[9]Moissan (1893) studied the compounds of Mo and W formed with carbon in the electrical furnace (they are extremely hard) from a mixture of the anhydrides and carbon. Poleck and Grützner obtained definite compounds FeW2and FeW2C3for tungsten. Metallic W and Mo displace Ag from its solutions but not Pb. There is reason for believing that the sp. gr. of pure molybdenum is higher than that (8·5) generally ascribed to it.[9 bis]We may conclude our description of tungsten and molybdenum by stating that their sulphur compounds have an acid character, like carbon bisulphide or stannic sulphide. If sulphuretted hydrogen be passed through a solution of a molybdate it does not give a precipitate unless sulphuric acid be present, when a dark brown precipitate ofmolybdenum trisulphide, MoS3, is formed. When this sulphide is ignited without access of air it gives the bisulphide MoS2; the latter is not able to combine with potassium sulphide like the trisulphide MoS3, which forms a salt, K2MoS4, corresponding with K2MoO4. This is soluble in water, and separates out from its solution in red crystals, which have a metallic lustre and reflect a green light. It is easily obtained by heating the native bisulphide, MoS2, with potash, sulphur, and a small amount of charcoal, which serves for deoxidising the oxygen compounds. Tungsten gives similar compounds, R2WS4, where R = NH4, K, Na. They are decomposed by acids, with the separation of tungsten trisulphide, WS3, and molybdenum trisulphide, MoS3. Rideal (1892) obtained W2N3by heating WO3in NH3. This compound exhibited the general properties of metallic nitrides.[9 tri]When peroxide of hydrogen acts upon a solution of potassium molybdate well-formed yellow crystals belonging to the triclinic system separate out in the cold. When these crystals are heated in vacuo they first lose water and then decompose, leaving a residue composed of the salt originally taken. They are soluble in water but insoluble in alcohol. Their composition is represented by the formula K2Mo2O82H2O. An ammonium salt is obtained by evaporating peroxide of hydrogen with ammonium molybdate. The following salts have also been obtained by the action of peroxide of hydrogen upon the corresponding molybdates: Na2Mo2O66H2O—in yellow prismatic crystals; MgMo2O810H2O—stellar needles; BaMoO82H2O—in microscopic yellow octahedra. A corresponding sodium pertungstate has been obtained by Péchard by boiling sodium tungstate with a solution of peroxide of hydrogen for several minutes. The solution rapidly turns yellow, and no longer gives a precipitate of tungstic anhydride when treated with nitric acid. When evaporated in vacuo the solution leaves a thick syrupy liquid from which ray-like crystals separate out; these crystals are more soluble in water than the salt originally taken. When heated they also lose water and oxygen. Their composition answers to the formula M2W2O82H2O, where M = Na, NH4, &c. The permolybdates and pertungstates have similar properties. When treated with oxygen acids they give peroxide of hydrogen, and disengage chlorine and iodine from hydrochloric acid and potassium iodide.Piccini (1891) showed that peroxide of hydrogen not only combines with the oxygen compounds of Mo and W, but also with their fluo-compounds, among which ammonium fluo-molybdate MoO2F22NH4and others have long been known. (A few new salts of similar composition have been obtained by F. Moureu in 1893.) The action of peroxide of hydrogen upon these compounds gives salts containing a larger amount of oxygen; for instance, a solution of MoO2F22KFH2O with peroxide of hydrogen gives a yellow solution which after cooling separates out yellow crystalline flakes of MoO3F22KFH2O, resembling the salt originally taken in their external appearance. By employing a similar method Piccini also obtained: MoO3F22RbFH2O—yellow monoclinic crystals; MoO3F2,2CsFH2O,—yellow flakes, and the corresponding tungstic compounds. All these salts react like peroxide of hydrogen.In speaking of these compounds I for my part think it may be well to call attention to the fact that, in the first place, the composition of Piccini's oxy-fluo compounds does not correspond to that of permolybdic and pertungstic acid. If the latter be expressed by formulæ with one equivalent of an element, they will be HMoO4and HWO4, and the oxy-fluo form corresponding to them should have the composition MoO3F and WO3F while it contains MO3F2and WO3F2,i.e.answers as it were to a higher degree of oxidation, MoH2O3and W3HO3. But if permolybdic acid be regarded as 2MoO3+ H2O2,i.e.as containing the elements of peroxide of hydrogen, then Piccini's compound will also be found to contain the original salts + H2O; for example, from MoO2F22KFH2O there is obtained a compound MoO2F22KFH2O2,i.e.instead of H2O they contain H2O2. In the second place the capacity of the salts of molybdenum and tungsten to retain a further amount of oxygen or H2O2probably bears some relation to their property of giving complex acids and of polymerising which has been considered in Note8 bis. There is, however, a great chemical interest in the accumulation of data respecting these high peroxide compounds corresponding to molybdic and tungstic acids. With regard to the peroxide form of uranium,seeChapter XX., Note66.[10]Uranium trioxide, or uranic oxide, shows its feeble basic and acid properties in a great number of its reactions. (1) Solutions of uranic salts give yellow precipitates with alkalis, but these precipitates do not contain the hydrate of the oxide, but compounds of it with bases; for example, 2UO2(NO3)2+ 6KHO = 4KNO3+ 3H2O + K2U2O7. There are otherurano-alkali compoundsof the same constitution; for example, (NH4)2U2O7(known commercially as uranic oxide), MgU2O7, BaU2O7. They are the analogues of the dichromates. Sodium uranate is the most generally used under the name of uranium yellow, Na2U2O7. It is used for imparting the characteristic yellow-green tint to glass and porcelain. Neither heat nor water nor acids are able to extract the alkali from sodium uranate, Na2U2O7, and therefore it is a true insoluble salt, of a yellow colour, and clearly indicates the acid character (although feeble) of uranic oxide. (2) The carbonates of the alkaline earths (for instance, barium carbonate) precipitate uranic oxide from its salts, as they do all the salts of feeble bases; for example, R2O3. (3) Thealkaline carbonates, when added to solutions of uranic salts, give aprecipitate, which is soluble inan excess of the reagent, and particularly so if the acid carbonates be taken. This is due to the fact that (4) the uranyl saltseasily form double saltswith the salts of the alkali metals, including the salts of ammonium. Uranium, in the form of these double salts, often gives salts of well-defined crystalline form, although the simple salts are little prone to appear in crystals. Such, for example, are the salts obtained by dissolving potassium uranate, K2U2O7, in acids, with the addition of potassium salts of the same acids. Thus, with hydrochloric acid and potassium chloride a well-formed crystalline salt, K2(UO2)Cl4,2H2O, belonging to the monoclinic system, is produced. This salt decomposes in dissolving in pure water. Among these double salts we may mention the double carbonate with the alkalis, R4(UO2)(CO3)3(equal to 2R2CO3+ UO2CO3); the acetates, R(UO2)(C2H3O2)3—for instance, the sodium salt, Na(UO2)(C2H3O2)3, and the potassium salt, K(UO2)(C2H3O2)3,H2O; the sulphates, R2(UO2)(SO4)3,2H2O, &c. In the preceding formula R = K, Na, NH4, or R2= Mg, Ba, &c.This property of giving comparatively stable double salts indicates feebly developed basic properties, because double salts are mainly formed by salts of distinctly basic metals (these form, as it were, the basic element of a double salt) and salts of feebly energetic bases (these form the acid element of a double salt), just as the former also give acid salts; the acid of the acid salts is replaced in the double salts by the salt of the feebly energetic base, which, like water, belongs to the class of intermediate bases. For this reason barium does not give double salts with alkalis as magnesium does, and this is why double salts are more easily formed by potassium than by lithium in the series of the alkali metals. (5) The most remarkable property, proving the feeble energy of uranic oxide as a base, is seen in the fact that when their composition is compared with that of other salts those of uranic oxidealways appear as basic salts. It is well known that a normal salt, R2X6, corresponds with the oxide R2O3, where X = Cl, NO3, &c., or X2= SO4, CO3, &c.; but there also exist basic salts of the same type where X = HO or X2= O. We saw salts of all kinds among the salts of aluminium, chromium, and others. With uranic oxide no salts are known of the types UX6(UCl6, U(SO4)3, alums, &c., are not known), nor even salts, U(HO)2X4or UOX4, but it always forms salts of the type U(HO)4X2, or UO2X2. Judging from the fact that nearly all the salts of uranic oxide retain water in crystallising from their solutions, and that this water is difficult to separate from them, it may be thought to be water of hydration. This is seen in part from the fact that the composition of many of the salts of uranic oxide may then be expressed without the presence of water of crystallisation; for instance, U(HO)4K2Cl4(and the salt of NH4, U(HO)4K2(SO4)2, U(HO)4(C2H3O2)2. Sodium uranyl acetate however does not contain water.