Chapter 40

[31]For this reason it is necessary that in the preparation of bleaching powder the chlorine should be free from hydrochloric acid, and even the lime from calcium chloride. An excess of chlorine, in acting on a solution of bleaching powder, may also give chlorine monoxide, because calcium carbonate also gives chlorine monoxide under the action of chlorine. This reaction may be brought about by treating freshly precipitated calcium carbonate with a stream of chlorine in water: 2Cl2+ CaCO3= CO2+ CaCl2+ Cl2O. From this we may conclude that, although carbonic anhydride displaces hypochlorous anhydride, it may be itself displaced by an excess of the latter.[32]Dry red mercury oxide acts on chlorine, forming dry hypochlorous anhydride (chlorine monoxide) (Balard); when mixed with water, red mercury oxide acts feebly on chlorine, and when freshly precipitated it evolves oxygen and chlorine. An oxide of mercury which easily and abundantly evolves chlorine monoxide under the action of chlorine in the presence of water may be prepared as follows: the oxide of mercury, precipitated from a mercuric salt by an alkali, is heated to 300° and cooled (Pelouze). If a salt, MClO, be added to a solution of mercuric salt, HgX2, mercuric oxide is liberated, because the hypochlorite is decomposed.[32 bis]A solution of hypochlorous anhydride is also obtained by the action of chlorine on many salts; for example, in the action of chlorine on a solution of sodium sulphate the following reaction takes place: Na2SO4+ H2O + Cl2= NaCl + HClO + NaHSO4. Here the hypochlorous acid is formed, together with HCl, at the expense of chlorine and water, for Cl2+ H2O = HCl + HClO. If the crystallo-hydrate of chlorine be mixed with mercury oxide, the hydrochloric acid formed in the reaction gives mercury chloride, and hypochlorous acid remains in solution. A dilute solution of hypochlorous acid or chlorine monoxide may be concentrated by distillation, and if a substance which takes up water (without destroying the acid)—for instance, calcium nitrate—be added to the stronger solution, then the anhydride of hypochlorous acid—i.e.chlorine monoxide—is disengaged.[33]All explosive substances are of this kind—ozone, hydrogen peroxide, chloride of nitrogen, nitro-compounds, &c. Hence they cannot be formed directly from the elements or their simplest compounds, but, on the contrary, decompose into them. In a liquid state chlorine monoxide explodes even on contact with powdery substances, or when rapidly agitated—for instance, if a file be rasped over the vessel in which it is contained.[34]A solution of chlorine monoxide, or hypochlorous acid, does not explode, owing to the presence of the mass of water. In dissolving, chlorine monoxide evolves about 9,000 heat units, so that its store of heat becomes less.The capacity of hypochlorous acid (studied by Carius and others) for entering into combination with the unsaturated hydrocarbons is very often taken advantage of in organic chemistry. Thus its solution absorbs ethylene, forming the chlorhydrin C2H4ClOH.The oxidising action of hypochlorous acid and its salts is not only applied to bleaching but also to many reactions of oxidation. Thus it converts the lower oxides of manganese into the peroxide.[35]Chlorous acid, HClO2(according to the data given by Millon, Brandau, and others) in many respects resembles hypochlorous acid, HClO, whilst they both differ from chloric and perchloric acids in their degree of stability, which is expressed, for instance, in their bleaching properties; the two higher acids do not bleach, but both the lower ones do so (oxidise at the ordinary temperature). On the other hand, chlorous acid is analogous to nitrous acid, HNO2. The anhydride of chlorous acid, Cl2O3, is not known in a pure state, but it probably occurs in admixture with chlorine dioxide, ClO2, which is obtained by the action of nitric and sulphuric acids on a mixture of potassium chlorate with such reducing substances as nitric oxide, arsenious oxide, sugar, &c. All that is at present known is that pure chlorine dioxide ClO2(seeNotes39–43) is gradually converted into a mixture of hypochlorous and chlorous acids under the action of water (and alkalis); that is, it acts like nitric peroxide, NO2(giving HNO3and HNO2), or as a mixed anhydride, 2ClO2+ H2O = HClO3+ HClO2. The silver salt, AgClO2, is sparingly soluble in water. The investigations of Garzarolli-Thurnlackh and others seem to show that the anhydride Cl2O3does not exist in a free state.[36]Hydrochloric acid, which is an example of compounds of this kind, is a saturated substance which does not combine directly with oxygen, but in which, nevertheless, a considerable quantity of oxygen may be inserted between the elements forming it. The same may be observed in a number of other cases. Thus oxygen may be added or inserted between the elements, sometimes in considerable quantities, in the saturated hydrocarbons; for instance, in C3H8, three atoms of oxygen produce an alcohol, glycerin or glycerol, C3H5(OH)3. We shall meet with similar examples hereafter. This is generally explained by regarding oxygen as a bivalent element—that is, as capable of combining with two different elements, such as chlorine, hydrogen, &c. On the basis of this view, it may be inserted between each pair of combined elements; the oxygen will then be combined with one of the elements by one of its affinities and with the other element by its other affinity. This view does not, however, express the entire truth of the matter, even when applied to the compounds of chlorine. Hypochlorous acid, HOCl—that is, hydrochloric acid in which one atom of oxygen is inserted—is, as we have already seen, a substance of small stability; it might therefore be expected that on the addition of a fresh quantity of oxygen, a still less stable substance would be obtained, because, according to the above view, the chlorine and hydrogen, which form such a stable compound together, are then still further removed from each other. But it appears that chloric and perchloric acid, HClO3and HClO4, are much more stable substances. Furthermore, the addition of oxygen has also its limit, it can only be added to a certain extent. If the above representation were true and not merely hypothetical, there would be no limit to the combination of oxygen, and the more it entered into one continuous chain the more unstable would be the resultant compound. But not more than four atoms of oxygen can be added to hydrogen sulphide, nor to hydrochloric acid, nor to hydrogen phosphide. This peculiarity must lie in the properties of oxygen itself; four atoms of oxygen seem to have the power of forming a kind of radicle which retains two or several atoms of various other substances—for example, chlorine and hydrogen, hydrogen and sulphur, sodium and manganese, phosphorus and metals, &c., forming comparatively stable compounds, NaClO4, Na2SO4, NaMnO4, Na3PO4, &c.SeeChapter X. Note 1 and Chapter XV.[37]If chlorine be passed through acoldsolution of potash, a bleaching compound, potassium chloride and hypochlorite, KCl + KClO, is formed, but if it be passed through ahotsolution potassium chlorate is formed. As this is sparingly soluble in water, it chokes the gas-conducting tube, which should therefore be widened out at the end.Potassium chlorate is usually obtained on a large scale from calcium chlorate, which is prepared by passing chlorine (as long as it is absorbed) into water containing lime, the mixture being kept warm. A mixture of calcium chlorate and chloride is thus formed in the solution. Potassium chloride is then added to the warm solution, and on cooling a precipitate of potassium chlorate is formed as a substance which is sparingly soluble in cold water, especially in the presence of other salts. The double decomposition taking place is Ca(ClO3)2+ 2KCl = CaCl2+ 2KClO3. On a small scale in the laboratory potassium chlorate is best prepared from a strong solution of bleaching powder by passing chlorine through it and then adding potassium chloride. KClO3is always formed by the action of an electric current on a solution of KCl, especially at 80° (Häussermann and Naschold, 1894), so that this method is now used on a large scale.Potassium chlorate crystallises easily in large colourless tabular crystals. Its solubility in 100 parts of water at 0° = 3 parts, 20° = 8 parts, 40° = 14 parts, 60° = 25 parts, 80° = 40 parts. For comparison we will cite the following figures showing the solubility of potassium chloride and perchlorate in 100 parts of water: potassium chloride at O° = 28 parts, 20° = 35 parts, 40° = 40 parts, 100° = 57 parts; potassium perchlorate at 0° about 1 part, 20° about 1¾ part, 100° about 18 