogen. V. Meyer and Langer (1885) observed the decomposition of hydrochloric acid at 1,690° in a platinum vessel; the decomposition in this instance was proved not only from the fact that hydrogen diffused through the platinum (p. 142), owing to which the volume was diminished, but also from chlorine being obtained in the residue (the hydrogen chloride was mixed with nitrogen), which liberated iodine from potassium iodide.[2] The usual method for the preparation of chlorine consists in the abstraction of the hydrogen by oxidising agents.[2 bis]

An aqueous solution of hydrochloric acid is generally employed for the evolution of chlorine. The hydrogen has to be abstracted from the hydrochloric acid. This is accomplished by nearly all oxidising substances, and especially by those which are able to evolve oxygen at a red heat (besides bases, such as mercury and silver oxides, which are able to give salts with hydrogen chloride); for example, manganese peroxide, potassium chlorate, chromic acid, &c. The decomposition essentially consists in the oxygen of the oxidising substance displacing the chlorine from 2HCl, forming water, H₂O, and setting the chlorine free, 2HCl + O (disengaged by the oxidising substances) = H₂O + Cl₂. Even nitric acid partially produces a like reaction; but as we shall afterwards see its action is more complicated, and it is therefore not suitable for the preparation of pure chlorine.[3] But other oxidising substances which do not give any other volatile products with hydrochloric acid may be employed for the preparation of chlorine. Among these may be mentioned: potassium chlorate, acid potassium chromate, sodium manganate, manganese peroxide, &c. Manganese peroxide is commonly employed in the laboratory, and on a large scale, for the preparation of chlorine. The chemical process in this case may be represented as follows: an exchange takes place between 4HCl and MnO₂, in which the manganese takes the place of the four atoms of hydrogen, or the chlorine and oxygen exchange places—that is, MnCl₄ and 2H₂O are produced. The chlorine compound, MnCl₄, obtained is very unstable; it splits up into chlorine, which as a gas passes from the sphere of action, and a lower compound containing less chlorine than the substance first formed, which remains in the apparatus in which the mixture is heated, MnCl₄ = MnCl₂ + Cl₂.[3 bis] The action of hydrochloric acid requires a temperature of about 100°. In the laboratory the preparation of chlorine is carried on in flasks, heated over a water-bath, by acting on manganese peroxide with hydrochloric acid or a mixture of common salt and sulphuric acid[4] and washing the gas with water to remove hydrochloric acid.[5] Chlorine cannot be collected over mercury, because it combines with it as with many other metals, and it is soluble in water; however, it is but slightly soluble in hot water or brine. Owing to its great weight, chlorine may be directly collected in a dry vessel by carrying the gas-conducting tube down to the bottom of the vessel. The chlorine will lie in a heavy layer at the bottom of the vessel, displace the air, and the extent to which it fills the vessel may be followed by its colour.[6]

Chlorine is a gas of a yellowish green colour, and has a very suffocating and characteristic odour. On lowering the temperature to -50° or increasing the pressure to six atmospheres (at 0°) chlorine condenses[7] into a liquid which has a yellowish-green colour, a density of 1·3, and boils at -34°. The density and atomic weight of chlorine is 35·5 times greater than that of hydrogen, hence the molecule contains Cl₂[8]. At 0° one volume of water dissolves about 1½ volume of chlorine, at 10° about 3 volumes, at 50° again 1½ volume.[9] Such a solution of chlorine is termed ‘chlorine water;’ and is employed in a diluted form in medicine and as a laboratory reagent. It is prepared by passing chlorine through a series of Woulfe's bottles or into an inverted retort filled with water. Under the action of light, chlorine water gives oxygen and hydrochloric acid. At 0° a saturated solution of chlorine yields a crystallo-hydrate, Cl₂,8H₂O, which easily splits up into chlorine and water when heated, so that if it be sealed up in a tube and heated to 35°, two layers of liquid are formed—a lower stratum of chlorine containing a small quantity of water, and an upper stratum of water containing a small quantity of chlorine.[10]

Chlorine explodes with hydrogen, if a mixture of equal volumes be exposed to the direct action of the sun's rays[11] or brought into contact with spongy platinum, or a strongly heated substance, or when subjected to the action of an electric spark. The explosion in this case takes place for exactly the same reasons—i.e. the evolution of heat and expansion of the resultant product—as in the case of detonating gas (Chapter III.) Diffused light acts in the same way, but slowly, whilst direct sunlight causes an explosion.[12] The hydrochloric acid gas produced by the reaction of chlorine on hydrogen occupies (at the original temperature and pressure) a volume equal to the sum of the original volumes; that is, a reaction of substitution here takes place: H₂ + Cl₂ = HCl + HCl. In this reaction twenty-two thousand heat units are evolved for one part by weight [1 gram] of hydrogen.[13]

These relations show that the affinity of chlorine for hydrogen is very great and analogous to the affinity between hydrogen and oxygen. Thus[14] on the one hand by passing a mixture of steam and chlorine through a red-hot tube, or by exposing water and chlorine to the sunlight, oxygen is disengaged, whilst on the other hand, as we saw above, oxygen in many cases displaces chlorine from its compound with hydrogen, and therefore the reaction H₂O + Cl₂ = 2HCl + O belongs to the number of reversible reactions, and hydrogen will distribute itself between oxygen and chlorine. This determines the relation of Cl to substances containing hydrogen and its reactions in the presence of water, to which we shall turn our attention after having pointed out the relation of chlorine to other elements.

Many metals when brought into contact with chlorine immediately combine with it, and form those metallic chlorides which correspond with hydrogen chloride and with the oxide of the metal taken. This combination may proceed rapidly with the evolution of heat and light; that is, metals are able to burn in chlorine. Thus, for example, sodium[15] burns in chlorine, synthesising common salt. Metals in the form of powders burn without the aid of heat, and become highly incandescent in the process; for instance, antimony, which is a metal easily converted into a powder.[16] Even such metals as gold and platinum,[17] which do not combine directly with oxygen and give very unstable compounds with it, unite directly with chlorine to form metallic chlorides. Either chlorine water or aqua regia may be employed for this purpose instead of gaseous chlorine. These dissolve gold and platinum, converting them into metallic chlorides. Aqua regia is a mixture of 1 part of nitric acid with 2 to 3 parts of hydrochloric acid. This mixture converts into soluble chlorides not only those metals which are acted on by hydrochloric and nitric acids, but also gold and platinum, which are insoluble in either acid separately. This action of aqua regia depends on the fact that nitric acid in acting on hydrochloric acid evolves chlorine. If the chlorine evolved be transferred to a metal, then a fresh quantity is formed from the remaining acids and also combines with the metal.[18] Thus the aqua regia acts by virtue of the chlorine which it contains and disengages.