[11]Uranyl nitrate, or uranium nitrate, UO2(NO3)2,6H2O, crystallises from its solutions in transparent yellowish-green prisms (from an acid solution), or in tabular crystals (from a neutral solution), which effloresce in the air and are easily soluble in water, alcohol, and ether, have a sp. gr. of 2·8, and fuse when heated, losing nitric acid and water in the process. If the salt itself (Berzelius) or its alcoholic solution (Malaguti) be heated up to the temperature at which oxides of nitrogen are evolved, there then remains a mass which, after being evaporated with water, leaves uranyl hydroxide, UO2(HO)2(sp. gr. 5·93), whilst if the salt be ignited there remains the dioxide, UO2, as a brick-red powder, which on further heating loses oxygen and forms the dark olive uranoso-uranic oxide, U3O8. The solution of the nitrate obtained from the ore is purified in the following manner: sulphurous anhydride is first passed through it in order to reduce the arsenic acid present into arsenious acid; the solution is then heated to 60°, and sulphuretted hydrogen passed through it; this precipitates the lead, arsenic, and tin, and certain other metals, as sulphides, insoluble in water and dilute nitric acid. This liquid is then filtered and evaporated with nitric acid to crystallisation, and the crystals are dissolved in ether. Or else the solution is first treated with chlorine in order to convert the ferrous chloride (produced by the action of the hydrogen sulphide) into ferric chloride, the oxides are then precipitated by ammonia, and the resultant precipitate, containing the oxides Fe2O3, UO3, and compounds of the latter with potash, lime, ammonia, and other bases present in the solution (the latter being due to the property of uranic oxide of combining with bases), is washed and dissolved in a strong, slightly-heated solution of ammonium carbonate, which dissolves the uranic oxide but not the ferric oxide. The solution is filtered, and on cooling deposits a well-crystallisinguranyl ammonium carbonate, UO2(NH4)4(CO3)3, in brilliant monoclinic crystals which on exposure to air slowly give off water, carbonic anhydride, and ammonia; the same decomposition is readily effected at 300°, the residue then consisting of uranic oxide. This salt is not very soluble in water, but is readily so in ammonium carbonate; it is obvious that it may readily be converted into all the other salts of oxides of uranium. Uranium salts are also purified in the form ofacetate, which is very sparingly soluble, and is therefore directly precipitated from a strong solution of the nitrate by mixing it with acetic acid.We may also mention theuranyl phosphate, HUPO6, which must be regarded as an orthophosphate in which two hydrogens are replaced by the radicle uranyl, UO2,i.e.as H(UO2)PO4. This salt is formed as a hydrated gelatinous yellow precipitate, on mixing a solution of uranyl nitrate with disodium phosphate. The precipitation occurs in the presence of acetic acid, but not in the presence of hydrochloric acid. If moreover an excess of an ammonium salt be present, the ammonia enters into the composition of the bright yellow gelatinous precipitate formed, in the proportion UO2NH4PO4. This precipitate is not soluble in water and acetic acid, and its solution in inorganic acids when boiled entirely expels all the phosphoric acid. This fact is taken advantage of for removing phosphoric acids from solutions—for instance, from those containing salts of calcium and magnesium.[12]Uranium dioxide, oruranyl, UO2, which is contained in the salts UO2X2, has the appearance and many of the properties of a metal. Uranic oxide may be regarded as uranyl oxide, (UO2)O, its salts as salts of this uranyl; its hydroxide, (UO2)H2O2, is constituted like CaH2O2. The green oxide of uranium, uranoso-uranic oxide (easily formed from uranic salts by the loss of oxygen), U3O8= UO2,2UO3, when ignited with charcoal or hydrogen (dry) gives a brilliant crystalline substance of sp. gr. about 11·0 (Urlaub), whose appearance resembles that of metals, and decomposes steam at a red heat with the evolution of hydrogen; it does not, however, decompose hydrochloric or sulphuric acid, but is oxidised by nitric acid. The same substance (i.e. uranium dioxide UO2) is also obtained by igniting the compound (UO2)K2Cl4in a stream of hydrogen, according to the equation UO2K4Cl4+ H2= UO2+ 2HCl + 2KCl. It was at first regarded as the metal. In 1841 Peligot found that it contained oxygen, because carbonic oxide and anhydride were evolved when it was ignited with charcoal in a stream of chlorine, and from 272 parts of the substance which was considered to be metal he obtained 382 parts of a volatile product containing 142 parts of chlorine. From this it was concluded that the substance taken contained an equivalent amount of oxygen. As 142 parts of chlorine correspond with 32 parts of oxygen, it followed that 272 - 32 = 240 parts of metal were combined in the substance with 32 parts of oxygen, and also in the chlorine compound obtained with 142 parts of chlorine. These calculations have been made for the now accepted atomic weight of uranium (U = 240,seeNote14). Peligot took another atomic weight, but this does not alter the principle of the argument.[13]Uranium tetrachloride, uranous chloride, UCl4, corresponds with uranous oxide as a base. It was obtained by Peligot by igniting uranic oxide mixed with charcoal in a stream ofdrychlorine: UO3+ 3C + 2Cl2= UCl4+ 3CO. This green volatile compound (Note12) crystallises in regular octahedra, is very hygroscopic, easily soluble in water, with the development of a considerable amount of heat, and no longer separates out from its solution in an anhydrous state, but disengages hydrochloric acid when evaporated. The solution of uranous chloride in water is green. It is also formed by the action of zinc and copper (forming cuprous chloride) on a solution of uranyl chloride, UO2Cl2, especially in the presence of hydrochloric acid and sal-ammoniac. Solutions of uranyl salts are converted into uranous salts by the action of various reducing agents, and among others by organic substances or by the action of light, whilst the salts UX4are converted into uranyl salts, UO2X2, by exposure to air or by oxidising agents. Solutions of the green uranyl salts act as powerful reducing agents, and give a brown precipitate of the uranous hydroxide, UH4O4, with potash and other alkalis. This hydroxide is easily soluble in acids but not in alkalis. On ignition it does not form the oxide UO2, because it decomposes water, but when the higher oxides of uranium are ignited in a stream of hydrogen or with charcoal they yield uranous oxide. Both it and the chloride UCl4, dissolve in strong sulphuric acid, forming a green salt, U(SO4)2,2H2O. The same salt, together with uranyl sulphate, UO2(SO4), is formed when the green oxide, U3O8, is dissolved in hot sulphuric acid. The salts obtained in the latter instance may be separated by adding alcohol to the solution, which is left exposed to the light; the alcohol reduces the uranyl salt to uranous salt, an excess of acid being required. An excess of water decomposes this salt, forming a basic salt, which is also easily produced under other circumstances, and contains UO(SO4),2H2O (which corresponds to the uranic salt).[14]The atomic weight of uranium was formerly taken as half the present one, U = 120, and the oxides U2O3, suboxide UO, and green oxide U3O4, were of the same types as the oxides of iron. With a certain resemblance to the elements of the iron group, uranium presents many points of distinction which do not permit its being grouped with them. Thus uranium forms a very stable oxide, U2O3(U = 120), but does not give the corresponding chloride U2Cl6(Roscoe, however, in 1874 obtained UCl5, like MoCl5and WCl5), and under those circumstances (the ignition of oxide of uranium mixed with charcoal, in a stream of chlorine), when the formation of this compound might be expected, it gives (U = 120) the chloride UCl2, which is characterised by its volatility; this is not a property, to such an extent, of any of the bichlorides, RCl2, of the iron group.The alteration or doubling of the atomic weight of uranium—i.e.the recognition of U = 240—was made for the first time in the first (Russian) edition of this work (1871), and in my memoir of the same year in Liebig'sAnnalen, on the ground that with an atomic weight 120, uranium could not be placed in the periodic system. I think it will not be superfluous to add the following remarks on this subject: (1) In the other groups (K—Rb—Cs, Ca—Sr—Ba, Cl—Br—I) the acid character of the oxides decreases and their basic character increases with the rise of atomic weight, and therefore we should expect to find the same in the group Cr—Mo—W—U, and if CrO3, MoO3, WO3be the anhydrides of acids then we indeed find a decrease in their acid character, and therefore uranium trioxide, UO3, should be a very feeble anhydride, but its basic properties should also be very feeble. Uranic oxide does indeed show these properties, as was pointed out above (Note10). (2) Chromium and its analogues, besides the oxides RO3, also form lower grades of oxidation RO2, R2O3, and the same is seen in uranium; it forms UO3, UO2, U2O3and their compounds. (3) Molybdenum and tungsten, in being reduced from RO3, easily and frequently give an intermediate oxide of a blue colour, and uranium shows the same property; giving the so-called green oxide which, according to present views, must be regarded as U3O8= UO22UO3, analogous to Mo3O8. (4) The higher chlorides, RCl6, possible for the elements of this group, are either unstable (WCl6) or do not exist at all (Cr); but there is one single lower volatile compound, which is decomposed by water, and liable to further reduction into a non-volatile chlorine product and the metal. The same is observed in uranium, which forms an easily volatile chloride, UCl4, decomposed by water. (5) The high sp. gr. of uranium (18·6) is explained by its analogy to tungsten (sp. gr. 19·1). (6) For uranium, as for chromium and tungsten, yellow tints predominate in the form RO3, whilst the lower forms are green and blue. (7) Zimmermann (1881) determined the vapour densities of uranous bromide, UBr4, and chloride, UCl4(19·4 and 13·2), and they were found to correspond to the formulæ given above—that is, they confirmed the higher atomic weight U = 240. Roscoe, a great authority on the metals of this group, was the first to accept the proposed atomic weight of uranium, U = 240, which since Zimmermann's work has been generally recognised.