parts. When heated, potassium chlorate melts (the melting point has been given as from 335°–376°; according to the latest determination by Carnelley, 359°) and decomposes with the evolution of oxygen, potassium perchlorate being at first formed, as will afterwards be described (seeNote47). A mixture of potassium chlorate and nitric and hydrochloric acids effects oxidation and chlorination in solutions. It deflagrates when thrown upon incandescent carbon, and when mixed with sulphur (⅓ by weight) it ignites it on being struck, in which case an explosion takes place. The same occurs with many metallic sulphides and organic substances. Such mixtures are also ignited by a drop of sulphuric acid. All these effects are due to the large amount of oxygen contained in potassium chlorate, and to the ease with which it is evolved. A mixture of two parts of potassium chlorate, one part of sugar, and one part of yellow prussiate of potash acts like gunpowder, but burns too rapidly, and therefore bursts the guns, and it also has a very strong oxidising action on their metal. The sodium salt, NaClO3, is much more soluble than the potassium salt, and it is therefore more difficult to free it from sodium chloride, &c. The barium salt is also more soluble than the potassium salt; O° = 24 parts, 20° = 37 parts, 80° = 98 parts of salt per 100 of water.[38]Barium chlorate, Ba(ClO3)2,H2O, is prepared in the following way: impure chloric acid is first prepared and saturated with baryta, and the barium salt purified by crystallisation. The impure free chloric acid is obtained by converting the potassium in potassium chlorate into an insoluble salt. This is done by adding tartaric or hydrofluosilicic acid to a solution of potassium chlorate, because potassium tartrate and potassium silicofluoride are very sparingly soluble in water. Chloric acid is easily soluble in water.[39]To prepare ClO2100 grams of sulphuric acid are cooled in a mixture of ice and salt, and 15 grams of powdered potassium chlorate are gradually added to the acid, which is then carefully distilled at 20° to 40°, the vapour given off being condensed in a freezing mixture. Potassium perchlorate is then formed: 3KClO3+ 2H2SO4= 2KHSO4+ KClO4+ 2ClO2+ H2O. The reaction may result in an explosion. Calvert and Davies obtained chloric peroxide without the least danger by heating a mixture of oxalic acid and potassium chlorate in a test tube in a water-bath. In this case 2KClO3+ 3C2H2O4,2H2O = 2C2HKO4+ 2CO2+ 2ClO2+ 8H2O. The reaction is still further facilitated by the addition of a small quantity of sulphuric acid. If a solution of HCl acts upon KClO3at the ordinary temperature, a mixture of Cl2and ClO2is formed, but if the temperature be raised to 80° the greater part of the ClO2decomposes, and when passed through a hot solution of MnCl2it oxidises it. Gooch and Kreider proposed (1894) to employ this method for preparing small quantities of chlorine in the laboratory.[40]By analogy with nitric peroxide it might be expected that at low temperatures a doubling of the molecule into Cl2O4would take place, as the reactions of ClO2point to its being a mixed anhydride of HClO2and HClO3.[41]Owing to the formation of this chlorine dioxide, a mixture of potassium chlorate and sugar is ignited by a drop of sulphuric acid. This property was formerly made use of for making matches, and is now sometimes employed for setting fire to explosive charges by means of an arrangement in which the acid is caused to fall on the mixture at the moment required. An interesting experiment on the combustion of phosphorus under water may be conducted with chlorine dioxide. Pieces of phosphorus and of potassium chlorate are placed under water, and sulphuric acid is poured on to them (through a long funnel); the phosphorus then burns at the expense of the chlorine dioxide.[42]Potassium permanganate oxidises chlorine dioxide into chloric acid (Fürst).[43]The euchlorine obtained by Davy by gently heating potassium chlorate with hydrochloric acid is (Pebal) a mixture of chlorine dioxide and free chlorine. The liquid and gaseous chlorine oxide (Note35), which Millon considered to be Cl2O3, probably contains a mixture of ClO2(vapour density 35), Cl2O3(whose vapour density should be 59), and chlorine (vapour density 35·5), since its vapour density was determined to be about 40.[44]If a solution of chloric acid, HClO3, be first concentrated over sulphuric acid under the receiver of an air-pump and afterwards distilled, chlorine and oxygen are evolved and perchloric acid is formed: 4HClO3= 2HClO4+ Cl2+ 3O + H2O. Roscoe accordingly decomposed directly a solution of potassium chlorate by hydrofluosilicic acid, decanted it from the precipitate of potassium silicofluoride, K2SiF6, concentrated the solution of chloric acid, and then distilled it, perchloric acid being then obtained (seefollowing footnote). That chloric acid is capable of passing into perchloric acid is also seen from the fact that potassium permanganate is decolorised, although slowly, by the action of a solution of chloric acid. On decomposing a solution of potassium chlorate by the action of an electric current, potassium perchlorate is obtained at the positive electrode (where the oxygen is evolved). Perchloric acid is also formed by the action of an electric current on solutions of chlorine and chlorine monoxide. Perchloric acid was obtained by Count Stadion and afterwards by Serullas, and was studied by Roscoe and others.[45]Perchloric acid, which is obtained in a free state by the action of sulphuric acid on its salts, may be separated from a solution very easily by distillation, being volatile, although it is partially decomposed by distillation. The solution obtained after distillation may be concentrated by evaporation in open vessels. In the distillation the solution reaches a temperature of 200°, and then a very constant liquid hydrate of the composition HClO4,2H2O is obtained in the distillate. If this hydrate be mixed with sulphuric acid, it begins to decompose at 100°, but nevertheless a portion of the acid passes over into the receiver without decomposing, forming a crystalline hydrate HClO4,H2O which melts at 50°. On carefully heating this hydrate it breaks up into perchloric acid, which distills over below 100°, and into the liquid hydrate HClO4,2H2O. The acid HClO4may also be obtained by adding one-fourth part of strong sulphuric acid to potassium chlorate, carefully distilling and subjecting the crystals of the hydrate HClO4,H2O obtained in the distillate to a fresh distillation. Perchloric acid, HClO4, itself does not distil, and is decomposed on distillation until the more stable hydrate HClO4, H2O is formed; this decomposes into HClO4and HClO4,2H2O, which latter hydrate distils without decomposition. This forms an excellent example of the influence of water on stability, and of the property of chlorine of giving compounds of the type ClX7, of which all the above hydrates, ClO3(OH), ClO2(OH)3, and ClO(OH)5, are members. Probably further research will lead to the discovery of a hydrate Cl(OH)7.[46]According to Roscoe the specific gravity of perchloric acid = 1·782 and of the hydrate HClO4,H2O in a liquid state (50°) 1·811; hence a considerable contraction takes place in the combination of HClO4with H2O.[47]The decomposition of salts analogous to potassium chlorate has been more fully studied in recent years by Potilitzin and P. Frankland. Professor Potilitzin, by decomposing, for example, lithium chlorate LiClO3, found (from the quantity of lithium chloride and oxygen) that at first the decomposition of the fused salt (368°) takes place according to the equation, 3LiClO3= 2LiCl + LiClO4+ 5O, and that towards the end the remaining salt is decomposed thus: 5LiClO3= 4LiCl + LiClO4+ 10O. The phenomena observed by Potilitzin obliged him to admit that lithium perchlorate is capable of decomposing simultaneously with lithium chlorate, with the formation of the latter salt and oxygen; and this was confirmed by direct experiment, which showed that lithium chlorate is always formed in the decomposition of the perchlorate. Potilitzin drew particular attention to the fact that the decomposition of potassium chlorate and of salts analogous to it, although exothermal (Chapter III., Note12), not only does not proceed spontaneously, but requires time and a rise of temperature in order to attain completion, which again shows that chemical equilibria are not determined by the heat effects of reactions only.P. Frankland and J. Dingwall (1887) showed that at 448° (in the vapour of sulphur) a mixture of potassium chlorate and powdered glass is decomposed almost in accordance with the equation 2KClO3= KClO4+ KCl + O2, whilst the salt by