The majority of non-metals also react directly on chlorine; hot sulphur and phosphorus burn in it and combine with it at the ordinary temperature. Only nitrogen, carbon, and oxygen do not combine directly with it. The chlorine compounds formed by the non-metals—for instance, phosphorus trichloride, PCl₃, and sulphurous chloride, &c., do not have the properties of salts, and, as we shall afterwards see more fully, correspond to acid anhydrides and acids; for example, PCl₃—to phosphorous acid, P(OH)₃:

As the above-mentioned relation in composition—i.e. substitution of Cl by the aqueous residue—exists between many chlorine compounds and their corresponding hydrates, and as furthermore some (acid) hydrates are obtained from chlorine compounds by the action of water, for instance,

whilst other chlorine compounds are formed from hydroxides and hydrochloric acid, with the liberation of water, for example,

NaHO + HCl = NaCl + H₂O

we endeavour to express this intimate connection between the hydrates and chlorine compounds by calling the latter chloranhydrides. In general terms, if the hydrate be basic, then,

and if the hydrate ROH be acid, then,

The chloranhydrides MCl corresponding to the bases are evidently metallic chlorides or salts corresponding to HCl. In this manner a distinct equivalency is marked between the compounds of chlorine and the so-called hydroxyl radicle (HO), which is also expressed in the analogy existing between chlorine, Cl₂, and hydrogen peroxide, (HO)₂.

As regards the chloranhydrides corresponding to acids and non-metals, they bear but little resemblance to metallic salts. They are nearly all volatile, and have a powerful suffocating smell which irritates the eyes and respiratory organs. They react on water like many anhydrides of the acids, with the evolution of heat and liberation of hydrochloric acid, forming acid hydrates. For this reason they cannot usually be obtained from hydrates—that is, acids—by the action of hydrochloric acid, as in that case water would be formed together with them, and water decomposes them, converting them into hydrates. There are many intermediate chlorine compounds between true saline metallic chlorides like sodium chloride and true acid chloranhydrides, just as there are all kinds of transitions between bases and acids. Acid chloranhydrides are not only obtained from chlorine and non-metals, but also from many lower oxides, by the aid of chlorine. Thus, for example, CO, NO, NO₂, SO₂, and other lower oxides which are capable of combining with oxygen may also combine with a corresponding quantity of chlorine. Thus COCl₂, NOCl, NO₂Cl, SO₂Cl₂, &c., are obtained. They correspond with the hydrates CO(OH)₂, NO(OH), NO₂(OH), SO₂(OH)₂, &c., and to the anhydrides CO₂, N₂O₃, N₂O₅, SO₃, &c. Here we should notice two aspects of the matter: (1) chlorine combines with that with which oxygen is able to combine, because it is in many respects equally if not more energetic than oxygen and replaces it in the proportion Cl₂: O; (2) that highest limit of possible combination which is proper to a given element or grouping of elements is very easily and often attained by combination with chlorine. If phosphorus gives PCl₃ and PCl₅, it is evident that PCl₅ is the higher form of combination compared with PCl₃. To the form PCl₅, or in general PX₅, correspond PH₄I, PO(OH)₃, POCl₃, &c. If chlorine does not always directly give compounds of the highest possible forms for a given element, then generally the lower forms combine with it in order to reach or approach the limit. This is particularly clear in hydrocarbons, where we see the limit C_nH_2n+2 very distinctly. The unsaturated hydrocarbons are sometimes able to combine with chlorine with the greatest ease and thus reach the limit. Thus ethylene, C₂H₄, combines with Cl₂, forming the so-called Dutch liquid or ethylene chloride, C₂H₄Cl₂, because it then reaches the limit C_nX_2n+2. In this and all similar cases the combined chlorine is able by reactions of substitution to give a hydroxide and a whole series of other derivatives. Thus a hydroxide called glycol, C₂H₄(OH)₂, is obtained from C₂H₄Cl₂.

Chlorine in the presence of water very often acts directly as an oxidising agent. A substance A combines with chlorine and gives, for example, ACl₂, and this in turn a hydroxide, A(OH)₂, which on losing water forms AO. Here the chlorine has oxidised the substance A. This frequently happens in the simultaneous action of water and chlorine: A + H₂O + Cl₂ = 2HCl + AO. Examples of this oxidising action of chlorine may frequently be observed both in practical chemistry and technical processes. Thus, for instance, chlorine in the presence of water oxidises sulphur and metallic sulphides. In this case the sulphur is converted into sulphuric acid, and the chlorine into hydrochloric acid, or a metallic chloride if a metallic sulphide be taken. A mixture of carbonic oxide and chlorine passed into water gives carbonic anhydride and hydrochloric acid. Sulphurous anhydride is oxidised by chlorine in the presence of water into sulphuric acid, just as it is by the action of nitric acid: SO₂ + 2H₂O + Cl₂ = H₂SO₄ + 2HCl.

The oxidising action of chlorine in the presence of water is taken advantage of in practice for the rapid bleaching of tissues and fibres. The colouring matter of the fibres is altered by oxidation and converted into a colourless substance, but the chlorine afterwards acts on the tissue itself. Bleaching by means of chlorine therefore requires a certain amount of technical skill in order that the chlorine should not act on the fibres themselves, but that its action should be limited to the colouring matter only. The fibre for making writing paper, for instance, is bleached in this manner. The bleaching property of chlorine was discovered by Berthollet, and forms an important acquisition to the arts, because it has in the majority of cases replaced that which before was the universal method of bleaching—namely, exposure to the sun of the fabrics damped with water, which is still employed for linens, &c. Time and great trouble, and therefore money also, have been considerably saved by this change.[19]

The power of chlorine for combination is intimately connected with its capacity for substitution, because, according to the law of substitution, if chlorine combines with hydrogen, then it also replaces hydrogen, and furthermore the combination and substitution are accomplished in the same quantities. Therefore the atom of chlorine which combines with the atom of hydrogen is also able to replace the atom of hydrogen. We mention this property of chlorine not only because it illustrates the application of the law of substitution in clear and historically important examples, but more especially because reactions of this kind explain those indirect methods of the formation of many substances which we have often mentioned and to which recourse is had in many cases in chemistry. Thus chlorine does not act on carbon,[20] oxygen, or nitrogen, but nevertheless its compounds with these elements may be obtained by the indirect method of the substitution of hydrogen by chlorine.

