Chapter II. Note 27. According to F. Freyer and V. Meyer (1892),

the following are the boiling points of some of the corresponding compounds of chlorine and bromine:

BCl_{5} 17° BBr_{3} 90° SiCl_{3} 59° SiBr_{4} 153° PCl_{3} 76° PBr_{3} 175° SbCl_{3} 223° SbBr_{3} 275° BiCl_{3} 447° BiBr_{3} 453° SnCl_{4} 606° SnBr_{4} 619° ZnCl_{2} 730° ZnBr_{2} 650°

Thus for all the more volatile compounds the replacement of chlorine by bromine raises the boiling point, but in the ease of ZnX_{2} it lowers it (Chapter XV. Note 19).

Their oxygen compounds exhibit a similar analogy. Only fluorine does not give any oxygen compounds. The iodine and bromine compounds corresponding with HClO_{3} and HClO_{4} are HBrO_{3} and HBrO_{4}, HIO_{3} and HIO_{4}. 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 NH_{3}) and quinqui-valent to oxygen (it gives N_{2}O_{5}), and C being quadrivalent to both H and O as it forms CH_{4} and CO_{2}. 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, F_{2} + H_{2}O = 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, CaF_{2}, 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, 2CeF_{4} = 2CeF_{3} + F_{2}; 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, SiF_{4}), boron (forming BF_{3}), sulphur, &c. Its density (H = 1) is 18, so that its molecule is F_{2}. 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]

[48] Even before free fluorine was obtained (1886) it was evident from experience gained in the efforts made to obtain it, and from analogy, that it would decompose water (_see_ first Russian edition of the _Principles of Chemistry_).

[48 bis] It is most likely that in this experiment of Fremy's, which corresponds with the action of oxygen on calcium chloride, fluorine was set free, but that a converse reaction also proceeded, CaO + F_{2} = CaF_{2} + O--that is, the calcium distributed itself between the oxygen and fluorine. MnF_{4}, which is capable of splitting up into MnF_{2} and F_{2}, is without doubt formed by the action of a strong solution of hydrofluoric acid on manganese peroxide, but under the action of water the fluorine gives hydrofluoric acid, and probably this is aided by the affinity of the manganese fluoride and hydrofluoric acid. In all the attempts made (by Davy, Knox, Louget, Fremy, Gore, and others) to decompose fluorides (those of lead, silver, calcium, and others) by chlorine, there were doubtless also cases of distribution, a portion of the metal combined with chlorine and a portion of the fluorine was evolved; but it is improbable that any decisive results were obtained. Fremy probably obtained fluorine, but not in a pure state.

[49] According to Moissan, fluorine is disengaged by the action of an electric current on fused hydrogen potassium fluoride, KHF_{2}. The present state of chemical knowledge is such that the knowledge of the properties of an element is much more general than the knowledge of the free element itself. It is useful and satisfactory to learn that even fluorine in the free state has not succeeded in eluding experiment and research, that the efforts to isolate it have been crowned with success, but the sum total of chemical data concerning fluorine as an element gains but little by this achievement. The gain will, however, be augmented if it be now possible to subject fluorine to a comparative study in relation to oxygen and chlorine. There is particular interest in the phenomena of the distribution of fluorine and oxygen, or fluorine and chlorine, competing under different conditions and relations. We may add that Moissan (1892) found that free fluorine decomposes H_{2}S, HCl, HBr, CS_{2}, and CNH with a flash; it does not act upon O_{2}, N_{2}, CO, and CO_{2}; Mg, Al, Ag, and Ni, when heated, burn in it, as also do S, Se, P (forms PF_{5}); it reacts upon H_{2} even in the dark, with the evolution of 366·00 units of heat. At a temperature of -95°, F_{2} still retains its gaseous state. Soot and carbon in general (but not the diamond) when heated in gaseous fluorine form _fluoride of carbon_, CF_{4} (Moissan, 1890); this compound is also formed at 300° by the double decomposition of CCl_{4} and AgF; it is a gas which liquefies at 10° under a pressure of 5 atmospheres. With an alcoholic solution of KHO, CF_{4} gives K_{2}CO_{3}, according to the equation CF_{4} + 6KHO = K_{2}CO_{3} + 4KF + 3H_{2}O. CF_{4} is not soluble in water, but it is easily soluble in CCl_{4} and alcohol.

In 1894 Brauner obtained fluorine directly by igniting the easily formed[49 bis] double lead salt HF,3KF,PbF_{4}, which first, at 230°, decomposes with the evolution of HF, and then splits up forming 3KF,PbF_{2} and fluorine F_{2}, which is recognised by the fact that it liberates iodine from KI and easily combines with silicon, forming SiF_{4}. This method gives chemically pure fluorine, and is based upon the breaking up of the higher compound--tetrafluoride of lead, PbF_{4}, corresponding to PbO_{2}, into free fluorine, F_{2}, and the lower more stable form--bifluoride of lead, PbF_{2}, which corresponds to PbO; that is, this method resembles the ordinary method of obtaining chlorine by means of MnO_{2}, as MnCl_{4} here breaks up into MnCl_{2} and chlorine, just as PbF_{4} splits up into PbF_{2} and fluorine.