itself evolves about half as much oxygen, in accordance with the equation, 8KClO3= 5KClO4+ 3KCl + 2O2. The decomposition of potassium perchlorate in admixture with manganese peroxide proceeds to completion, KClO4= KCl + 2O2. But in decomposing by itself the salt at first gives potassium chlorate, approximately according to the equation 7KClO4= 2KClO3+ 5KCl + 11O2. Thus there is now no doubt that when potassium chlorate is heated, the perchlorate is formed, and that this salt, in decomposing with evolution of oxygen, again gives the former salt.In the decomposition of barium hypochlorite, 50 per cent. of the whole amount passes into chlorate, in the decomposition of strontium hypochlorite (Potilitzin, 1890) 12·5 per cent., and of calcium hypochlorite about 2·5 per cent. Besides which Potilitzin showed that the decomposition of the hypochlorites and also of the chlorates is always accompanied by the formation of a certain quantity of the oxides and by the evolution of chlorine, the chlorine being displaced by the oxygen disengaged. Spring and Prost (1889) represent the evolution of oxygen from KClO3as due to the salt first splitting up into base and anhydride, thus (1) 2MClO3= M2O + Cl2O5; (2) Cl2O5= Cl2+ O3; and (3) M2O + Cl = 2MCl + O.I may further remark that the decomposition of potassium chlorate as a reaction evolving heat easily lends itself for this very reason to the contact action of manganese peroxide and other similar admixtures; for such very feeble influences as those of contact may become evident either in those cases (for instance, detonating gas, hydrogen peroxide, &c.), when the reaction is accompanied by the evolution of heat, or when (for instance, H2+ I2, &c.) little heat is absorbed or evolved. In these cases it is evident that the existing equilibrium is not very stable, and that a small alteration in the conditions at the surfaces of contact may suffice to upset it. In order to conceive themodus operandiof contact phenomena, it is enough to imagine, for instance, that at the surface of contact the movement of the atoms in the molecules changes from a circular to an elliptical path. Momentary and transitory compounds may he formed, but their formation cannot affect the explanation of the phenomena.[47 bis]See, for example the melting point of NaCl, NaBr, NaI in Chapter II. Note27. According to F. Freyer and V. Meyer (1892), the following are the boiling points of some of the corresponding compounds of chlorine and bromine:BCl317°BBr390°SiCl359°SiBr4153°PCl376°PBr3175°SbCl3223°SbBr3275°BiCl3447°BiBr3453°SnCl4606°SnBr4619°ZnCl2730°ZnBr2650°Thus for all the more volatile compounds the replacement of chlorine by bromine raises the boiling point, but in the ease of ZnX2it lowers it (Chapter XV. Note 19).[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 (seefirst Russian edition of thePrinciples 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 + F2= CaF2+ O—that is, tbe calcium distributed itself between the oxygen and fluorine. MnF4, which is capable of splitting up into MnF2and F2, 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, KHF2. 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 H2S, HCl, HBr, CS2, and CNH with a flash; it does not act upon O2, N2, CO, and CO2; Mg, Al, Ag, and Ni, when heated, burn in it, as also do S, Se, P (forms PF5); it reacts upon H2even in the dark, with the evolution of 366·00 units of heat. At a temperature of -95°, F2still retains its gaseous state. Soot and carbon in general (but not the diamond) when heated in gaseous fluorine formfluoride of carbon, CF4(Moissan, 1890); this compound is also formed at 300° by the double decomposition of CCl4and AgF; it is a gas which liquefies at 10° under a pressure of 5 atmospheres. With an alcoholic solution of KHO, CF4gives K2CO3, according to the equation CF4+ 6KHO = K2CO3+ 4KF + 3H2O. CF4is not soluble in water, but it is easily soluble in CCl4and alcohol.[49 bis]T. Nikolukin (1885) and subsequently Friedrich and Classen obtained PbCl4and a double ammonium salt of tetrachloride of lead (starting from the binoxide), PbCl42NH4Cl; Hutchinson and Pallard obtained a similar salt of acetic acid (1893) corresponding to PbX4by treating red lead with strong acetic acid; the composition of this salt is Pb(C2H3O2)4; it melts (and decomposes) at about 175°. Brauner (1894) obtained a salt corresponding to tetrafluoride of lead, PbF4, and the acid corresponding to it, H4PbF8. For example, by treating potassium plumbate (Chapter XVIII. Note 55) with strong HF, and also the above-mentioned tetra-acetate with a solution of KHF2, Brauner obtained crystalline HK3PbF8—i.e. the salt from which he obtained fluorine.[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, SiO2+ 2CaF2= 2CaO + SiF4; 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: CaF2+ H2O = 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(NO3)2+ 2KF = CaF2(precipitate) + 2KNO3(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.[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 HKF2, 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,2H2O. 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, H2F2being 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, SiO2, as we shall consider more fully in describing that compound, forming gaseous silicon fluoride, SiO2+ 4HF = SiF4+ 2H2O. 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. Inengraving on glassthe 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 H2F2is probably connected with the property of many fluorides of forming acid salts—for example, KHF2and H2SiF6. We saw above that HCl has the same property (forming, for instance, H2PtCl6, &c., p. 457), and hence this property of hydrofluoric acid does not stand isolated from the properties of the other halogens.[55]For instance, the experiment with Dutch metal foil (Note16) 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:—123K2+X2211191160Na2+X2195172138Ag2+X2594528Hg2+X2836848Hg+X2635134Ca+X2170141—Ba+X2195170—Zn+X2977649Pb+X2836440Al+X216112070We may remark that the latent heat of vaporisation of the molecular weight Br2is about 7·2, and of iodine 6·0 thousand heat units, whilst the latent heat of fusion of Br2is about 0·3, and of I2about 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 ZnI2would give 48 + 8 + 3, or about sixty thousand heat units, or 1½ times less than Zn + Cl2.[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 + KBrO3+ 6H2SO4= 6KHSO4+ 3H2O + 3Br2. 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 CaBr2.[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 tensionpin mm. attwas found to be—p=202530354045 mm.t=-16°·6-14°-12°-10°-8·5°-7°For liquid bromine—p=50100200400600760 mm.t=-5°·0+8°·223°·440°451°·958°·7These curves intersect at -7°·05. Besides which, in comparing the vapour tension of many liquids (for example, those given in Chapter II., Note27), Ramsay and Young observed that the ratio of the absolute temperatures (t+ 273) corresponding with equal tensionvariesfor every pair of substances in rectilinear proportion in dependence upont, and, therefore, for the above pressurep, Ramsay and Young determined the ratio oft+ 273 for water and bromine, and found that the straight lines expressing these ratios for liquid and solid bromine intersect also at 7°·05; thus, for example, for solid bromine—p=202530354045273 +t=256·4259261263264·6266273 +t′ =295·3299302·1304·8307·2309·3c=1·1521·1541·1571·1591·1611·163wheret′ indicates the temperature of water corresponding with a vapour tensionp, and wherecis the ratio of 273 +t′ to 273 +t. The magnitude ofcis evidently expressed with great accuracy by the straight linec= 1·1703 + 0·0011t. In exactly the same way we find the ratio for liquid bromine and water to bec1= 1·1585 + 0·00057t. The intersection of these straight lines in fact corresponds with -7°·06, which again confirms the melting point given above for bromine. In this manner it is possible with the existing store of data to accurately establish andverifythe melting point of substances. Ramsay and Young established the thermal constants of iodine by exactly the same method.[60]The observations made by Paterno and Nasini (by Raoult's method, Chapter I. Note49) on the temperature of the formation of ice ( -1°·115, with 1·391 gram of bromine in 100 grams of water) in an aqueous solution of bromine, showed that bromine is contained in solutions as the molecule Br2. Similar experiments conducted on iodine (Kloboukoff 1889 and Beckmann 1890) show that in solution the molecule is I2.B. Roozeboom investigated the hydrate of bromine as completely as the hydrate of chlorine (Notes9,10). The temperature of the complete decomposition of the hydrate is +6°·2; the density of Br2,10H2O = 1·49.