As chlorine easily combines with hydrogen, and does not act on carbon, it decomposes hydrocarbons (and many of their derivatives) at a high temperature, depriving them of their hydrogen and liberating the carbon, as, for example, is clearly seen when a lighted candle is placed in a vessel containing chlorine. The flame becomes smaller, but continues to burn for a certain time, a large amount of soot is obtained, and hydrochloric acid is formed. In this case the gaseous and incandescent substances of the flame are decomposed by the chlorine, the hydrogen combines with it, and the carbon is disengaged as soot.[21] This action of chlorine on hydrocarbons, &c., proceeds otherwise at lower temperatures, as we will now consider.

A very important epoch in the history of chemistry was inaugurated by the discovery of Dumas and Laurent that chlorine is able to displace and replace hydrogen. This discovery is important from the fact that chlorine proved to be an element which combines with great ease simultaneously with both the hydrogen and the element with which the hydrogen was combined. This clearly proved that there is no opposite polarity between elements forming stable compounds. Chlorine does not combine with hydrogen because it has opposite properties, as Dumas and Laurent stated previously, accounting hydrogen to be electro-positive and chlorine electro-negative; this is not the reason of their combining together, for the same chlorine which combines with hydrogen is also able to replace it without altering many of the properties of the resultant substance. This substitution of hydrogen by chlorine is termed metalepsis. The mechanism of this substitution is very constant. If we take a hydrogen compound, preferably a hydrocarbon, and if chlorine acts directly on it, then there is produced on the one hand hydrochloric acid and on the other hand a compound containing chlorine in the place of the hydrogen—so that the chlorine divides itself into two equal portions, one portion is evolved as hydrochloric acid, and the other portion takes the place of the hydrogen thus liberated. Hence this metalepsis is always accompanied by the formation of hydrochloric acid.[22] The scheme of the process is as follows:

Or, in general terms—

The conditions under which metalepsis takes place are also very constant. In the dark chlorine does not usually act on hydrogen compounds, but the action commences under the influence of light. The direct action of the sun's rays is particularly propitious to metalepsis. It is also remarkable that the presence of traces of certain substances,[23] especially of iodine, aluminium chloride, antimony chloride, &c., promotes the action. A trace of iodine added to the substance subjected to metalepsis often produces the same effect as sunlight.[24]

If marsh gas be mixed with chlorine and the mixture ignited, then the hydrogen is entirely taken up from the marsh gas and hydrochloric acid and carbon formed, but there is no metalepsis.[25] But if a mixture of equal volumes of chlorine and marsh gas be exposed to the action of diffused light, then the greenish yellow mixture gradually becomes colourless, and hydrochloric acid and the first product of metalepsis—namely, methyl chloride—are formed:

The volume of the mixture remains unaltered. The methyl chloride which is formed is a gas. If it be separated from the hydrochloric acid (it is soluble in acetic acid, in which hydrochloric acid is but sparingly soluble) and be again mixed with chlorine, then it may be subjected to a further metalepsical substitution—the second atom of hydrogen may be substituted by chlorine, and a liquid substance, CH₂Cl₂, called methylene chloride, will be obtained. In the same manner the substitution may be carried on still further, and CHCl₃, or chloroform, and lastly carbon tetrachloride, CCl₄, will be produced. Of these substances the best known is chloroform, owing to its being formed from many organic substances (by the action of bleaching powder) and to its being used in medicine as an anæsthetic; chloroform boils at 62° and carbon tetrachloride at 78°. They are both colourless odoriferous liquids, heavier than water. The progressive substitution of hydrogen by chlorine is thus evident, and it can be clearly seen that the double decompositions are accomplished between molecular quantities of the substance—that is, between equal volumes in a gaseous state.

Carbon tetrachloride, which is obtained by the metalepsis of marsh gas, cannot be obtained directly from chlorine and carbon, but it may be obtained from certain compounds of carbon—for instance, from carbon bisulphide—if its vapour mixed with chlorine be passed through a red-hot tube. Both the sulphur and carbon then combine with the chlorine. It is evident that by ultimate metalepsis a corresponding carbon chloride may be obtained from any hydrocarbon—indeed, the number of chlorides of carbon C_nCl_2m already known is very large.

As a rule, the fundamental chemical characters of hydrocarbons are not changed by metalepsis; that is, if a neutral substance be taken, then the product of metalepsis is also a neutral substance, or if an acid be taken the product of metalepsis also has acid properties. Even the crystalline form not unfrequently remains unaltered after metalepsis. The metalepsis of acetic acid, CH₃·COOH, is historically the most important. It contains three of the atoms of the hydrogen of marsh gas, the fourth being replaced by carboxyl, and therefore by the action of chlorine it gives three products of metalepsis (according to the amount of the chlorine and conditions under which the reaction takes place), mono-, di-, and tri-chloracetic acids—CH₂Cl·COOH, CHCl₂·COOH, and CCl₃·COOH; they are all, like acetic acid, monobasic. The resulting products of metalepsis, in containing an element which so easily acts on metals as chlorine, possess the possibility of attaining a further complexity of molecules of which the original hydrocarbon is often in no way capable. Thus on treating with an alkali (or first with a salt and then with an alkali, or with a basic oxide and water, &c.) the chlorine forms a salt with its metal, and the hydroxyl radicle takes the place of the chlorine—for example, CH₃·OH is obtained from CH₃Cl. By the action of metallic derivatives of hydrocarbons—for example, CH₃Na—the chlorine also gives a salt, and the hydrocarbon radicle—for instance, CH₃—takes the place of the chlorine. In this, or in a similar manner, CH₃·CH₃, or C₂H₆ is obtained from CH₃Cl and C₆H₅·CH₃ from C₆H₆. The products of metalepsis also often react on ammonia, forming hydrochloric acid (and thence NH₄Cl) and an amide; that is, the product of metalepsis, with the ammonia radicle NH₂, &c. in the place of chlorine. Thus by means of metalepsical substitution methods were found in chemistry for an artificial and general means of the formation of complex carbon compounds from more simple compounds which are often totally incapable of direct reaction. Besides which, this key opened the doors of that secret edifice of complex organic compounds into which man had up to then feared to enter, supposing the hydrocarbon elements to be united only under the influence of those mystic forces acting in organisms.[26]