[49 bis] T. Nikolukin (1885) and subsequently Friedrich and Classen obtained PbCl_{4} and a double ammonium salt of tetrachloride of lead (starting from the binoxide), PbCl_{4}2NH_{4}Cl; Hutchinson and Pallard obtained a similar salt of acetic acid (1893) corresponding to PbX_{4} by treating red lead with strong acetic acid; the composition of this salt is Pb(C_{2}H_{3}O_{2})_{4}; it melts (and decomposes) at about 175°. Brauner (1894) obtained a salt corresponding to tetrafluoride of lead, PbF_{4}, and the acid corresponding to it, H_{4}PbF_{8}. For example, by treating potassium plumbate (Chapter XVIII. Note 55) with strong HF, and also the above-mentioned tetra-acetate with a solution of KHF_{2}, Brauner obtained crystalline HK_{3}PbF_{8}--i.e. the salt from which he obtained fluorine.

Among the compounds of fluorine, calcium fluoride, CaF_{2}, is somewhat widely distributed in nature as fluor spar,[50] whilst _cryolite_, or aluminium sodium fluoride, Na_{3}AlF_{6}, 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.

[50] It is called spar because it very frequently occurs as crystals of a clearly laminar structure, and is therefore easily split up into pieces bounded by planes. It is called fluor spar because when used as a flux it renders ores fusible, owing to its reacting with silica, SiO_{2} + 2CaF_{2} = 2CaO + SiF_{4}; the silicon fluoride escapes as a gas and the lime combines with a further quantity of silica, and gives a vitreous slag. Fluor spar occurs in mineral veins and rocks, sometimes in considerable quantities. It always crystallises in the cubic system, sometimes in very large semi-transparent cubic crystals, which are colourless or of different colours. It is insoluble in water. It melts under the action of heat, and crystallises on cooling. The specific gravity is 3·1. When steam is passed over incandescent fluor spar, lime and hydrofluoric acid are formed: CaF_{2} + H_{2}O = CaO + 2HF. A double decomposition is also easily produced by fusing fluor spar with sodium or potassium hydroxides, or potash, or even with their carbonates; the fluorine then passes over to the potassium or sodium, and the oxygen to the calcium. In solutions--for example, Ca(NO_{3})_{2} + 2KF = CaF_{2} (precipitate) + 2KNO_{3} (in solution)--the formation of calcium fluoride takes place, owing to its very sparing solubility. 26,000 parts of water dissolve one part of fluor spar.

Hydrifluoric 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, CaF_{2} + H_{2}SO_{4} = CaSO_{4} + 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,2H_{2}O.[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 HKF_{2} (and by dissolving KF in liquid HF, KHF_{2}2HF 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 H_{2}F_{2}, 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]

[51] According to Gore. Fremy obtained anhydrous hydrofluoric acid by decomposing lead fluoride at a red heat, by hydrogen, or by beating the double salt HKF_{2}, which easily crystallises (in cubes) from a solution of hydrofluoric acid, half of which has been saturated with potassium hydroxide. Its vapour density corresponds to the formula HF.

[52] This composition corresponds to the crystallo-hydrate HCl,2H_{2}O. All the properties of hydrofluoric acid recall those of hydrochloric acid, and therefore the comparative ease with which hydrofluoric acid is liquefied (it boils at +19°, hydrochloric acid at -35°) must be explained by a polymerisation taking place at low temperatures, as will be afterwards explained, H_{2}F_{2} being formed, and therefore in a liquid state it differs from hydrochloric acid, in which a phenomenon of a similar kind has not yet been observed.

[53] The corrosive action of hydrofluoric acid on glass and similar siliceous compounds is based upon the fact that it acts on silica, SiO_{2}, as we shall consider more fully in describing that compound, forming gaseous silicon fluoride, SiO_{2} + 4HF = SiF_{4} + 2H_{2}O. Silica, on the other hand, forms the binding (acid) element of glass and of the mass of mineral substances forming the salts of silica. When it is removed the cohesion is destroyed. This is made use of in the arts, and in the laboratory, for etching designs and scales, &c., on glass. In _engraving on glass_ the surface is covered with a varnish composed of four parts of wax and one part of turpentine. This varnish is not acted on by hydrofluoric acid, and it is soft enough to allow of designs being drawn upon it whose lines lay bare the glass. The drawing is made with a steel point, and the glass is afterwards laid in a lead trough in which a mixture of fluor spar and sulphuric acid is placed. The sulphuric acid must be used in considerable excess, as otherwise transparent lines are obtained (owing to the formation of hydrofluosilicic acid). After being exposed for some time, the varnish is removed (melted) and the design drawn by the steel point is found reproduced in dull lines. The drawing may be also made by the direct application of a mixture of a silicofluoride and sulphuric acid, which forms hydrofluoric acid.