[15]Uranium glass, obtained by the addition of the yellow salt K2U2O7to glass, has a green yellow fluorescence, and is sometimes employed for ornaments; it absorbs the violet rays, like the other salts of uranic oxide—that is, it possesses an absorption spectrum in which the violet rays are absent. The index of refraction of the absorbed rays is altered, and they are given out again as greenish-yellow rays; hence, compounds of uranic acid, when placed in the violet portion of the spectrum, emit a greenish-yellow light, and this forms one of the best examples (another is found in a solution of quinine sulphate) of the phenomenon of fluorescence. The rays of light which pass through uranic compounds do not contain the rays which excite the phenomena of fluorescence and of chemical transformation, as the researches of Stokes prove.[16]The comparison of potassium permanganate with potassium perchlorate, or of potassium manganate with potassium sulphate, shows directly that many of the physical and chemical properties of substances do not depend on the nature of the elements, but on the atomic types in which they appear, on the kind of movements, or on the positions in which the atoms forming the molecule occur.[17]If, however, we compare the spectra (Vol. I. p.565) of chlorine, bromine, and iodine with that of manganese, a certain resemblance or analogy is to be found connecting manganese both to iron and to chlorine, bromine, and iodine.[18]The name ‘peroxide’ should only be retained for thosehighestoxides (and MnO2stands between MnO and MnO3) which either by a direct method of double decomposition are able to give hydrogen peroxide or contain a larger proportion of oxygen than the base or the acid, just as hydrogen peroxide contains more oxygen than water. Their type will be H2O2, and they are exemplified by barium peroxide, BaO2, and sulphur peroxide, S2O7, &c. Such a dioxide as MnO2is, in all probability, a salt—that is, a manganous manganate, MnO3MnO, and also, as a basic salt of a feeble base, capable of combining with alkalis and acids. Hence the name of manganese peroxide should be abandoned, and replaced by manganese dioxide. PbO2is better termed lead dioxide than peroxide. Bisulphide of manganese, MnS2, corresponding to iron pyrites, FeS2, sometimes occurs in nature in fine octahedra (and cube combinations), for instance, in Sicily; it is called Hauerite.[18 bis]On comparing the manganates with the permanganates—for example, K2MnO4with KMnO4—we find that they differ in composition by the abstraction of one equivalent of the metal. Such a relation in composition produced by oxidation is of frequent occurrence—for instance, K4Fe(CN)6in oxidising gives K3Fe(CN)6; H2SO4in oxidising gives persulphuric acid, HSO4, or H2S7O8; H2O forms HO or H2O2, &c.[19]In the preparation of oxygen from the dioxide by means of H2SO4, MnSO4is formed; in the preparation of chlorine from HCl and MnO2, MnCl2is obtained. These two manganous salts may be taken as examples of compounds MnX2. Manganous sulphate generally contains various impurities, and also a large amount of iron salt (from the native MnO2), from which it cannot be freed by crystallisation. Their removal may, however, be effected by mixing a portion of the liquid with a solution of sodium carbonate; a precipitate of manganous carbonate is then formed. This precipitate is collected and washed, and then added to the remaining mass of the impure solution of manganous sulphate; on heating the solution with this precipitate, the whole of the iron is precipitated as oxide. This is due to the fact that in the solution of the manganese dioxide in sulphuric acid the whole of the iron is converted into the ferric state (because the dioxide acts as an oxidising agent), which, as an exceedingly feeble base precipitated by calcium carbonate and other kindred salts, is also precipitated by manganous carbonate. After being treated in this manner, the solution of manganous sulphate is further purified by crystallisation. If it be a bright red colour, it is due to the presence of higher grades of oxidation of manganese; they may be destroyed by boiling the solution, when the oxygen from the oxides of manganese is evolved and a very faintly coloured solution of manganous sulphate is obtained. This salt is remarkable for the facility with which it gives various combinations with water. By evaporating the almost colourless solution ofmanganous sulphateat very low temperatures, and by cooling the saturated solution at about 0°, crystals are obtained containing 7 atoms of water of crystallisation, MnSO4,7H2O, which are isomorphous with cobaltous and ferrous sulphates. These crystals, even at 10°, lose 5 p.c. of water, and completely effloresce at 15°, losing about 20 p.c. of water. By evaporating a solution of the salt at the ordinary temperature, but not above 20°, crystals are obtained containing 5 mol. H2O, which are isomorphous with copper sulphate; whilst if the crystallisation be carried on between 20° and 30°, large transparent prismatic crystals are formed containing 4 mol. H2O (see Nickel). A boiling solution also deposits these crystals together with crystals containing 3 mol. H2O, whilst the first salt, when fused and boiled with alcohol, gives crystals containing 2 mol. H2O. Graham obtained a monohydrated salt by drying the salt at about 200°. The last atom of water is eliminated with difficulty, as is the case with all salts like MnSO4nH2O. The crystals containing a considerable amount of water are rose-coloured, and the anhydrous crystals are colourless. The solubility of MnSO4,4H2O (Chapter I., Note24) per 100 parts of water is: at 10°, 127 parts; at 37°·5, 149 parts; at 75°, 145 parts; and at 101°, 92 parts. Whence it is seen that at the boiling-point this salt is less soluble than at lower temperatures, and therefore a solution saturated at the ordinary temperature becomes turbid when boiled. Manganous sulphate, being analogous to magnesium sulphate, is decomposed, like the latter, when ignited, but it does not then leave manganous oxide, but the intermediate oxide, Mn3O4. It gives double salts with the alkali sulphates. With aluminium sulphate it forms fine radiated crystals, whose composition resembles that of the alums—namely, MnAl2(SO4)4,24H2O. This salt is easily soluble in water, and occurs in nature.Manganous chloride, MCl2, crystallises with 4 mol. H2O, like the ferrous salt, and not with 6 mol. H2O like many kindred salts—for example, those of cobalt, calcium, and magnesium; 100 parts of water dissolve 38 parts of the anhydrous salt at 10° and 55 parts at 62°. Alcohol also dissolves manganous chloride, and the alcoholic solution burns with a red flame. This salt, like magnesium chloride, readily forms double salts. A solution of borax gives a dirty rose-coloured precipitate having the composition MnH4(BO3)2H2O, which is used as a drier in paint-making. Potassium cyanide produces a yellowish-grey precipitate, MnC2N2, with manganous salts, soluble in an excess of the reagent, a double salt, K4MnC6N6, corresponding with potassium ferrocyanide, being formed. On evaporation of this solution, a portion of the manganese is oxidised and precipitated, whilst a salt corresponding to Gmelin's red salt, K3,MnC6N6(seeChapterXXII.), remains in solution. Sulphuretted hydrogen does not precipitate salts of manganese, not even the acetate, but ammonium sulphide gives a flesh-coloured precipitate, MnS; at 320° this sulphide of manganese passes into a green variety (Antony). Oxalic acid in strong solutions of manganous salts gives a white precipitate of the oxalate, MnC2O4. This precipitate is insoluble in water, and is used for the preparation of manganous oxide itself because it decomposes like oxalic acid when ignited (in a tube without access of air), with the formation of carbonic anhydride, carbonic oxide, and manganous oxide.Manganous oxidethus obtained is a green powder, which however oxidises with such facility that it burns in air when brought into contact with an incandescent substance, and passes into the red intermediate oxide Mn3O4. In solutions of manganous salts, alkalis produce a precipitate of the hydroxide MnH2O2, which rapidly absorbs oxygen in the presence of air and gives the brown intermediate oxide, or, more correctly speaking, its hydrate.Manganous oxide, besides being obtained by the above-described method from manganous oxalate, may also be obtained by igniting the higher oxides in a stream of hydrogen, and also from manganese carbonate. The manganous oxide ignited in the presence of hydrogen acquires a great density, and is no longer so easily oxidised. It may also be obtained in a crystalline form, if during the ignition of the carbonate or higher oxide a trace of dry hydrochloric acid gas be passed into the current of hydrogen. It is thus obtained in the form of transparent emerald green crystals of the regular system, and in this state is easily soluble in acids.Manganous oxide in oxidising gives thered oxide of manganese, Mn5O4. This is the most stable of all the oxides of manganese; it is not only stable at the ordinary but also at a high temperature—that is, it does not absorb or disengage oxygen spontaneously. When ignited, all the higher oxides of manganese pass into it by losing oxygen, and manganous oxide by absorbing oxygen. This oxide does not give any distinct salts, but it dissolves in sulphuric acid, forming a dark red solution, which contains both manganous and manganic (of theoxide, Mn2O3) sulphates. The latter with potassium sulphate gives a manganese alum, in which the alumina is replaced by its isomorphous oxide of manganese. But this alum, like the solution