It is not only hydrocarbons which are subject to metalepsis. Certain other hydrogen compounds, under the action of chlorine, also give corresponding chlorine derivatives in exactly the same manner; for instance, ammonia, caustic potash, caustic lime, and a whole series of alkaline substances.[27] In fact, just as the hydrogen in marsh gas can be replaced by chlorine and form methyl chloride, so the hydrogen in caustic potash, KHO, ammonia, NH₃, and calcium hydroxide, CaH₂O₂ or Ca(OH)₂, may be replaced by chlorine and give potassium hypochlorite, KClO, calcium hypochlorite, CaCl₂O₂, and the so-called chloride of nitrogen, NCl₃. For not only is the correlation in composition the same as in the substitution in marsh gas, but the whole mechanism of the reaction is the same. Here also two atoms of chlorine act: one takes the place of the hydrogen whilst the other is evolved as hydrochloric acid, only in the former case the hydrochloric acid evolved remained free, and in the latter, in presence of alkaline substances, it reacts on them. Thus, in the action of chlorine on caustic potash, the hydrochloric acid formed acts on another quantity of caustic potash and gives potassium chloride and water, and therefore not only KHO + Cl₂ = HCl + KClO, but also KHO + HCl = H₂O + KCl, and the result of both simultaneous phases will be 2KHO + Cl₂ = H₂O + KCl + KClO. We will here discuss certain special cases.

The action of chlorine on ammonia may either result in the entire breaking up of the ammonia, with the evolution of gaseous nitrogen, or in a product of metalepsis (as with CH₄). With an excess of chlorine and the aid of heat the ammonia is decomposed, with the disengagement of free nitrogen.[28] This reaction evidently results in the formation of sal-ammoniac, 8NH₃ + 3Cl₂ = 6NH₄Cl + N₂. But if the ammonium salt be in excess, then the reaction takes the direction of the replacement of the hydrogen in the ammonia by chlorine. The principal result is that NH₃ + 3Cl₂ forms NCl₃ + 3HCl.[29] The resulting product of metalepsis, or chloride of nitrogen, NCl₃, discovered by Dulong, is a liquid having the property of decomposing with excessive ease not only when heated, but even under the action of mechanical influences, as by a blow or by contact with certain solid substances. The explosion which accompanies the decomposition is due to the fact that the liquid chloride of nitrogen gives gaseous products, nitrogen and chlorine.[29 bis]

Chloride of nitrogen is a yellow oily liquid of sp. gr. 1·65, which boils at 71°, and breaks up into N + Cl₃ at 97°. The contact of phosphorus, turpentine, india-rubber, &c. causes an explosion, which is sometimes so violent that a small drop will pierce through a thick board. The great ease with which chloride of nitrogen decomposes is dependent upon the fact that it is formed with an absorption of heat, which it evolves when decomposed, to the amount of about 38,000 heat units for NCl₃, as Deville and Hautefeuille determined.

Chlorine, when absorbed by a solution of caustic soda (and also of other alkalis) at the ordinary temperature, causes the replacement of the hydrogen in the caustic soda by the chlorine, with the formation of sodium chloride by the hydrochloric acid, so that the reaction may be represented in two phases, as described above. In this manner, sodium hypochlorite, NaClO, and sodium chloride are simultaneously formed: 2NaHO + Cl₂ = NaCl + NaClO + H₂O. The resultant solution contains NaClO and is termed ‘eau de Javelle.’ An exactly similar reaction takes place when chlorine is passed over dry hydrate of lime at the ordinary temperature: 2Ca(HO)₂ + 2Cl₂ = CaCl₂O₂ + CaCl₂ + 2H₂O. A mixture of the product of metalepsis with calcium chloride is obtained. This mixture is employed in practice on a large scale, and is termed ‘bleaching powder,’ owing to its acting, especially when mixed with acids, as a bleaching agent on tissues, so that it resembles chlorine in this respect. It is however preferable to chlorine, because the destructive action of the chlorine can be moderated in this case, and because it is much more convenient to deal with a solid substance than with gaseous chlorine. Bleaching powder is also called chloride of lime, because it is obtained from chlorine and hydrate of lime, and contains[30] both these substances. It may be prepared in the laboratory by passing a current of chlorine through a cold mixture of water and lime (milk of lime). The mixture must be kept cold, as otherwise 3Ca(ClO)₂ passes into 2CaCl₂ + Ca(ClO₃)₂. In the manufacture of bleaching powder in large quantities at chemical works, the purest possible slaked lime is taken and laid in a thin layer in large flat chambers, M (whose walls are made of Yorkshire flags or tarred wood, on which chlorine has no action), and into which chlorine gas is introduced by lead tubes. The distribution of the plant is shown in the annexed drawing (fig. 67).