[32]Dry red mercury oxide acts on chlorine, forming dry hypochlorous anhydride (chlorine monoxide) (Balard); when mixed with water, red mercury oxide acts feebly on chlorine, and when freshly precipitated it evolves oxygen and chlorine. An oxide of mercury which easily and abundantly evolves chlorine monoxide under the action of chlorine in the presence of water may be prepared as follows: the oxide of mercury, precipitated from a mercuric salt by an alkali, is heated to 300° and cooled (Pelouze). If a salt, MClO, be added to a solution of mercuric salt, HgX2, mercuric oxide is liberated, because the hypochlorite is decomposed.

[32 bis]A solution of hypochlorous anhydride is also obtained by the action of chlorine on many salts; for example, in the action of chlorine on a solution of sodium sulphate the following reaction takes place: Na2SO4+ H2O + Cl2= NaCl + HClO + NaHSO4. Here the hypochlorous acid is formed, together with HCl, at the expense of chlorine and water, for Cl2+ H2O = HCl + HClO. If the crystallo-hydrate of chlorine be mixed with mercury oxide, the hydrochloric acid formed in the reaction gives mercury chloride, and hypochlorous acid remains in solution. A dilute solution of hypochlorous acid or chlorine monoxide may be concentrated by distillation, and if a substance which takes up water (without destroying the acid)—for instance, calcium nitrate—be added to the stronger solution, then the anhydride of hypochlorous acid—i.e.chlorine monoxide—is disengaged.

[33]All explosive substances are of this kind—ozone, hydrogen peroxide, chloride of nitrogen, nitro-compounds, &c. Hence they cannot be formed directly from the elements or their simplest compounds, but, on the contrary, decompose into them. In a liquid state chlorine monoxide explodes even on contact with powdery substances, or when rapidly agitated—for instance, if a file be rasped over the vessel in which it is contained.

[34]A solution of chlorine monoxide, or hypochlorous acid, does not explode, owing to the presence of the mass of water. In dissolving, chlorine monoxide evolves about 9,000 heat units, so that its store of heat becomes less.The capacity of hypochlorous acid (studied by Carius and others) for entering into combination with the unsaturated hydrocarbons is very often taken advantage of in organic chemistry. Thus its solution absorbs ethylene, forming the chlorhydrin C2H4ClOH.The oxidising action of hypochlorous acid and its salts is not only applied to bleaching but also to many reactions of oxidation. Thus it converts the lower oxides of manganese into the peroxide.

The capacity of hypochlorous acid (studied by Carius and others) for entering into combination with the unsaturated hydrocarbons is very often taken advantage of in organic chemistry. Thus its solution absorbs ethylene, forming the chlorhydrin C2H4ClOH.

The oxidising action of hypochlorous acid and its salts is not only applied to bleaching but also to many reactions of oxidation. Thus it converts the lower oxides of manganese into the peroxide.