of the intermediate oxide in sulphuric acid, evolves oxygen and leaves a manganous salt when slightly heated.Manganese dioxideis still less basic than the oxide, and disengages oxygen or a halogen in the presence of acids, forming manganous salts, like the oxide. However, if it be suspended in ether, and hydrochloric acid gas passed into the mixture, which is kept cool, the ether acquires a green colour, owing to the formation of tetrachloride of manganese, MnCl4, corresponding with the dioxide which passes into solution. It is however very unstable, being exceedingly easily decomposed with the evolution of chlorine. The corresponding fluoride, MnF4, obtained by Nicklés is much more stable. At all events, manganese dioxide does not exhibit any well-defined basic character, but has rather an acid character, which is particularly shown in the compounds MnF4and MnCl4just mentioned, and in the property of manganese dioxide of combining with alkalis. If the higher grades of oxidation of manganese be deoxidised in the presence of alkalis, they frequently give the dioxide combined with the alkali—for example, in the presence of potash a compound is formed which contains K2O,5MnO2, which shows the weak acid character of this oxide. When ignited in the presence of sodium compounds manganese dioxide frequently forms Na2O,8MnO2and Na2O,12MnO2, and lime when heated with MnO2gives from CaO,3MnO2to (CaO)2,MnO2(Rousseau) according to the temperature. Besides which, perhaps, MnO2is a saline compound, containing MnOMnO3or (MnO)3Mn2O7, and there are reactions which support such a view (Spring, Richards, Traube, and others); for instance it is known that manganous chloride and potassium permanganate give the dioxide in the presence of alkalis.Manganese dioxide may be obtained from manganous salts by the action of oxidising agents. If manganous hydroxide or carbonate be shaken up in water through which chlorine is passed, the hypochlorite of the metal is not formed, as is the case with certain other oxides, but manganese dioxide is precipitated: 2MnO2H2+ Cl2= MnCl2+ MnO2,H2O + H2O. Owing to this fact, hypochlorites in the presence of alkalis and acetic acid when added to a solution of manganous salts give hydrated manganese dioxide, as was mentioned above. Manganous nitrate also leaves manganese dioxide when heated to 200°. It is also obtained from manganous and manganic salts of the alkalis, when they are decomposed in the presence of a small amount of acid; the practical method of converting the salts MnX2into the higher grades of oxidation is given in Chapter II., Note6.[20]Other chemists have obtained manganese by different methods, and attributed different properties to it. This difference probably depends on the presence of carbon in different proportions. Deville obtained manganese by subjecting the pure dioxide, mixed with pure charcoal (from burnt sugar), to a strong heat in a lime crucible until the resultant metal fused. The metal obtained had a rose tint, like bismuth, and like it was very brittle, although exceedingly hard. It decomposed water at the ordinary temperature. Brunner obtained manganese having a specific gravity of about 7·2, which decomposed water very feebly at the ordinary temperature, did not oxidise in air, and was capable of taking a bright polish, like steel; it had the grey colour of cast iron, was very brittle, and hard enough to scratch steel and glass, like a diamond. Brunner's method was as follows: He decomposed the manganese fluoride (obtained as a soluble compound by the action of hydrofluoric acid on manganese carbonate) with sodium, by mixing these substances together in a crucible and covering the mixture with a layer of salt and fluor spar; after which the crucible was first gradually heated until the reaction began, and then strongly heated in order to fuse the metal separated. Glatzel (1889) obtained 25 grms. of manganese, having a grey colour and sp. gr. 7·39, by heating a mixture of 100 grms. of MnCl2with 200 grms. KCl and 15 grms. Mg to a bright white heat. Moissan and others, by heating the oxides of manganese with carbon in the electric furnace, obtained carbides of manganese—for example, Mn3C—and remarked that the metal volatilised in the heat of the voltaic arc. Metallic manganese is, however, not prepared on a large scale, but only its alloys with carbon (they readily and rapidly oxidise) andferro-manganeseor a coarsely crystalline alloy of iron, manganese and carbon, which is smelted in blast-furnaces like pig-iron (seeChapterXXII.) This ferro-manganese is employed in the manufacture of steel by Bessemer's and other processes (see ChapterXXII.) and for the manufacture of manganese bronze. However, in America, Green and Wahl (1895) obtained almost pure metallic manganese on a large scale. They first treat the ore of MnO2with 30 p.c. sulphuric acid (which extracts all the oxides of iron present in the ore), and then heat it in a reducing flame to convert it into MnO, which they mix with a powder of Al, lime and CaF2(as a flux), and heat the mixture in a crucible lined with magnesia; a reaction immediately takes place at a certain temperature, and a metal of specific gravity 7·3 is obtained, which only contains a small trace of iron.Manganese gives two compounds withnitrogen, Mn5N2and Mn3N2. They were obtained by Prelinger (1894) from the amalgam of manganese Mn2Hg5(obtained on a mercury anode by the action of an electric current upon a solution of MnCl2); the mercury may be removed from this amalgam by heating it in an atmosphere of hydrogen, and then metallic manganese is obtained as a grey porous mass of specific gravity 7·42. If this amalgam be heated in dry nitrogen it gives Mn5N2(grey powder, sp. gr. 6·58), but if heated in an atmosphere of NH3it gives (as also does Mn5N2) Mn3N2, (a dark mass with a metallic lustre, sp. gr. 6·21), which, when heated in nitrogen is converted into Mn5N2, and if heated in hydrogen evolves NH3and disengages hydrogen from a solution of NH4Cl. At all events, manganese is a metal which decomposes water more easily than iron, nickel, and cobalt.[21]Volume I. p. 157, Note7.[22]It was known to the alchemists by this name, but the true explanation of the change in colour is due to the researches of Chevillot, Edwards, Mitscherlich, and Forchhammer. The change in colour of potassium manganate is due to its instability and to its splitting up into two other manganese compounds, a higher and a lower: 3MnO3= Mn2O7+ MnO2. Manganese trioxide is really decomposed in this manner by the action of water (see later): 3MnO3+ H2O = 2MnHO4+ MnO2(Franke, Thorpe, and Humbly). The instability of the salt is proved by the fact of its being deoxidised by organic matter, with the formation of manganese dioxide and alkali, so that, for instance, a solution of this salt cannot be filtered through paper. The presence of an excess of alkali increases the stability of the salt; when heated it breaks up in the presence of water, with the evolution of oxygen.The method of preparingpotassium permanganatewill be understood from the above. There are many recipes for preparing this substance, as it is now used in considerable quantities both for technical and laboratory purposes. But in all cases the essence of the methods is one and the same: a mixture of alkali with any oxide of manganese (even manganous hydroxide, which may be obtained from manganous chloride) is first heated in the presence of air or of an oxidising substance (for the sake of rapidity, with potassium chlorate); the resultant mass is then treated with water and heated, when manganese dioxide is precipitated and potassium permanganate remains in solution. This solution may be boiled, as the liquid will contain free alkali; but the solution cannot be evaporated to dryness, because a strong solution, as well as the solid salt, is decomposed by heat.By adding a dilute solution of manganous sulphate to a boiling mixture of lead dioxide and dilute nitric acid, the whole of the manganese may be converted into permanganic acid (Crum).[22 bis]The solution of this salt with an excess of impure commercial alkali generally acquires a green tint.[23]A solution of potassium permanganate gives a beautiful absorption spectrum (ChapterXIII.) If the light in passing through this solution loses a portion of its rays in it (if one may so account for it), this is partially explained by the increased oxidising power which the solution then acquires. We may here also remark that a dilute solution of permanganate of potassium forms a colourless solution with nickel salts, because the green colour of the solution of nickel salts is complementary to the red. Such a decolorised solution, containing a large proportion of nickel and a small proportion of manganese, decomposes after a time, throws down a precipitate, and re-acquires the green colour proper to the nickel salts. The addition of a solution of a cobalt salt (rose-red) to the nickel salt also destroys the colour of both salts.