The products of the metalepsis of alkaline hydrates, NaClO and Ca(ClO)₂, which are present in solutions of ‘Javelle salt’ and bleaching powder (they are not obtained free from metallic chlorides), must be counted as salts, because their metals are capable of substitution. But the hydrate HClO corresponding with these salts, or hypochlorous acid, is not obtained in a free or pure state, for two reasons: in the first place, because this hydrate, as a very feeble acid, splits up (like H₂CO₃ or HNO₃) into water and the anhydride, or chlorine monoxide, Cl₂O = 2HClO - H₂O; and, in the second place, because, in a number of instances, it evolves oxygen with great facility, forming hydrochloric acid: HClO = HCl + O. Both hypochlorous acid and chlorine monoxide may be regarded as products of the metalepsis of water, because HOH corresponds with ClOH and ClOCl. Hence in many instances bleaching salts (a mixture of hypochlorites and chlorides) break up, with the evolution of (1) chlorine, under the action of an excess of a powerful acid capable of evolving hydrochloric acid from sodium or calcium chlorides, and this takes place most simply under the action of hydrochloric acid itself, because (p. 462) NaCl + NaClO + 3HCl = 2NaCl + HCl + Cl₂ + H₂O; (2) oxygen, as we saw in Chapter III.—The bleaching properties and, in general, oxidising action of bleaching salts is based on this evolution of oxygen (or chlorine); oxygen is also disengaged on heating the dry salts—for instance, NaCl + NaClO = 2NaCl + O; (3) and, lastly, chlorine monoxide, which contains both chlorine and oxygen. Thus, if a little sulphuric, nitric, or similar acid (not enough to liberate hydrochloric acid from the CaCl₂) be added to a solution of a bleaching salt (which has an alkaline reaction, owing either to an excess of alkali or to the feeble acid properties of HClO), then the hypochlorous acid set free gives water and chlorine monoxide. If carbonic anhydride (or boracic or a similar very feeble acid) act on the solution of a bleaching salt, then hydrochloric acid is not evolved from the sodium or calcium chlorides, but the hypochlorous acid is displaced and gives chlorine monoxide,[31] because hypochlorous acid is one of the most feeble acids. Another method for the preparation of chlorine monoxide is based on these feeble acid properties of hypochlorous acid. Zinc oxide and mercury oxide, under the action of chlorine in the presence of water, do not give a salt of hypochlorous acid, but form a chloride and hypochlorous acid, which fact shows the incapacity of this acid to combine with the bases mentioned. Therefore, if such oxides as those of zinc or mercury be shaken up in water, and chlorine be passed through the turbid liquid,[32] a reaction occurs which may be expressed in the following manner: 2HgO + 2Cl₂ = Hg₂OCl₂ + Cl₂O. In this case, a compound of mercury oxide with mercury chloride, or the so-called mercury oxychloride, is obtained: Hg₂OCl₂ = HgO + HgCl₂. This is insoluble in water, and is not affected by hypochlorous anhydride, so that the solution will contain hypochlorous acid only, but the greater part of it splits up into the anhydride and water.[32 bis]

Chlorine monoxide, which corresponds to bleaching and hypochlorous salts, containing as it does the two elements oxygen and chlorine, forms a characteristic example of a compound of elements which, in the majority of cases, act chemically in an analogous manner. Chlorine monoxide, as prepared from an aqueous solution by the abstraction of water or by the action of dry chlorine on cold mercury oxide, is, at the ordinary temperature, a gas or vapour which condenses into a red liquid boiling at +20° and giving a vapour whose density (43 referred to hydrogen) shows that 2 vols. of chlorine and 1 vol. of oxygen give 2 vols. of chlorine monoxide. In an anhydrous form the gas or liquid easily explodes, splitting up into chlorine and oxygen. This explosiveness is determined by the fact that heat is evolved in the decomposition to the amount of about 15,000 heat units for Cl₂O.[33] The explosion may even take place spontaneously, and also in the presence of many oxidisable substances (for instance, sulphur, organic compounds, &c.), but the solution, although unstable and showing a strong oxidising tendency, does not explode.[34] It is evident that the presence of hypochlorous acid, HClO, may be assumed in an aqueous solution of Cl₂O, since Cl₂O + H₂O = 2HClO.

Hypochlorous acid, its salts, and chlorine monoxide serve as a transition between hydrochloric acid, chlorides, and chlorine, and a whole series of compounds containing the same elements combined with a still greater quantity of oxygen. The higher oxides of chlorine, as their origin indicates, are closely connected with hypochlorous acid and its salts:

When heated, solutions of hypochlorites undergo a remarkable change. Themselves so unstable, they, without any further addition, yield two fresh salts which are both much more stable; one contains more oxygen than MClO, the other contains none at all.

Part of the salt—namely, two-thirds of it—parts with its oxygen in order to oxidise the remaining third.[36] From an intermediate substance, RX, two extremes, R and RX₃ are formed, just as nitrous anhydride splits up into nitric oxide and nitric anhydride (or nitric acid). The resulting salt, MClO₃, corresponds with chloric acid and potassium chlorate, KClO₃. It is evident that a similar salt may be obtained directly by the action of chlorine on an alkali if its solution be heated, because RClO will be first formed, and then RClO₃; for example, 6KHO + 3Cl₂ = KClO₃ + 5KCl + 3H₂O. Chlorates are so prepared; for instance, potassium chlorate, which is easily separated from potassium chloride, being sparingly soluble in cold water.[37]

If dilute sulphuric acid be added to a solution of potassium chlorate, chloric acid is liberated, but it cannot be separated by distillation, as it is decomposed in the process. To obtain the free acid, sulphuric acid must be added to a solution of barium chlorate.[38] The sulphuric acid gives a precipitate of barium sulphate, and free chloric acid remains in solution. The solution may be evaporated under the receiver of an air-pump. This solution is colourless, has no smell, and acts as a powerful acid (it neutralises sodium hydroxide, decomposes sodium carbonate, gives hydrogen with zinc, &c.); when heated above 40°, however, it decomposes, forming chlorine, oxygen, and perchloric acid: 4HClO₃ = 2HClO₄ + H₂O + Cl₂ + O₃. In a concentrated condition the acid acts as an exceedingly energetic oxidiser, so that organic substances brought into contact with it burst into flame. Iodine, sulphurous acid, and similar oxidisable substances form higher oxidation products and reduce the chloric acid to hydrochloric acid. Hydrochloric acid gas gives chlorine with chloric acid (and consequently with KClO₃ also) acting in the same manner as it acts on the lower acids: HClO₃ + 5HCl = 3H₂O + 3Cl₂.