[35]Chlorous acid, HClO2(according to the data given by Millon, Brandau, and others) in many respects resembles hypochlorous acid, HClO, whilst they both differ from chloric and perchloric acids in their degree of stability, which is expressed, for instance, in their bleaching properties; the two higher acids do not bleach, but both the lower ones do so (oxidise at the ordinary temperature). On the other hand, chlorous acid is analogous to nitrous acid, HNO2. The anhydride of chlorous acid, Cl2O3, is not known in a pure state, but it probably occurs in admixture with chlorine dioxide, ClO2, which is obtained by the action of nitric and sulphuric acids on a mixture of potassium chlorate with such reducing substances as nitric oxide, arsenious oxide, sugar, &c. All that is at present known is that pure chlorine dioxide ClO2(seeNotes39–43) is gradually converted into a mixture of hypochlorous and chlorous acids under the action of water (and alkalis); that is, it acts like nitric peroxide, NO2(giving HNO3and HNO2), or as a mixed anhydride, 2ClO2+ H2O = HClO3+ HClO2. The silver salt, AgClO2, is sparingly soluble in water. The investigations of Garzarolli-Thurnlackh and others seem to show that the anhydride Cl2O3does not exist in a free state.

[36]Hydrochloric acid, which is an example of compounds of this kind, is a saturated substance which does not combine directly with oxygen, but in which, nevertheless, a considerable quantity of oxygen may be inserted between the elements forming it. The same may be observed in a number of other cases. Thus oxygen may be added or inserted between the elements, sometimes in considerable quantities, in the saturated hydrocarbons; for instance, in C3H8, three atoms of oxygen produce an alcohol, glycerin or glycerol, C3H5(OH)3. We shall meet with similar examples hereafter. This is generally explained by regarding oxygen as a bivalent element—that is, as capable of combining with two different elements, such as chlorine, hydrogen, &c. On the basis of this view, it may be inserted between each pair of combined elements; the oxygen will then be combined with one of the elements by one of its affinities and with the other element by its other affinity. This view does not, however, express the entire truth of the matter, even when applied to the compounds of chlorine. Hypochlorous acid, HOCl—that is, hydrochloric acid in which one atom of oxygen is inserted—is, as we have already seen, a substance of small stability; it might therefore be expected that on the addition of a fresh quantity of oxygen, a still less stable substance would be obtained, because, according to the above view, the chlorine and hydrogen, which form such a stable compound together, are then still further removed from each other. But it appears that chloric and perchloric acid, HClO3and HClO4, are much more stable substances. Furthermore, the addition of oxygen has also its limit, it can only be added to a certain extent. If the above representation were true and not merely hypothetical, there would be no limit to the combination of oxygen, and the more it entered into one continuous chain the more unstable would be the resultant compound. But not more than four atoms of oxygen can be added to hydrogen sulphide, nor to hydrochloric acid, nor to hydrogen phosphide. This peculiarity must lie in the properties of oxygen itself; four atoms of oxygen seem to have the power of forming a kind of radicle which retains two or several atoms of various other substances—for example, chlorine and hydrogen, hydrogen and sulphur, sodium and manganese, phosphorus and metals, &c., forming comparatively stable compounds, NaClO4, Na2SO4, NaMnO4, Na3PO4, &c.SeeChapter X. Note 1 and Chapter XV.

[37]If chlorine be passed through acoldsolution of potash, a bleaching compound, potassium chloride and hypochlorite, KCl + KClO, is formed, but if it be passed through ahotsolution potassium chlorate is formed. As this is sparingly soluble in water, it chokes the gas-conducting tube, which should therefore be widened out at the end.Potassium chlorate is usually obtained on a large scale from calcium chlorate, which is prepared by passing chlorine (as long as it is absorbed) into water containing lime, the mixture being kept warm. A mixture of calcium chlorate and chloride is thus formed in the solution. Potassium chloride is then added to the warm solution, and on cooling a precipitate of potassium chlorate is formed as a substance which is sparingly soluble in cold water, especially in the presence of other salts. The double decomposition taking place is Ca(ClO3)2+ 2KCl = CaCl2+ 2KClO3. On a small scale in the laboratory potassium chlorate is best prepared from a strong solution of bleaching powder by passing chlorine through it and then adding potassium chloride. KClO3is always formed by the action of an electric current on a solution of KCl, especially at 80° (Häussermann and Naschold, 1894), so that this method is now used on a large scale.Potassium chlorate crystallises easily in large colourless tabular crystals. Its solubility in 100 parts of water at 0° = 3 parts, 20° = 8 parts, 40° = 14 parts, 60° = 25 parts, 80° = 40 parts. For comparison we will cite the following figures showing the solubility of potassium chloride and perchlorate in 100 parts of water: potassium chloride at O° = 28 parts, 20° = 35 parts, 40° = 40 parts, 100° = 57 parts; potassium perchlorate at 0° about 1 part, 20° about 1¾ part, 100° about 18 parts. When heated, potassium chlorate melts (the melting point has been given as from 335°–376°; according to the latest determination by Carnelley, 359°) and decomposes with the evolution of oxygen, potassium perchlorate being at first formed, as will afterwards be described (seeNote47). A mixture of potassium chlorate and nitric and hydrochloric acids effects oxidation and chlorination in solutions. It deflagrates when thrown upon incandescent carbon, and when mixed with sulphur (⅓ by weight) it ignites it on being struck, in which case an explosion takes place. The same occurs with many metallic sulphides and organic substances. Such mixtures are also ignited by a drop of sulphuric acid. All these effects are due to the large amount of oxygen contained in potassium chlorate, and to the ease with which it is evolved. A mixture of two parts of potassium chlorate, one part of sugar, and one part of yellow prussiate of potash acts like gunpowder, but burns too rapidly, and therefore bursts the guns, and it also has a very strong oxidising action on their metal. The sodium salt, NaClO3, is much more soluble than the potassium salt, and it is therefore more difficult to free it from sodium chloride, &c. The barium salt is also more soluble than the potassium salt; O° = 24 parts, 20° = 37 parts, 80° = 98 parts of salt per 100 of water.

Potassium chlorate is usually obtained on a large scale from calcium chlorate, which is prepared by passing chlorine (as long as it is absorbed) into water containing lime, the mixture being kept warm. A mixture of calcium chlorate and chloride is thus formed in the solution. Potassium chloride is then added to the warm solution, and on cooling a precipitate of potassium chlorate is formed as a substance which is sparingly soluble in cold water, especially in the presence of other salts. The double decomposition taking place is Ca(ClO3)2+ 2KCl = CaCl2+ 2KClO3. On a small scale in the laboratory potassium chlorate is best prepared from a strong solution of bleaching powder by passing chlorine through it and then adding potassium chloride. KClO3is always formed by the action of an electric current on a solution of KCl, especially at 80° (Häussermann and Naschold, 1894), so that this method is now used on a large scale.