[24]If sulphuric acid is allowed to act on potassium permanganate without any special precautions, a large amount of oxygen is evolved (it may even explode and inflame), and a violet spray of the decomposing permanganic acid is given off. But if the pure salt (i.e.free from chlorine) be dissolved in pure well-cooled sulphuric acid, without any rise in temperature, a green-coloured liquid settles at the bottom of the vessel. This liquid does not contain any sulphuric acid, and consists of permanganic anhydride, Mn2O7(Aschoff, Terreil). It is impossible to prepare any considerable quantity of the anhydride by this method, as it decomposes with an explosion as it collects, evolving oxygen and leaving red oxide of manganese.Permanganic anhydride, Mn2O7, in dissolving in sulphuric acid, gives a green solution, which (according to Franke, 1887) contains a compound Mn2SO10= (MnO3)2SO4—that is, sulphuric acid in which both hydrogens are replaced by the group MnO3, which is combined with OK in permanganate of potassium. This mixture with a small quantity of water gives Mn2O7, according to the equation: (MnO3)2SO4+ H2O = H2SO4+ Mn2O7, and when heated to 30° it givesmanganese trioxide, (MnO3)2SO4+ H2O = 2MnO2+ H2SO4+ O. Pure manganese trioxide is obtained if the solution of (MnO3)2SO4be poured in drops on to sodium carbonate. Then, together with carbonic anhydride, a spray of manganese trioxide passes over, which may be collected in a well-cooled receiver, and this shows that the reaction proceeds according to the equation: (MnO3)2SO4+ Na2CO3= Na2SO4+ 2MnO3+ CO2+ O (Thorpe). The trioxide is decomposed by water, forming manganese dioxide and a solution ofpermanganic acid: 3MnO3+ H2O = MnO2+ 2HMnO4. The same acid is obtained by dissolving permanganic anhydride in water.Barium permanganate when treated with sulphuric acid gives the same acid. This barium salt may be prepared by the action of barium chloride on the difficultly soluble silver permanganate, AgMnO4, which is precipitated on mixing a strong solution of the potassium salt with silver nitrate. The solution of permanganic acid forms a bright red liquid which reflects a dark violet tint. A dilute solution has exactly the same colour as that of the potassium salt. It deposits manganese dioxide when exposed to the action of light, and also when heated above 60°, and this proceeds the more rapidly the more dilute the solution. It shows its oxidising properties in many cases, as already mentioned. Even hydrogen gas is absorbed by a solution of permanganic acid; and charcoal and sulphur are also oxidised by it, as they are by potassium permanganate. This may be taken advantage of in analysing gunpowder, because when it is treated with a solution of potassium permanganate, all the sulphur is converted into sulphuric acid and all the charcoal into carbonic anhydride. Finely-divided platinum immediately decomposes permanganic acid. With potassium iodide it liberates iodine (which may afterwards be oxidised into iodic acid) (Mitscherlich, Fromherz, Aschoff, and others). Ammonia does not form a corresponding salt with free permanganic acid, because it is oxidised with evolution of nitrogen. The oxidising action of permanganic acid in a strong solution may be accompanied by flame and the formation of violet fumes of permanganic acid; thus a strong solution of it takes fire when brought into contact with paper, alcohol, alkaline sulphides, fats, &c.We may add that, according to Franke, 1 part of potassium permanganate with 13 parts of sulphuric acid at 100° gives brown crystals of the salt Mn2(SO4)3,H2SO4,4H2O, which gives a precipitate of hydrated manganese dioxide, H2MnO3= MnO2H2O, when treated with water.Spring, by precipitating potassium permanganate with sodium sulphite and washing the precipitate by decantation, obtained a soluble colloidal manganese oxide, whose composition was the mean between Mn2O3and MnO2—namely, Mn2O3,4(MnO2H2O).
[9]Moissan (1893) studied the compounds of Mo and W formed with carbon in the electrical furnace (they are extremely hard) from a mixture of the anhydrides and carbon. Poleck and Grützner obtained definite compounds FeW2and FeW2C3for tungsten. Metallic W and Mo displace Ag from its solutions but not Pb. There is reason for believing that the sp. gr. of pure molybdenum is higher than that (8·5) generally ascribed to it.
[9]Moissan (1893) studied the compounds of Mo and W formed with carbon in the electrical furnace (they are extremely hard) from a mixture of the anhydrides and carbon. Poleck and Grützner obtained definite compounds FeW2and FeW2C3for tungsten. Metallic W and Mo displace Ag from its solutions but not Pb. There is reason for believing that the sp. gr. of pure molybdenum is higher than that (8·5) generally ascribed to it.
[9 bis]We may conclude our description of tungsten and molybdenum by stating that their sulphur compounds have an acid character, like carbon bisulphide or stannic sulphide. If sulphuretted hydrogen be passed through a solution of a molybdate it does not give a precipitate unless sulphuric acid be present, when a dark brown precipitate ofmolybdenum trisulphide, MoS3, is formed. When this sulphide is ignited without access of air it gives the bisulphide MoS2; the latter is not able to combine with potassium sulphide like the trisulphide MoS3, which forms a salt, K2MoS4, corresponding with K2MoO4. This is soluble in water, and separates out from its solution in red crystals, which have a metallic lustre and reflect a green light. It is easily obtained by heating the native bisulphide, MoS2, with potash, sulphur, and a small amount of charcoal, which serves for deoxidising the oxygen compounds. Tungsten gives similar compounds, R2WS4, where R = NH4, K, Na. They are decomposed by acids, with the separation of tungsten trisulphide, WS3, and molybdenum trisulphide, MoS3. Rideal (1892) obtained W2N3by heating WO3in NH3. This compound exhibited the general properties of metallic nitrides.
[9 bis]We may conclude our description of tungsten and molybdenum by stating that their sulphur compounds have an acid character, like carbon bisulphide or stannic sulphide. If sulphuretted hydrogen be passed through a solution of a molybdate it does not give a precipitate unless sulphuric acid be present, when a dark brown precipitate ofmolybdenum trisulphide, MoS3, is formed. When this sulphide is ignited without access of air it gives the bisulphide MoS2; the latter is not able to combine with potassium sulphide like the trisulphide MoS3, which forms a salt, K2MoS4, corresponding with K2MoO4. This is soluble in water, and separates out from its solution in red crystals, which have a metallic lustre and reflect a green light. It is easily obtained by heating the native bisulphide, MoS2, with potash, sulphur, and a small amount of charcoal, which serves for deoxidising the oxygen compounds. Tungsten gives similar compounds, R2WS4, where R = NH4, K, Na. They are decomposed by acids, with the separation of tungsten trisulphide, WS3, and molybdenum trisulphide, MoS3. Rideal (1892) obtained W2N3by heating WO3in NH3. This compound exhibited the general properties of metallic nitrides.
[9 tri]When peroxide of hydrogen acts upon a solution of potassium molybdate well-formed yellow crystals belonging to the triclinic system separate out in the cold. When these crystals are heated in vacuo they first lose water and then decompose, leaving a residue composed of the salt originally taken. They are soluble in water but insoluble in alcohol. Their composition is represented by the formula K2Mo2O82H2O. An ammonium salt is obtained by evaporating peroxide of hydrogen with ammonium molybdate. The following salts have also been obtained by the action of peroxide of hydrogen upon the corresponding molybdates: Na2Mo2O66H2O—in yellow prismatic crystals; MgMo2O810H2O—stellar needles; BaMoO82H2O—in microscopic yellow octahedra. A corresponding sodium pertungstate has been obtained by Péchard by boiling sodium tungstate with a solution of peroxide of hydrogen for several minutes. The solution rapidly turns yellow, and no longer gives a precipitate of tungstic anhydride when treated with nitric acid. When evaporated in vacuo the solution leaves a thick syrupy liquid from which ray-like crystals separate out; these crystals are more soluble in water than the salt originally taken. When heated they also lose water and oxygen. Their composition answers to the formula M2W2O82H2O, where M = Na, NH4, &c. The permolybdates and pertungstates have similar properties. When treated with oxygen acids they give peroxide of hydrogen, and disengage chlorine and iodine from hydrochloric acid and potassium iodide.Piccini (1891) showed that peroxide of hydrogen not only combines with the oxygen compounds of Mo and W, but also with their fluo-compounds, among which ammonium fluo-molybdate MoO2F22NH4and others have long been known. (A few new salts of similar composition have been obtained by F. Moureu in 1893.) The action of peroxide of hydrogen upon these compounds gives salts containing a larger amount of oxygen; for instance, a solution of MoO2F22KFH2O with peroxide of hydrogen gives a yellow solution which after cooling separates out yellow crystalline flakes of MoO3F22KFH2O, resembling the salt originally taken in their external appearance. By employing a similar method Piccini also obtained: MoO3F22RbFH2O—yellow monoclinic crystals; MoO3F2,2CsFH2O,—yellow flakes, and the corresponding tungstic compounds. All these salts react like peroxide of hydrogen.In speaking of these compounds I for my part think it may be well to call attention to the fact that, in the first place, the composition of Piccini's oxy-fluo compounds does not correspond to that of permolybdic and pertungstic acid. If the latter be expressed by formulæ with one equivalent of an element, they will be HMoO4and HWO4, and the oxy-fluo form corresponding to them should have the composition MoO3F and WO3F while it contains MO3F2and WO3F2,i.e.answers as it were to a higher degree of oxidation, MoH2O3and W3HO3. But if permolybdic acid be regarded as 2MoO3+ H2O2,i.e.as containing the elements of peroxide of hydrogen, then Piccini's compound will also be found to contain the original salts + H2O; for example, from MoO2F22KFH2O there is obtained a compound MoO2F22KFH2O2,i.e.instead of H2O they contain H2O2. In the second place the capacity of the salts of molybdenum and tungsten to retain a further amount of oxygen or H2O2probably bears some relation to their property of giving complex acids and of polymerising which has been considered in Note8 bis. There is, however, a great chemical interest in the accumulation of data respecting these high peroxide compounds corresponding to molybdic and tungstic acids. With regard to the peroxide form of uranium,seeChapter XX., Note66.