By cautiously acting on potassium chlorate with sulphuric acid, the dioxide (chloric peroxide), ClO₂,[39] is obtained (Davy, Millon). This gas is easily liquefied in a freezing mixture, and boils at +10°. The vapour density (about 35 if H = 1) shows that the molecule of this substance is ClO₂.[40] In a gaseous or liquid state it very easily explodes (for instance, at 60°, or by contact with organic compounds or finely divided substances, &c.), forming Cl and O₂, and in many instances[41] therefore it acts as an oxidising agent, although (like nitric peroxide) it may itself be further oxidised.[42] In dissolving in water or alkalis chloric peroxide gives chlorous and hypochlorous acids—2ClO₂ + 2KHO = KClO₃ + KClO₂ + H₂O—and therefore, like nitric peroxide, the dioxide may be regarded as an intermediate oxide between the (unknown) anhydrides of chlorous and chloric acids: 4ClO₂ = Cl₂O₃ + Cl₂O₃.[43]

As the salts of chloric acid, HClO₃, are produced by the splitting up of the salts of hypochlorous acid, so in the same way the salts of perchloric acid, HClO₄, are produced from the salts of chloric acid, HClO₃. But this is the highest form of the oxidation of HCl. Perchloric acid, HClO₄, is the most stable of all the acids of chlorine. When fused potassium chlorate begins to swell up and solidify, after having parted with one-third of its oxygen, potassium chloride and potassium perchlorate have been formed according to the equation 2KClO₃ = KClO₄ + KCl + O₂.

The formation of this salt is easily observed in the preparation of oxygen from potassium chlorate, owing to the fact that the potassium perchlorate fuses with greater difficulty than the chlorate, and therefore appears in the molten salt as solid grains (see Chapter III. Note 12). Under the action of certain acids—for instance, sulphuric and nitric—potassium chlorate also gives potassium perchlorate. This latter may be easily purified, because it is but sparingly soluble in water, although all the other salts of perchloric acid are very soluble and even deliquesce in the air. The perchlorates, although they contain more oxygen than the chlorates, are decomposed with greater difficulty, and even when thrown on ignited charcoal give a much feebler deflagration than the chlorates. Sulphuric acid (at a temperature not below 100°) evolves volatile and to a certain extent stable perchloric acid from potassium perchlorate. Neither sulphuric nor any other acid will further decompose perchloric acid as it decomposes chloric acid. Of all the acids of chlorine, perchloric acid alone can be distilled.[44] The pure hydrate HClO₄[45] is a colourless and exceedingly caustic substance which fumes in the air and has a specific gravity 1·78 at 15° (sometimes, after being kept for some time, it decomposes with a violent explosion). It explodes violently when brought into contact with charcoal, paper, wood, and other organic substances. If a small quantity of water be added to this hydrate, and it be cooled, a crystallo-hydrate, ClHO₄,H₂O, separates out. This is much more stable, but the liquid hydrate HClO₄,2H₂O is still more so. The acid dissolves in water in all proportions, and its solutions are distinguished for their stability.[46] When ignited both the acid and its salts are decomposed, with the evolution of oxygen.[47]

On comparing chlorine as an element not only with nitrogen and carbon but with all the other non-metallic elements (chlorine has so little analogy with the metals that a comparison with them would be superfluous), we find in it the following fundamental properties of the halogens or salt-producers. With metals chlorine gives salts (such as sodium chloride, &c.); with hydrogen a very energetic and monobasic acid HCl, and the same quantity of chlorine is able by metalepsis to replace the hydrogen; with oxygen it forms unstable oxides of an acid character. These properties of chlorine are possessed by three other elements, bromine, iodine, and fluorine. They are members of one natural family. Each representative has its peculiarities, its individual properties and points of distinction, in combination and in the free state—otherwise they would not be independent elements; but the repetition in all of them of the same chief characteristics of the family enables one more quickly to grasp all their various properties and to classify the elements themselves.

In order to have a guiding thread in forming comparisons between the elements, attention must however be turned not only to their points of resemblance but also to those of their properties and characters in which they differ most from each other. And the atomic weights of the elements must be considered as their most elementary property, since this is a quantity which is most firmly established, and must be taken account of in all the reactions of the element. The halogens have the following atomic weights—

F = 19, Cl = 35·5, Br = 80, I = 127.

All the properties, physical and chemical, of the elements and their corresponding compounds must evidently be in a certain dependence on this fundamental point, if the grouping in one family be natural.[47 bis] And we find in reality that, for instance, the properties of bromine, whose atomic weight is almost the mean between those of iodine and chlorine, occupy a mean position between those of these two elements. The second measurable property of the elements is their equivalence or their capacity for forming compounds of definite forms. Thus carbon or nitrogen in this respect differs widely from the halogens. Although the form ClO₂ corresponds with NO₂ and CO₂, yet the last is the highest oxide of carbon, whilst that of nitrogen is N₂O₅, and for chlorine, if there were an anhydride of perchloric acid, its composition would be Cl₂O₇, which is quite different from that of carbon. In respect to the forms of their compounds the halogens, like all elements of one family or group, are perfectly analogous to each other, as is seen from their hydrogen compounds:

HF, HCl, HBr, HI.

Their oxygen compounds exhibit a similar analogy. Only fluorine does not give any oxygen compounds. The iodine and bromine compounds corresponding with HClO₃ and HClO₄ are HBrO₃ and HBrO₄, HIO₃ and HIO₄. On comparing the properties of these acids we can even predict that fluorine will not form any oxygen compound. For iodine is easily oxidised—for instance, by nitric acid—whilst chlorine is not directly oxidised. The oxygen acids of iodine are comparatively more stable than those of chlorine; and, generally speaking, the affinity of iodine for oxygen is much greater than that of chlorine. Here also bromine occupies an intermediate position. In fluorine we may therefore expect a still smaller affinity for oxygen than in chlorine—and up to now it has not been combined with oxygen. If any oxygen compounds of fluorine should be obtained, they will naturally be exceedingly unstable. The relation of these elements to hydrogen is the reverse of the above. Fluorine has so great an affinity for hydrogen that it decomposes water at the ordinary temperature; whilst iodine has so little affinity for hydrogen that hydriodic acid, HI, is formed with difficulty, is easily decomposed, and acts as a reducing agent in a number of cases.