[54] Mallet (1881) determined the density at 30° and 100°, previous to which Gore (1869) had determined the vapour density at 100°, whilst Thorpe and Hambly (1888) made fourteen determinations between 26° and 88°, and showed that within this limit of temperature the density gradually diminishes, just like the vapour of acetic acid, nitrogen dioxide, and others. The tendency of HF to polymerise into H_{2}F_{2} is probably connected with the property of many fluorides of forming acid salts--for example, KHF_{2} and H_{2}SiF_{6}. We saw above that HCl has the same property (forming, for instance, H_{2}PtCl_{6}, &c., p. 457), and hence this property of hydrofluoric acid does not stand isolated from the properties of the other halogens.

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.

[55] For instance, the experiment with Dutch metal foil (Note 16) may be made with bromine just as well as with chlorine. A very instructive experiment on the direct combination of the halogens with metals maybe made by throwing a small piece (a shaving) of aluminium into a vessel containing liquid bromine; the aluminium, being lighter, floats on the bromine, and after a certain time reaction sets in accompanied by the evolution of heat, light, and fumes of bromine. The incandescent piece of metal moves rapidly over the surface of the bromine in which the resultant aluminium bromide dissolves. For the sake of comparison we will proceed to cite several thermochemical data (Thomsen) for analogous actions of (1) chlorine, (2) bromine, and (3) iodine, with respect to metals; the halogen being expressed by the symbol X, and the plus sign connecting the reacting substances. All the figures are given in thousands of calories, and refer to molecular quantities in grams and to the ordinary temperature:--

1 2 3 K_{2} + X_{2} 211 191 160 Na_{2} + X_{2} 195 172 138 Ag_{2} + X_{2} 59 45 28 Hg_{2} + X_{2} 83 68 48 Hg + X_{2} 63 51 34 Ca + X_{2} 170 141 -- Ba + X_{2} 195 170 -- Zn + X_{2} 97 76 49 Pb + X_{2} 83 64 40 Al + X_{2} 161 120 70

We may remark that the latent heat of vaporisation of the molecular weight Br_{2} is about 7·2, and of iodine 6·0 thousand heat units, whilst the latent heat of fusion of Br_{2} is about 0·3, and of I_{2} about 3·0 thousand heat units. From this it is evident that the difference between the amounts of heat evolved does not depend on the difference in physical state. For instance, the vapour of iodine in combining with Zn to form ZnI_{2} would give 48 + 8 + 3, or about sixty thousand heat units, or 1-1/2 times less than Zn + Cl_{2}.

_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.

[56] One litre of sea-water contains about 20 grams of chlorine, and about 0·07 gram of bromine. The Dead Sea contains about ten times as much bromine.

[57] But there is no iodine in Stassfurt carnallite.

[58] The chlorine must not, however, be in large excess, as otherwise the bromine would contain chlorine. Commercial bromine not unfrequently contains chlorine, as bromine chloride; this is more soluble in water than bromine, from which it may thus be freed. To obtain pure bromine the commercial bromine is washed with water, dried by sulphuric acid, and distilled, the portion coming over at 58° being collected; the greater part is then converted into potassium bromide and dissolved, and the remainder is added to the solution in order to separate iodine, which is removed by shaking with carbon bisulphide. By heating the potassium bromide thus obtained with manganese peroxide and sulphuric acid, bromine is obtained quite free from iodine, which, however, is not present in certain kinds of commercial bromine (the Stassfurt, for instance). By treatment with potash, the bromine is then converted into a mixture of potassium bromide and bromate, and the mixture (which is in the proportion given in the equation) is distilled with sulphuric acid, bromine being then evolved: 5KBr + KBrO_{3} + 6H_{2}SO_{4} = 6KHSO_{4} + 3H_{2}O + 3Br_{2}. After dissolving the bromine in a strong solution of calcium bromide and precipitating with an excess of water, it loses all the chlorine it contained, because chlorine forms calcium chloride with CaBr_{2}.

Bromine is a _dark brown liquid_, giving brown fumes, and having a poisonous suffocating smell, whence its name (from the Greek [Greek: brômos], signifying evil smelling). The vapour density of bromine shows that its molecule is Br_{2}. 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.

[59] There has long existed a difference of opinion as to the melting point of pure bromine. By some investigators (Regnault, Pierre) it was given as between -7° and -8°, and by others (Balard, Liebig, Quincke, Baumhauer) as between -20° and -25°. There is now no doubt, thanks more especially to the researches of Ramsay and Young (1885), that pure bromine melts at about -7°. This figure is not only established by direct experiment (Van der Plaats confirmed it), but also by means of the determination of the vapour tensions. For solid bromine the vapour tension _p_ in mm. at _t_ was found to be--

_p_ = 20 25 30 35 40 45 mm. _t_ = -16°·6 -14° -12° -10° -8·5° -7°

For liquid bromine--

_p_ = 50 100 200 400 600 760 mm. _t_ = -5°·0 8°·2 23°·4 40°·4 51°·9 58°·7

These curves intersect at -7°·05. Besides which, in comparing the vapour tension of many liquids (for example, those given in