[9 tri]When peroxide of hydrogen acts upon a solution of potassium molybdate well-formed yellow crystals belonging to the triclinic system separate out in the cold. When these crystals are heated in vacuo they first lose water and then decompose, leaving a residue composed of the salt originally taken. They are soluble in water but insoluble in alcohol. Their composition is represented by the formula K2Mo2O82H2O. An ammonium salt is obtained by evaporating peroxide of hydrogen with ammonium molybdate. The following salts have also been obtained by the action of peroxide of hydrogen upon the corresponding molybdates: Na2Mo2O66H2O—in yellow prismatic crystals; MgMo2O810H2O—stellar needles; BaMoO82H2O—in microscopic yellow octahedra. A corresponding sodium pertungstate has been obtained by Péchard by boiling sodium tungstate with a solution of peroxide of hydrogen for several minutes. The solution rapidly turns yellow, and no longer gives a precipitate of tungstic anhydride when treated with nitric acid. When evaporated in vacuo the solution leaves a thick syrupy liquid from which ray-like crystals separate out; these crystals are more soluble in water than the salt originally taken. When heated they also lose water and oxygen. Their composition answers to the formula M2W2O82H2O, where M = Na, NH4, &c. The permolybdates and pertungstates have similar properties. When treated with oxygen acids they give peroxide of hydrogen, and disengage chlorine and iodine from hydrochloric acid and potassium iodide.
Piccini (1891) showed that peroxide of hydrogen not only combines with the oxygen compounds of Mo and W, but also with their fluo-compounds, among which ammonium fluo-molybdate MoO2F22NH4and others have long been known. (A few new salts of similar composition have been obtained by F. Moureu in 1893.) The action of peroxide of hydrogen upon these compounds gives salts containing a larger amount of oxygen; for instance, a solution of MoO2F22KFH2O with peroxide of hydrogen gives a yellow solution which after cooling separates out yellow crystalline flakes of MoO3F22KFH2O, resembling the salt originally taken in their external appearance. By employing a similar method Piccini also obtained: MoO3F22RbFH2O—yellow monoclinic crystals; MoO3F2,2CsFH2O,—yellow flakes, and the corresponding tungstic compounds. All these salts react like peroxide of hydrogen.
In speaking of these compounds I for my part think it may be well to call attention to the fact that, in the first place, the composition of Piccini's oxy-fluo compounds does not correspond to that of permolybdic and pertungstic acid. If the latter be expressed by formulæ with one equivalent of an element, they will be HMoO4and HWO4, and the oxy-fluo form corresponding to them should have the composition MoO3F and WO3F while it contains MO3F2and WO3F2,i.e.answers as it were to a higher degree of oxidation, MoH2O3and W3HO3. But if permolybdic acid be regarded as 2MoO3+ H2O2,i.e.as containing the elements of peroxide of hydrogen, then Piccini's compound will also be found to contain the original salts + H2O; for example, from MoO2F22KFH2O there is obtained a compound MoO2F22KFH2O2,i.e.instead of H2O they contain H2O2. In the second place the capacity of the salts of molybdenum and tungsten to retain a further amount of oxygen or H2O2probably bears some relation to their property of giving complex acids and of polymerising which has been considered in Note8 bis. There is, however, a great chemical interest in the accumulation of data respecting these high peroxide compounds corresponding to molybdic and tungstic acids. With regard to the peroxide form of uranium,seeChapter XX., Note66.
[10]Uranium trioxide, or uranic oxide, shows its feeble basic and acid properties in a great number of its reactions. (1) Solutions of uranic salts give yellow precipitates with alkalis, but these precipitates do not contain the hydrate of the oxide, but compounds of it with bases; for example, 2UO2(NO3)2+ 6KHO = 4KNO3+ 3H2O + K2U2O7. There are otherurano-alkali compoundsof the same constitution; for example, (NH4)2U2O7(known commercially as uranic oxide), MgU2O7, BaU2O7. They are the analogues of the dichromates. Sodium uranate is the most generally used under the name of uranium yellow, Na2U2O7. It is used for imparting the characteristic yellow-green tint to glass and porcelain. Neither heat nor water nor acids are able to extract the alkali from sodium uranate, Na2U2O7, and therefore it is a true insoluble salt, of a yellow colour, and clearly indicates the acid character (although feeble) of uranic oxide. (2) The carbonates of the alkaline earths (for instance, barium carbonate) precipitate uranic oxide from its salts, as they do all the salts of feeble bases; for example, R2O3. (3) Thealkaline carbonates, when added to solutions of uranic salts, give aprecipitate, which is soluble inan excess of the reagent, and particularly so if the acid carbonates be taken. This is due to the fact that (4) the uranyl saltseasily form double saltswith the salts of the alkali metals, including the salts of ammonium. Uranium, in the form of these double salts, often gives salts of well-defined crystalline form, although the simple salts are little prone to appear in crystals. Such, for example, are the salts obtained by dissolving potassium uranate, K2U2O7, in acids, with the addition of potassium salts of the same acids. Thus, with hydrochloric acid and potassium chloride a well-formed crystalline salt, K2(UO2)Cl4,2H2O, belonging to the monoclinic system, is produced. This salt decomposes in dissolving in pure water. Among these double salts we may mention the double carbonate with the alkalis, R4(UO2)(CO3)3(equal to 2R2CO3+ UO2CO3); the acetates, R(UO2)(C2H3O2)3—for instance, the sodium salt, Na(UO2)(C2H3O2)3, and the potassium salt, K(UO2)(C2H3O2)3,H2O; the sulphates, R2(UO2)(SO4)3,2H2O, &c. In the preceding formula R = K, Na, NH4, or R2= Mg, Ba, &c.This property of giving comparatively stable double salts indicates feebly developed basic properties, because double salts are mainly formed by salts of distinctly basic metals (these form, as it were, the basic element of a double salt) and salts of feebly energetic bases (these form the acid element of a double salt), just as the former also give acid salts; the acid of the acid salts is replaced in the double salts by the salt of the feebly energetic base, which, like water, belongs to the class of intermediate bases. For this reason barium does not give double salts with alkalis as magnesium does, and this is why double salts are more easily formed by potassium than by lithium in the series of the alkali metals. (5) The most remarkable property, proving the feeble energy of uranic oxide as a base, is seen in the fact that when their composition is compared with that of other salts those of uranic oxidealways appear as basic salts. It is well known that a normal salt, R2X6, corresponds with the oxide R2O3, where X = Cl, NO3, &c., or X2= SO4, CO3, &c.; but there also exist basic salts of the same type where X = HO or X2= O. We saw salts of all kinds among the salts of aluminium, chromium, and others. With uranic oxide no salts are known of the types UX6(UCl6, U(SO4)3, alums, &c., are not known), nor even salts, U(HO)2X4or UOX4, but it always forms salts of the type U(HO)4X2, or UO2X2. Judging from the fact that nearly all the salts of uranic oxide retain water in crystallising from their solutions, and that this water is difficult to separate from them, it may be thought to be water of hydration. This is seen in part from the fact that the composition of many of the salts of uranic oxide may then be expressed without the presence of water of crystallisation; for instance, U(HO)4K2Cl4(and the salt of NH4, U(HO)4K2(SO4)2, U(HO)4(C2H3O2)2. Sodium uranyl acetate however does not contain water.