From the form of their compounds the halogens are univalent elements with respect to hydrogen and septivalent with respect to oxygen, N being trivalent to hydrogen (it gives NH₃) and quinqui-valent to oxygen (it gives N₂O₅), and C being quadrivalent to both H and O as it forms CH₄ and CO₂. And as not only their oxygen compounds, but also their hydrogen compounds, have acid properties, the halogens are elements of an exclusively acid character. Such metals as sodium, potassium, barium only give basic oxides. In the case of nitrogen, although it forms acid oxides, still in ammonia we find that capacity to give an alkali with hydrogen which indicates a less distinctly acid character than in the halogens. In no other elements is the acid-giving property so strongly developed as in the halogens.

In describing certain peculiarities characterising the halogens, we shall at every step encounter a confirmation of the above-mentioned general relations.

As fluorine decomposes water with the evolution of oxygen, F₂ + H₂O = 2HF + O, for a long time all efforts to obtain it in free state by means of methods similar to those for the preparation of chlorine proved fruitless.[48] Thus by the action of hydrofluoric acid on manganese peroxide, or by decomposing a solution of hydrofluoric acid by an electric current, either oxygen or a mixture of oxygen and fluorine were obtained instead of fluorine. Probably a certain quantity of fluorine[48 bis] was set free by the action of oxygen or an electric current on incandescent and fused calcium fluoride, but at a high temperature fluorine acts even on platinum, and therefore it was not obtained. When chlorine acted on silver fluoride, AgF, in a vessel of natural fluor spar, CaF₂, fluorine was also liberated; but it was mixed with chlorine, and it was impossible to study the properties of the resultant gas. Brauner (1881) also obtained fluorine by igniting cerium fluoride, 2CeF₄ = 2CeF₃ + F₂; but this, like all preceding efforts, only showed fluorine to be a gas which decomposes water, and is capable of acting in a number of instances like chlorine, but gave no possibility of testing its properties. It was evident that it was necessary to avoid as far as possible the presence of water and a rise of temperature; this Moissan succeeded in doing in 1886. He decomposed anhydrous hydrofluoric acid, liquefied at a temperature of -23° and contained in a U-shaped tube (to which a small quantity of potassium fluoride had been added to make it a better conductor), by the action of a powerful electric current (twenty Bunsen's elements in series). Hydrogen was then evolved at the negative pole, and fluorine appeared at the positive pole (of iridium platinum) as a pale green gas which decomposed water with the formation of ozone and hydrofluoric acid, and combined directly with silicon (forming silicon fluoride, SiF₄), boron (forming BF₃), sulphur, &c. Its density (H = 1) is 18, so that its molecule is F₂. But the action of fluorine on metals at the ordinary temperature is comparatively feeble, because the metallic fluoride formed coats the remaining mass of the metals; it is, however, completely absorbed by iron. Hydrocarbons (such as naphtha), alcohol, &c., immediately absorb fluorine, with the formation of hydrofluoric acid. Fluorine when mixed with hydrogen can easily be made to explode violently, forming hydrofluoric acid.[49]

In 1894, Brauner obtained fluorine directly by igniting the easily formed[49 bis] double lead salt HF,3KF,PbF₄, which first, at 230°, decomposes with the evolution of HF, and then splits up forming 3KF,PbF₂ and fluorine F₂, which is recognised by the fact that it liberates iodine from KI and easily combines with silicon, forming SiF₄. This method gives chemically pure fluorine, and is based upon the breaking up of the higher compound—tetrafluoride of lead, PbF₄, corresponding to PbO₂, into free fluorine, F₂, and the lower more stable form—bifluoride of lead, PbF₂, which corresponds to PbO; that is, this method resembles the ordinary method of obtaining chlorine by means of MnO₂, as MnCl₄ here breaks up into MnCl₂ and chlorine, just as PbF₄ splits up into PbF₂ and fluorine.

Among the compounds of fluorine, calcium fluoride, CaF₂, is somewhat widely distributed in nature as fluor spar,[50] whilst cryolite, or aluminium sodium fluoride, Na₃AlF₆, is found more rarely (in large masses in Greenland). Cryolite, like fluor spar, is also insoluble in water, and gives hydrofluoric acid with sulphuric acid. Small quantities of fluorine have also in a number of cases been found in the bodies of animals, in the blood, urine, and bones. If fluorides occur in the bodies of animals, they must have been introduced in food, and must occur in plants and in water. And as a matter of fact river, and especially sea, water always contains a certain, although small, quantity of fluorine compounds.

Hydrofluoric acid, HF, cannot be obtained from fluor spar in glass retorts, because glass is acted on by and destroys the acid. It is prepared in lead vessels, and when it is required pure, in platinum vessels, because lead also acts on hydrofluoric acid, although only very feebly on the surface, and when once a coating of fluoride and sulphate of lead is formed no further action takes place. Powdered fluor spar and sulphuric acid evolve hydrofluoric acid (which fumes in the air) even at the ordinary temperature, CaF₂ + H₂SO₄ = CaSO₄ + 2HF. At 130° fluor spar is completely decomposed by sulphuric acid. The acid is then evolved as vapour, which may be condensed by a freezing mixture into an anhydrous acid. The condensation is aided by pouring water into the receiver of the condenser, as the acid is easily soluble in cold water.

In the liquid anhydrous form hydrofluoric acid boils at +19°, and its specific gravity at 12·8° = 0·9849.[51] It dissolves in water with the evolution of a considerable amount of heat, and gives a solution of constant boiling point which distils over at 120°; showing that the acid is able to combine with water. The specific gravity of the compound is 1·15, and its composition HF,2H₂O.[52] With an excess of water a dilute solution distils over first. The aqueous solution and the acid itself must be kept in platinum vessels, but the dilute acid may be conveniently preserved in vessels made of various organic materials, such as gutta-percha, or even in glass vessels having an interior coating of paraffin. Hydrofluoric acid does not act on hydrocarbons and many other substances, but it acts in a highly corrosive manner on metals, glass, porcelain, and the majority of rock substances.[53] It also attacks the skin, and is distinguished by its poisonous properties, so that in working with the acid a strong draught must be kept up, to prevent the possibility of the fumes being inhaled. The non-metals do not act on hydrofluoric acid, but all metals—with the exception of mercury, silver, gold, and platinum, and, to a certain degree, lead—decompose it with the evolution of hydrogen. With bases it gives directly metallic fluorides, and behaves in many respects like hydrochloric acid. There are, however, several distinct individual differences, which are furthermore much greater than those between hydrochloric, hydrobromic, and hydriodic acids. Thus the silver compounds of the latter are insoluble in water, whilst silver fluoride is soluble. Calcium fluoride, on the contrary, is insoluble in water, whilst calcium chloride, bromide, and iodide are not only soluble, but attract water with great energy. Neither hydrochloric, hydrobromic, nor hydriodic acid acts on sand and glass, whilst hydrofluoric acid corrodes them, forming gaseous silicon fluoride. The other halogen acids only form normal salts, KCl, NaCl, with Na or K, whilst hydrofluoric acid gives acid salts, for instance HKF₂ (and by dissolving KF in liquid HF, KHF₂2HF is obtained). This latter property is in close connection with the fact that at the ordinary temperature the vapour density of hydrofluoric acid is nearly 20, which corresponds with a formula H₂F₂, as Mallet (1881) showed; but a depolymerisation occurs with a rise of temperature, and the density approaches 10, which answers to the formula HF.[54]