Potassium chlorate crystallises easily in large colourless tabular crystals. Its solubility in 100 parts of water at 0° = 3 parts, 20° = 8 parts, 40° = 14 parts, 60° = 25 parts, 80° = 40 parts. For comparison we will cite the following figures showing the solubility of potassium chloride and perchlorate in 100 parts of water: potassium chloride at O° = 28 parts, 20° = 35 parts, 40° = 40 parts, 100° = 57 parts; potassium perchlorate at 0° about 1 part, 20° about 1¾ part, 100° about 18 parts. When heated, potassium chlorate melts (the melting point has been given as from 335°–376°; according to the latest determination by Carnelley, 359°) and decomposes with the evolution of oxygen, potassium perchlorate being at first formed, as will afterwards be described (seeNote47). A mixture of potassium chlorate and nitric and hydrochloric acids effects oxidation and chlorination in solutions. It deflagrates when thrown upon incandescent carbon, and when mixed with sulphur (⅓ by weight) it ignites it on being struck, in which case an explosion takes place. The same occurs with many metallic sulphides and organic substances. Such mixtures are also ignited by a drop of sulphuric acid. All these effects are due to the large amount of oxygen contained in potassium chlorate, and to the ease with which it is evolved. A mixture of two parts of potassium chlorate, one part of sugar, and one part of yellow prussiate of potash acts like gunpowder, but burns too rapidly, and therefore bursts the guns, and it also has a very strong oxidising action on their metal. The sodium salt, NaClO3, is much more soluble than the potassium salt, and it is therefore more difficult to free it from sodium chloride, &c. The barium salt is also more soluble than the potassium salt; O° = 24 parts, 20° = 37 parts, 80° = 98 parts of salt per 100 of water.

[38]Barium chlorate, Ba(ClO3)2,H2O, is prepared in the following way: impure chloric acid is first prepared and saturated with baryta, and the barium salt purified by crystallisation. The impure free chloric acid is obtained by converting the potassium in potassium chlorate into an insoluble salt. This is done by adding tartaric or hydrofluosilicic acid to a solution of potassium chlorate, because potassium tartrate and potassium silicofluoride are very sparingly soluble in water. Chloric acid is easily soluble in water.

[39]To prepare ClO2100 grams of sulphuric acid are cooled in a mixture of ice and salt, and 15 grams of powdered potassium chlorate are gradually added to the acid, which is then carefully distilled at 20° to 40°, the vapour given off being condensed in a freezing mixture. Potassium perchlorate is then formed: 3KClO3+ 2H2SO4= 2KHSO4+ KClO4+ 2ClO2+ H2O. The reaction may result in an explosion. Calvert and Davies obtained chloric peroxide without the least danger by heating a mixture of oxalic acid and potassium chlorate in a test tube in a water-bath. In this case 2KClO3+ 3C2H2O4,2H2O = 2C2HKO4+ 2CO2+ 2ClO2+ 8H2O. The reaction is still further facilitated by the addition of a small quantity of sulphuric acid. If a solution of HCl acts upon KClO3at the ordinary temperature, a mixture of Cl2and ClO2is formed, but if the temperature be raised to 80° the greater part of the ClO2decomposes, and when passed through a hot solution of MnCl2it oxidises it. Gooch and Kreider proposed (1894) to employ this method for preparing small quantities of chlorine in the laboratory.

[40]By analogy with nitric peroxide it might be expected that at low temperatures a doubling of the molecule into Cl2O4would take place, as the reactions of ClO2point to its being a mixed anhydride of HClO2and HClO3.

[41]Owing to the formation of this chlorine dioxide, a mixture of potassium chlorate and sugar is ignited by a drop of sulphuric acid. This property was formerly made use of for making matches, and is now sometimes employed for setting fire to explosive charges by means of an arrangement in which the acid is caused to fall on the mixture at the moment required. An interesting experiment on the combustion of phosphorus under water may be conducted with chlorine dioxide. Pieces of phosphorus and of potassium chlorate are placed under water, and sulphuric acid is poured on to them (through a long funnel); the phosphorus then burns at the expense of the chlorine dioxide.

[42]Potassium permanganate oxidises chlorine dioxide into chloric acid (Fürst).

[43]The euchlorine obtained by Davy by gently heating potassium chlorate with hydrochloric acid is (Pebal) a mixture of chlorine dioxide and free chlorine. The liquid and gaseous chlorine oxide (Note35), which Millon considered to be Cl2O3, probably contains a mixture of ClO2(vapour density 35), Cl2O3(whose vapour density should be 59), and chlorine (vapour density 35·5), since its vapour density was determined to be about 40.

[44]If a solution of chloric acid, HClO3, be first concentrated over sulphuric acid under the receiver of an air-pump and afterwards distilled, chlorine and oxygen are evolved and perchloric acid is formed: 4HClO3= 2HClO4+ Cl2+ 3O + H2O. Roscoe accordingly decomposed directly a solution of potassium chlorate by hydrofluosilicic acid, decanted it from the precipitate of potassium silicofluoride, K2SiF6, concentrated the solution of chloric acid, and then distilled it, perchloric acid being then obtained (seefollowing footnote). That chloric acid is capable of passing into perchloric acid is also seen from the fact that potassium permanganate is decolorised, although slowly, by the action of a solution of chloric acid. On decomposing a solution of potassium chlorate by the action of an electric current, potassium perchlorate is obtained at the positive electrode (where the oxygen is evolved). Perchloric acid is also formed by the action of an electric current on solutions of chlorine and chlorine monoxide. Perchloric acid was obtained by Count Stadion and afterwards by Serullas, and was studied by Roscoe and others.

[45]Perchloric acid, which is obtained in a free state by the action of sulphuric acid on its salts, may be separated from a solution very easily by distillation, being volatile, although it is partially decomposed by distillation. The solution obtained after distillation may be concentrated by evaporation in open vessels. In the distillation the solution reaches a temperature of 200°, and then a very constant liquid hydrate of the composition HClO4,2H2O is obtained in the distillate. If this hydrate be mixed with sulphuric acid, it begins to decompose at 100°, but nevertheless a portion of the acid passes over into the receiver without decomposing, forming a crystalline hydrate HClO4,H2O which melts at 50°. On carefully heating this hydrate it breaks up into perchloric acid, which distills over below 100°, and into the liquid hydrate HClO4,2H2O. The acid HClO4may also be obtained by adding one-fourth part of strong sulphuric acid to potassium chlorate, carefully distilling and subjecting the crystals of the hydrate HClO4,H2O obtained in the distillate to a fresh distillation. Perchloric acid, HClO4, itself does not distil, and is decomposed on distillation until the more stable hydrate HClO4, H2O is formed; this decomposes into HClO4and HClO4,2H2O, which latter hydrate distils without decomposition. This forms an excellent example of the influence of water on stability, and of the property of chlorine of giving compounds of the type ClX7, of which all the above hydrates, ClO3(OH), ClO2(OH)3, and ClO(OH)5, are members. Probably further research will lead to the discovery of a hydrate Cl(OH)7.

[46]According to Roscoe the specific gravity of perchloric acid = 1·782 and of the hydrate HClO4,H2O in a liquid state (50°) 1·811; hence a considerable contraction takes place in the combination of HClO4with H2O.