[10]Uranium trioxide, or uranic oxide, shows its feeble basic and acid properties in a great number of its reactions. (1) Solutions of uranic salts give yellow precipitates with alkalis, but these precipitates do not contain the hydrate of the oxide, but compounds of it with bases; for example, 2UO2(NO3)2+ 6KHO = 4KNO3+ 3H2O + K2U2O7. There are otherurano-alkali compoundsof the same constitution; for example, (NH4)2U2O7(known commercially as uranic oxide), MgU2O7, BaU2O7. They are the analogues of the dichromates. Sodium uranate is the most generally used under the name of uranium yellow, Na2U2O7. It is used for imparting the characteristic yellow-green tint to glass and porcelain. Neither heat nor water nor acids are able to extract the alkali from sodium uranate, Na2U2O7, and therefore it is a true insoluble salt, of a yellow colour, and clearly indicates the acid character (although feeble) of uranic oxide. (2) The carbonates of the alkaline earths (for instance, barium carbonate) precipitate uranic oxide from its salts, as they do all the salts of feeble bases; for example, R2O3. (3) Thealkaline carbonates, when added to solutions of uranic salts, give aprecipitate, which is soluble inan excess of the reagent, and particularly so if the acid carbonates be taken. This is due to the fact that (4) the uranyl saltseasily form double saltswith the salts of the alkali metals, including the salts of ammonium. Uranium, in the form of these double salts, often gives salts of well-defined crystalline form, although the simple salts are little prone to appear in crystals. Such, for example, are the salts obtained by dissolving potassium uranate, K2U2O7, in acids, with the addition of potassium salts of the same acids. Thus, with hydrochloric acid and potassium chloride a well-formed crystalline salt, K2(UO2)Cl4,2H2O, belonging to the monoclinic system, is produced. This salt decomposes in dissolving in pure water. Among these double salts we may mention the double carbonate with the alkalis, R4(UO2)(CO3)3(equal to 2R2CO3+ UO2CO3); the acetates, R(UO2)(C2H3O2)3—for instance, the sodium salt, Na(UO2)(C2H3O2)3, and the potassium salt, K(UO2)(C2H3O2)3,H2O; the sulphates, R2(UO2)(SO4)3,2H2O, &c. In the preceding formula R = K, Na, NH4, or R2= Mg, Ba, &c.This property of giving comparatively stable double salts indicates feebly developed basic properties, because double salts are mainly formed by salts of distinctly basic metals (these form, as it were, the basic element of a double salt) and salts of feebly energetic bases (these form the acid element of a double salt), just as the former also give acid salts; the acid of the acid salts is replaced in the double salts by the salt of the feebly energetic base, which, like water, belongs to the class of intermediate bases. For this reason barium does not give double salts with alkalis as magnesium does, and this is why double salts are more easily formed by potassium than by lithium in the series of the alkali metals. (5) The most remarkable property, proving the feeble energy of uranic oxide as a base, is seen in the fact that when their composition is compared with that of other salts those of uranic oxidealways appear as basic salts. It is well known that a normal salt, R2X6, corresponds with the oxide R2O3, where X = Cl, NO3, &c., or X2= SO4, CO3, &c.; but there also exist basic salts of the same type where X = HO or X2= O. We saw salts of all kinds among the salts of aluminium, chromium, and others. With uranic oxide no salts are known of the types UX6(UCl6, U(SO4)3, alums, &c., are not known), nor even salts, U(HO)2X4or UOX4, but it always forms salts of the type U(HO)4X2, or UO2X2. Judging from the fact that nearly all the salts of uranic oxide retain water in crystallising from their solutions, and that this water is difficult to separate from them, it may be thought to be water of hydration. This is seen in part from the fact that the composition of many of the salts of uranic oxide may then be expressed without the presence of water of crystallisation; for instance, U(HO)4K2Cl4(and the salt of NH4, U(HO)4K2(SO4)2, U(HO)4(C2H3O2)2. Sodium uranyl acetate however does not contain water.
[11]Uranyl nitrate, or uranium nitrate, UO2(NO3)2,6H2O, crystallises from its solutions in transparent yellowish-green prisms (from an acid solution), or in tabular crystals (from a neutral solution), which effloresce in the air and are easily soluble in water, alcohol, and ether, have a sp. gr. of 2·8, and fuse when heated, losing nitric acid and water in the process. If the salt itself (Berzelius) or its alcoholic solution (Malaguti) be heated up to the temperature at which oxides of nitrogen are evolved, there then remains a mass which, after being evaporated with water, leaves uranyl hydroxide, UO2(HO)2(sp. gr. 5·93), whilst if the salt be ignited there remains the dioxide, UO2, as a brick-red powder, which on further heating loses oxygen and forms the dark olive uranoso-uranic oxide, U3O8. The solution of the nitrate obtained from the ore is purified in the following manner: sulphurous anhydride is first passed through it in order to reduce the arsenic acid present into arsenious acid; the solution is then heated to 60°, and sulphuretted hydrogen passed through it; this precipitates the lead, arsenic, and tin, and certain other metals, as sulphides, insoluble in water and dilute nitric acid. This liquid is then filtered and evaporated with nitric acid to crystallisation, and the crystals are dissolved in ether. Or else the solution is first treated with chlorine in order to convert the ferrous chloride (produced by the action of the hydrogen sulphide) into ferric chloride, the oxides are then precipitated by ammonia, and the resultant precipitate, containing the oxides Fe2O3, UO3, and compounds of the latter with potash, lime, ammonia, and other bases present in the solution (the latter being due to the property of uranic oxide of combining with bases), is washed and dissolved in a strong, slightly-heated solution of ammonium carbonate, which dissolves the uranic oxide but not the ferric oxide. The solution is filtered, and on cooling deposits a well-crystallisinguranyl ammonium carbonate, UO2(NH4)4(CO3)3, in brilliant monoclinic crystals which on exposure to air slowly give off water, carbonic anhydride, and ammonia; the same decomposition is readily effected at 300°, the residue then consisting of uranic oxide. This salt is not very soluble in water, but is readily so in ammonium carbonate; it is obvious that it may readily be converted into all the other salts of oxides of uranium. Uranium salts are also purified in the form ofacetate, which is very sparingly soluble, and is therefore directly precipitated from a strong solution of the nitrate by mixing it with acetic acid.We may also mention theuranyl phosphate, HUPO6, which must be regarded as an orthophosphate in which two hydrogens are replaced by the radicle uranyl, UO2,i.e.as H(UO2)PO4. This salt is formed as a hydrated gelatinous yellow precipitate, on mixing a solution of uranyl nitrate with disodium phosphate. The precipitation occurs in the presence of acetic acid, but not in the presence of hydrochloric acid. If moreover an excess of an ammonium salt be present, the ammonia enters into the composition of the bright yellow gelatinous precipitate formed, in the proportion UO2NH4PO4. This precipitate is not soluble in water and acetic acid, and its solution in inorganic acids when boiled entirely expels all the phosphoric acid. This fact is taken advantage of for removing phosphoric acids from solutions—for instance, from those containing salts of calcium and magnesium.
[11]Uranyl nitrate, or uranium nitrate, UO2(NO3)2,6H2O, crystallises from its solutions in transparent yellowish-green prisms (from an acid solution), or in tabular crystals (from a neutral solution), which effloresce in the air and are easily soluble in water, alcohol, and ether, have a sp. gr. of 2·8, and fuse when heated, losing nitric acid and water in the process. If the salt itself (Berzelius) or its alcoholic solution (Malaguti) be heated up to the temperature at which oxides of nitrogen are evolved, there then remains a mass which, after being evaporated with water, leaves uranyl hydroxide, UO2(HO)2(sp. gr. 5·93), whilst if the salt be ignited there remains the dioxide, UO2, as a brick-red powder, which on further heating loses oxygen and forms the dark olive uranoso-uranic oxide, U3O8. The solution of the nitrate obtained from the ore is purified in the following manner: sulphurous anhydride is first passed through it in order to reduce the arsenic acid present into arsenious acid; the solution is then heated to 60°, and sulphuretted hydrogen passed through it; this precipitates the lead, arsenic, and tin, and certain other metals, as sulphides, insoluble in water and dilute nitric acid. This liquid is then filtered and evaporated with nitric acid to crystallisation, and the crystals are dissolved in ether. Or else the solution is first treated with chlorine in order to convert the ferrous chloride (produced by the action of the hydrogen sulphide) into ferric chloride, the oxides are then precipitated by ammonia, and the resultant precipitate, containing the oxides Fe2O3, UO3, and compounds of the latter with potash, lime, ammonia, and other bases present in the solution (the latter being due to the property of uranic oxide of combining with bases), is washed and dissolved in a strong, slightly-heated solution of ammonium carbonate, which dissolves the uranic oxide but not the ferric oxide. The solution is filtered, and on cooling deposits a well-crystallisinguranyl ammonium carbonate, UO2(NH4)4(CO3)3, in brilliant monoclinic crystals which on exposure to air slowly give off water, carbonic anhydride, and ammonia; the same decomposition is readily effected at 300°, the residue then consisting of uranic oxide. This salt is not very soluble in water, but is readily so in ammonium carbonate; it is obvious that it may readily be converted into all the other salts of oxides of uranium. Uranium salts are also purified in the form ofacetate, which is very sparingly soluble, and is therefore directly precipitated from a strong solution of the nitrate by mixing it with acetic acid.