The analogy between chlorine and the other two halogens, bromine and iodine, is much more perfect. Not only have their hydrates or halogen acids much in common, but they themselves resemble chlorine in many respects,[55] and even the properties of the corresponding metallic compounds of bromine and iodine are very much alike. Thus, the chlorides, bromides, and iodides of sodium and potassium crystallise in the cubic system, and are soluble in water; the chlorides of calcium, aluminium, magnesium, and barium are just as soluble in water as the bromides and iodides of these metals. The iodides and bromides of silver and lead are sparingly soluble in water, like the chlorides of these metals. The oxygen compounds of bromine and iodine also present a very strong analogy to the corresponding compounds of chlorine. A hypobromous acid is known corresponding with hypochlorous acid. The salts of this acid have the same bleaching property as the salts of hypochlorous acid. Iodine was discovered in 1811 by Courtois in kelp, and was shortly afterwards investigated by Clement, Gay-Lussac, and Davy. Bromine was discovered in 1826 by Balard in the mother liquor of sea water.

Bromine and iodine, like chlorine, occur in sea water in combination with metals. However, the amount of bromides, and especially of iodides, in sea water is so small that their presence can only be discovered by means of sensitive reactions.[56] In the extraction of salt from sea water the bromides remain in the mother liquor. Iodine and bromine also occur combined with silver, in admixture with silver chloride, as a rare ore which is mainly found in America. Certain mineral waters (those of Kreuznach and Staro-rossüsk) contain metallic bromides and iodides, always in admixture with an excess of sodium chloride. Those upper strata of the Stassfurt rock salt (Chapter X.) which are a source of potassium salts also contain metallic bromides,[57] which collect in the mother liquors left after the crystallisation of the potassium salts; and this now forms the chief source (together with certain American springs) of the bromine in common use. Bromine may be easily liberated from a mixture of bromides and chlorides, owing to the fact that chlorine displaces bromine from its compounds with sodium, magnesium, calcium, &c. A colourless solution of bromides and chlorides turns an orange colour after the passage of chlorine, owing to the disengagement of bromine.[58] Bromine may be extracted on a large scale by a similar method, but it is simpler to add a small quantity of manganese peroxide and sulphuric acid to the mother liquid direct. This sets free a portion of the chlorine, and this chlorine liberates the bromine.

Bromine is a dark brown liquid, giving brown fumes, and having a poisonous suffocating smell, whence its name (from the Greek βρῶμος, signifying evil smelling). The vapour density of bromine shows that its molecule is Br₂. In the cold bromine freezes into brown-grey scales like iodine. The melting point of pure bromine is -7°·05.[59] The density of liquid bromide at 0° is 3·187, and at 15° about 3·0. The boiling point of bromine is about 58°·7. Bromine, like chlorine, is soluble in water; 1 part of bromine at 5° requires 27 parts of water, and at 15° 29 parts of water. The aqueous solution of bromine is of an orange colour, and when cooled to -2° yields crystals containing 10 molecules of water to 1 molecule of bromine.[60] Alcohol dissolves a greater quantity of bromine, and ether a still greater amount. But after a certain time products of the action of the bromine on these organic substances are formed in the solutions. Aqueous solutions of the bromides also absorb a large amount of bromine.

With respect to iodine, it is almost exclusively extracted from the mother liquors after the crystallisation of natural sodium nitrate (Chili saltpetre) and from the ashes of the sea-weed cast upon the shores of France, Great Britain, and Spain, sometimes in considerable quantities, by the high tides. The majority of these sea-weeds are of the genera Fucus, Laminaria, &c. The fused ashes of these sea-weeds are called ‘kelp’ in Scotland and ‘varech’ in Normandy. A somewhat considerable quantity of iodine is contained in these sea-weeds. After being burnt (or subjected to dry distillation) an ash is left which chiefly contains salts of potassium, sodium, and calcium. The metals occur in the sea-weed as salts of organic acids. On being burnt these organic salts are decomposed, forming carbonates of potassium and sodium. Hence, sodium carbonate is found in the ash of sea plants. The ash is dissolved in hot water, and on evaporation sodium carbonate and other salts separate, but a portion of the substances remains in solution. These mother liquors left after the separation of the sodium carbonate contain chlorine, bromine, and iodine in combination with metals, the chlorine and iodine being in excess of the bromine. 13,000 kilos of kelp give about 1,000 kilos of sodium carbonate and 15 kilos of iodine.

The liberation of the iodine from the mother liquor is effected with comparative ease, because chlorine disengages iodine from potassium iodide and its other combinations with the metals. Not only chlorine, but also sulphuric acid, liberates iodine from sodium iodide. Sulphuric acid, in acting on an iodide, sets hydriodic acid free, but the latter easily decomposes, especially in the presence of substances capable of evolving oxygen, such as chromic acid, nitrous acid, and even ferric salts.[61] Owing to its sparing solubility in water, the iodine liberated separates as a precipitate. To obtain pure iodine it is sufficient to distil it, and neglect the first and last portions of the distillate, the middle portion only being collected. Iodine passes directly from a state of vapour into a crystalline form, and settles on the cool portions of the apparatus in tabular crystals, having a black grey colour and metallic lustre.[62]