[47]The decomposition of salts analogous to potassium chlorate has been more fully studied in recent years by Potilitzin and P. Frankland. Professor Potilitzin, by decomposing, for example, lithium chlorate LiClO3, found (from the quantity of lithium chloride and oxygen) that at first the decomposition of the fused salt (368°) takes place according to the equation, 3LiClO3= 2LiCl + LiClO4+ 5O, and that towards the end the remaining salt is decomposed thus: 5LiClO3= 4LiCl + LiClO4+ 10O. The phenomena observed by Potilitzin obliged him to admit that lithium perchlorate is capable of decomposing simultaneously with lithium chlorate, with the formation of the latter salt and oxygen; and this was confirmed by direct experiment, which showed that lithium chlorate is always formed in the decomposition of the perchlorate. Potilitzin drew particular attention to the fact that the decomposition of potassium chlorate and of salts analogous to it, although exothermal (Chapter III., Note12), not only does not proceed spontaneously, but requires time and a rise of temperature in order to attain completion, which again shows that chemical equilibria are not determined by the heat effects of reactions only.P. Frankland and J. Dingwall (1887) showed that at 448° (in the vapour of sulphur) a mixture of potassium chlorate and powdered glass is decomposed almost in accordance with the equation 2KClO3= KClO4+ KCl + O2, whilst the salt by itself evolves about half as much oxygen, in accordance with the equation, 8KClO3= 5KClO4+ 3KCl + 2O2. The decomposition of potassium perchlorate in admixture with manganese peroxide proceeds to completion, KClO4= KCl + 2O2. But in decomposing by itself the salt at first gives potassium chlorate, approximately according to the equation 7KClO4= 2KClO3+ 5KCl + 11O2. Thus there is now no doubt that when potassium chlorate is heated, the perchlorate is formed, and that this salt, in decomposing with evolution of oxygen, again gives the former salt.In the decomposition of barium hypochlorite, 50 per cent. of the whole amount passes into chlorate, in the decomposition of strontium hypochlorite (Potilitzin, 1890) 12·5 per cent., and of calcium hypochlorite about 2·5 per cent. Besides which Potilitzin showed that the decomposition of the hypochlorites and also of the chlorates is always accompanied by the formation of a certain quantity of the oxides and by the evolution of chlorine, the chlorine being displaced by the oxygen disengaged. Spring and Prost (1889) represent the evolution of oxygen from KClO3as due to the salt first splitting up into base and anhydride, thus (1) 2MClO3= M2O + Cl2O5; (2) Cl2O5= Cl2+ O3; and (3) M2O + Cl = 2MCl + O.I may further remark that the decomposition of potassium chlorate as a reaction evolving heat easily lends itself for this very reason to the contact action of manganese peroxide and other similar admixtures; for such very feeble influences as those of contact may become evident either in those cases (for instance, detonating gas, hydrogen peroxide, &c.), when the reaction is accompanied by the evolution of heat, or when (for instance, H2+ I2, &c.) little heat is absorbed or evolved. In these cases it is evident that the existing equilibrium is not very stable, and that a small alteration in the conditions at the surfaces of contact may suffice to upset it. In order to conceive themodus operandiof contact phenomena, it is enough to imagine, for instance, that at the surface of contact the movement of the atoms in the molecules changes from a circular to an elliptical path. Momentary and transitory compounds may he formed, but their formation cannot affect the explanation of the phenomena.

P. Frankland and J. Dingwall (1887) showed that at 448° (in the vapour of sulphur) a mixture of potassium chlorate and powdered glass is decomposed almost in accordance with the equation 2KClO3= KClO4+ KCl + O2, whilst the salt by itself evolves about half as much oxygen, in accordance with the equation, 8KClO3= 5KClO4+ 3KCl + 2O2. The decomposition of potassium perchlorate in admixture with manganese peroxide proceeds to completion, KClO4= KCl + 2O2. But in decomposing by itself the salt at first gives potassium chlorate, approximately according to the equation 7KClO4= 2KClO3+ 5KCl + 11O2. Thus there is now no doubt that when potassium chlorate is heated, the perchlorate is formed, and that this salt, in decomposing with evolution of oxygen, again gives the former salt.

In the decomposition of barium hypochlorite, 50 per cent. of the whole amount passes into chlorate, in the decomposition of strontium hypochlorite (Potilitzin, 1890) 12·5 per cent., and of calcium hypochlorite about 2·5 per cent. Besides which Potilitzin showed that the decomposition of the hypochlorites and also of the chlorates is always accompanied by the formation of a certain quantity of the oxides and by the evolution of chlorine, the chlorine being displaced by the oxygen disengaged. Spring and Prost (1889) represent the evolution of oxygen from KClO3as due to the salt first splitting up into base and anhydride, thus (1) 2MClO3= M2O + Cl2O5; (2) Cl2O5= Cl2+ O3; and (3) M2O + Cl = 2MCl + O.

I may further remark that the decomposition of potassium chlorate as a reaction evolving heat easily lends itself for this very reason to the contact action of manganese peroxide and other similar admixtures; for such very feeble influences as those of contact may become evident either in those cases (for instance, detonating gas, hydrogen peroxide, &c.), when the reaction is accompanied by the evolution of heat, or when (for instance, H2+ I2, &c.) little heat is absorbed or evolved. In these cases it is evident that the existing equilibrium is not very stable, and that a small alteration in the conditions at the surfaces of contact may suffice to upset it. In order to conceive themodus operandiof contact phenomena, it is enough to imagine, for instance, that at the surface of contact the movement of the atoms in the molecules changes from a circular to an elliptical path. Momentary and transitory compounds may he formed, but their formation cannot affect the explanation of the phenomena.

[47 bis]See, for example the melting point of NaCl, NaBr, NaI in Chapter II. Note27. According to F. Freyer and V. Meyer (1892), the following are the boiling points of some of the corresponding compounds of chlorine and bromine:BCl317°BBr390°SiCl359°SiBr4153°PCl376°PBr3175°SbCl3223°SbBr3275°BiCl3447°BiBr3453°SnCl4606°SnBr4619°ZnCl2730°ZnBr2650°Thus for all the more volatile compounds the replacement of chlorine by bromine raises the boiling point, but in the ease of ZnX2it lowers it (Chapter XV. Note 19).

Thus for all the more volatile compounds the replacement of chlorine by bromine raises the boiling point, but in the ease of ZnX2it lowers it (Chapter XV. Note 19).

[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 (seefirst Russian edition of thePrinciples 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 + F2= CaF2+ O—that is, tbe calcium distributed itself between the oxygen and fluorine. MnF4, which is capable of splitting up into MnF2and F2, 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, KHF2. 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 H2S, HCl, HBr, CS2, and CNH with a flash; it does not act upon O2, N2, CO, and CO2; Mg, Al, Ag, and Ni, when heated, burn in it, as also do S, Se, P (forms PF5); it reacts upon H2even in the dark, with the evolution of 366·00 units of heat. At a temperature of -95°, F2still retains its gaseous state. Soot and carbon in general (but not the diamond) when heated in gaseous fluorine formfluoride of carbon, CF4(Moissan, 1890); this compound is also formed at 300° by the double decomposition of CCl4and AgF; it is a gas which liquefies at 10° under a pressure of 5 atmospheres. With an alcoholic solution of KHO, CF4gives K2CO3, according to the equation CF4+ 6KHO = K2CO3+ 4KF + 3H2O. CF4is not soluble in water, but it is easily soluble in CCl4and alcohol.

[49 bis]T. Nikolukin (1885) and subsequently Friedrich and Classen obtained PbCl4and a double ammonium salt of tetrachloride of lead (starting from the binoxide), PbCl42NH4Cl; Hutchinson and Pallard obtained a similar salt of acetic acid (1893) corresponding to PbX4by treating red lead with strong acetic acid; the composition of this salt is Pb(C2H3O2)4; it melts (and decomposes) at about 175°. Brauner (1894) obtained a salt corresponding to tetrafluoride of lead, PbF4, and the acid corresponding to it, H4PbF8. For example, by treating potassium plumbate (Chapter XVIII. Note 55) with strong HF, and also the above-mentioned tetra-acetate with a solution of KHF2, Brauner obtained crystalline HK3PbF8—i.e. the salt from which he obtained fluorine.

[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, SiO2+ 2CaF2= 2CaO + SiF4; 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: CaF2+ H2O = 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(NO3)2+ 2KF = CaF2(precipitate) + 2KNO3(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.

[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 HKF2, 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,2H2O. 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, H2F2being 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, SiO2, as we shall consider more fully in describing that compound, forming gaseous silicon fluoride, SiO2+ 4HF = SiF4+ 2H2O. 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. Inengraving on glassthe 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 H2F2is probably connected with the property of many fluorides of forming acid salts—for example, KHF2and H2SiF6. We saw above that HCl has the same property (forming, for instance, H2PtCl6, &c., p. 457), and hence this property of hydrofluoric acid does not stand isolated from the properties of the other halogens.

[55]For instance, the experiment with Dutch metal foil (Note16) 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:—123K2+X2211191160Na2+X2195172138Ag2+X2594528Hg2+X2836848Hg+X2635134Ca+X2170141—Ba+X2195170—Zn+X2977649Pb+X2836440Al+X216112070We may remark that the latent heat of vaporisation of the molecular weight Br2is about 7·2, and of iodine 6·0 thousand heat units, whilst the latent heat of fusion of Br2is about 0·3, and of I2about 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 ZnI2would give 48 + 8 + 3, or about sixty thousand heat units, or 1½ times less than Zn + Cl2.

We may remark that the latent heat of vaporisation of the molecular weight Br2is about 7·2, and of iodine 6·0 thousand heat units, whilst the latent heat of fusion of Br2is about 0·3, and of I2about 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 ZnI2would give 48 + 8 + 3, or about sixty thousand heat units, or 1½ times less than Zn + Cl2.

[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 + KBrO3+ 6H2SO4= 6KHSO4+ 3H2O + 3Br2. 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 CaBr2.

[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 tensionpin mm. attwas found to be—p=202530354045 mm.t=-16°·6-14°-12°-10°-8·5°-7°For liquid bromine—p=50100200400600760 mm.t=-5°·0+8°·223°·440°451°·958°·7These curves intersect at -7°·05. Besides which, in comparing the vapour tension of many liquids (for example, those given in Chapter II., Note27), Ramsay and Young observed that the ratio of the absolute temperatures (t+ 273) corresponding with equal tensionvariesfor every pair of substances in rectilinear proportion in dependence upont, and, therefore, for the above pressurep, Ramsay and Young determined the ratio oft+ 273 for water and bromine, and found that the straight lines expressing these ratios for liquid and solid bromine intersect also at 7°·05; thus, for example, for solid bromine—p=202530354045273 +t=256·4259261263264·6266273 +t′ =295·3299302·1304·8307·2309·3c=1·1521·1541·1571·1591·1611·163wheret′ indicates the temperature of water corresponding with a vapour tensionp, and wherecis the ratio of 273 +t′ to 273 +t. The magnitude ofcis evidently expressed with great accuracy by the straight linec= 1·1703 + 0·0011t. In exactly the same way we find the ratio for liquid bromine and water to bec1= 1·1585 + 0·00057t. The intersection of these straight lines in fact corresponds with -7°·06, which again confirms the melting point given above for bromine. In this manner it is possible with the existing store of data to accurately establish andverifythe melting point of substances. Ramsay and Young established the thermal constants of iodine by exactly the same method.

For liquid bromine—

These curves intersect at -7°·05. Besides which, in comparing the vapour tension of many liquids (for example, those given in Chapter II., Note27), Ramsay and Young observed that the ratio of the absolute temperatures (t+ 273) corresponding with equal tensionvariesfor every pair of substances in rectilinear proportion in dependence upont, and, therefore, for the above pressurep, Ramsay and Young determined the ratio oft+ 273 for water and bromine, and found that the straight lines expressing these ratios for liquid and solid bromine intersect also at 7°·05; thus, for example, for solid bromine—

wheret′ indicates the temperature of water corresponding with a vapour tensionp, and wherecis the ratio of 273 +t′ to 273 +t. The magnitude ofcis evidently expressed with great accuracy by the straight linec= 1·1703 + 0·0011t. In exactly the same way we find the ratio for liquid bromine and water to bec1= 1·1585 + 0·00057t. The intersection of these straight lines in fact corresponds with -7°·06, which again confirms the melting point given above for bromine. In this manner it is possible with the existing store of data to accurately establish andverifythe melting point of substances. Ramsay and Young established the thermal constants of iodine by exactly the same method.

[60]The observations made by Paterno and Nasini (by Raoult's method, Chapter I. Note49) on the temperature of the formation of ice ( -1°·115, with 1·391 gram of bromine in 100 grams of water) in an aqueous solution of bromine, showed that bromine is contained in solutions as the molecule Br2. Similar experiments conducted on iodine (Kloboukoff 1889 and Beckmann 1890) show that in solution the molecule is I2.B. Roozeboom investigated the hydrate of bromine as completely as the hydrate of chlorine (Notes9,10). The temperature of the complete decomposition of the hydrate is +6°·2; the density of Br2,10H2O = 1·49.

B. Roozeboom investigated the hydrate of bromine as completely as the hydrate of chlorine (Notes9,10). The temperature of the complete decomposition of the hydrate is +6°·2; the density of Br2,10H2O = 1·49.

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