We may also mention theuranyl phosphate, HUPO6, which must be regarded as an orthophosphate in which two hydrogens are replaced by the radicle uranyl, UO2,i.e.as H(UO2)PO4. This salt is formed as a hydrated gelatinous yellow precipitate, on mixing a solution of uranyl nitrate with disodium phosphate. The precipitation occurs in the presence of acetic acid, but not in the presence of hydrochloric acid. If moreover an excess of an ammonium salt be present, the ammonia enters into the composition of the bright yellow gelatinous precipitate formed, in the proportion UO2NH4PO4. This precipitate is not soluble in water and acetic acid, and its solution in inorganic acids when boiled entirely expels all the phosphoric acid. This fact is taken advantage of for removing phosphoric acids from solutions—for instance, from those containing salts of calcium and magnesium.
[12]Uranium dioxide, oruranyl, UO2, which is contained in the salts UO2X2, has the appearance and many of the properties of a metal. Uranic oxide may be regarded as uranyl oxide, (UO2)O, its salts as salts of this uranyl; its hydroxide, (UO2)H2O2, is constituted like CaH2O2. The green oxide of uranium, uranoso-uranic oxide (easily formed from uranic salts by the loss of oxygen), U3O8= UO2,2UO3, when ignited with charcoal or hydrogen (dry) gives a brilliant crystalline substance of sp. gr. about 11·0 (Urlaub), whose appearance resembles that of metals, and decomposes steam at a red heat with the evolution of hydrogen; it does not, however, decompose hydrochloric or sulphuric acid, but is oxidised by nitric acid. The same substance (i.e. uranium dioxide UO2) is also obtained by igniting the compound (UO2)K2Cl4in a stream of hydrogen, according to the equation UO2K4Cl4+ H2= UO2+ 2HCl + 2KCl. It was at first regarded as the metal. In 1841 Peligot found that it contained oxygen, because carbonic oxide and anhydride were evolved when it was ignited with charcoal in a stream of chlorine, and from 272 parts of the substance which was considered to be metal he obtained 382 parts of a volatile product containing 142 parts of chlorine. From this it was concluded that the substance taken contained an equivalent amount of oxygen. As 142 parts of chlorine correspond with 32 parts of oxygen, it followed that 272 - 32 = 240 parts of metal were combined in the substance with 32 parts of oxygen, and also in the chlorine compound obtained with 142 parts of chlorine. These calculations have been made for the now accepted atomic weight of uranium (U = 240,seeNote14). Peligot took another atomic weight, but this does not alter the principle of the argument.
[12]Uranium dioxide, oruranyl, UO2, which is contained in the salts UO2X2, has the appearance and many of the properties of a metal. Uranic oxide may be regarded as uranyl oxide, (UO2)O, its salts as salts of this uranyl; its hydroxide, (UO2)H2O2, is constituted like CaH2O2. The green oxide of uranium, uranoso-uranic oxide (easily formed from uranic salts by the loss of oxygen), U3O8= UO2,2UO3, when ignited with charcoal or hydrogen (dry) gives a brilliant crystalline substance of sp. gr. about 11·0 (Urlaub), whose appearance resembles that of metals, and decomposes steam at a red heat with the evolution of hydrogen; it does not, however, decompose hydrochloric or sulphuric acid, but is oxidised by nitric acid. The same substance (i.e. uranium dioxide UO2) is also obtained by igniting the compound (UO2)K2Cl4in a stream of hydrogen, according to the equation UO2K4Cl4+ H2= UO2+ 2HCl + 2KCl. It was at first regarded as the metal. In 1841 Peligot found that it contained oxygen, because carbonic oxide and anhydride were evolved when it was ignited with charcoal in a stream of chlorine, and from 272 parts of the substance which was considered to be metal he obtained 382 parts of a volatile product containing 142 parts of chlorine. From this it was concluded that the substance taken contained an equivalent amount of oxygen. As 142 parts of chlorine correspond with 32 parts of oxygen, it followed that 272 - 32 = 240 parts of metal were combined in the substance with 32 parts of oxygen, and also in the chlorine compound obtained with 142 parts of chlorine. These calculations have been made for the now accepted atomic weight of uranium (U = 240,seeNote14). Peligot took another atomic weight, but this does not alter the principle of the argument.
[13]Uranium tetrachloride, uranous chloride, UCl4, corresponds with uranous oxide as a base. It was obtained by Peligot by igniting uranic oxide mixed with charcoal in a stream ofdrychlorine: UO3+ 3C + 2Cl2= UCl4+ 3CO. This green volatile compound (Note12) crystallises in regular octahedra, is very hygroscopic, easily soluble in water, with the development of a considerable amount of heat, and no longer separates out from its solution in an anhydrous state, but disengages hydrochloric acid when evaporated. The solution of uranous chloride in water is green. It is also formed by the action of zinc and copper (forming cuprous chloride) on a solution of uranyl chloride, UO2Cl2, especially in the presence of hydrochloric acid and sal-ammoniac. Solutions of uranyl salts are converted into uranous salts by the action of various reducing agents, and among others by organic substances or by the action of light, whilst the salts UX4are converted into uranyl salts, UO2X2, by exposure to air or by oxidising agents. Solutions of the green uranyl salts act as powerful reducing agents, and give a brown precipitate of the uranous hydroxide, UH4O4, with potash and other alkalis. This hydroxide is easily soluble in acids but not in alkalis. On ignition it does not form the oxide UO2, because it decomposes water, but when the higher oxides of uranium are ignited in a stream of hydrogen or with charcoal they yield uranous oxide. Both it and the chloride UCl4, dissolve in strong sulphuric acid, forming a green salt, U(SO4)2,2H2O. The same salt, together with uranyl sulphate, UO2(SO4), is formed when the green oxide, U3O8, is dissolved in hot sulphuric acid. The salts obtained in the latter instance may be separated by adding alcohol to the solution, which is left exposed to the light; the alcohol reduces the uranyl salt to uranous salt, an excess of acid being required. An excess of water decomposes this salt, forming a basic salt, which is also easily produced under other circumstances, and contains UO(SO4),2H2O (which corresponds to the uranic salt).
[13]Uranium tetrachloride, uranous chloride, UCl4, corresponds with uranous oxide as a base. It was obtained by Peligot by igniting uranic oxide mixed with charcoal in a stream ofdrychlorine: UO3+ 3C + 2Cl2= UCl4+ 3CO. This green volatile compound (Note12) crystallises in regular octahedra, is very hygroscopic, easily soluble in water, with the development of a considerable amount of heat, and no longer separates out from its solution in an anhydrous state, but disengages hydrochloric acid when evaporated. The solution of uranous chloride in water is green. It is also formed by the action of zinc and copper (forming cuprous chloride) on a solution of uranyl chloride, UO2Cl2, especially in the presence of hydrochloric acid and sal-ammoniac. Solutions of uranyl salts are converted into uranous salts by the action of various reducing agents, and among others by organic substances or by the action of light, whilst the salts UX4are converted into uranyl salts, UO2X2, by exposure to air or by oxidising agents. Solutions of the green uranyl salts act as powerful reducing agents, and give a brown precipitate of the uranous hydroxide, UH4O4, with potash and other alkalis. This hydroxide is easily soluble in acids but not in alkalis. On ignition it does not form the oxide UO2, because it decomposes water, but when the higher oxides of uranium are ignited in a stream of hydrogen or with charcoal they yield uranous oxide. Both it and the chloride UCl4, dissolve in strong sulphuric acid, forming a green salt, U(SO4)2,2H2O. The same salt, together with uranyl sulphate, UO2(SO4), is formed when the green oxide, U3O8, is dissolved in hot sulphuric acid. The salts obtained in the latter instance may be separated by adding alcohol to the solution, which is left exposed to the light; the alcohol reduces the uranyl salt to uranous salt, an excess of acid being required. An excess of water decomposes this salt, forming a basic salt, which is also easily produced under other circumstances, and contains UO(SO4),2H2O (which corresponds to the uranic salt).
[14]The atomic weight of uranium was formerly taken as half the present one, U = 120, and the oxides U2O3, suboxide UO, and green oxide U3O4, were of the same types as the oxides of iron. With a certain resemblance to the elements of the iron group, uranium presents many points of distinction which do not permit its being grouped with them. Thus uranium forms a very stable oxide, U2O3(U = 120), but does not give the corresponding chloride U2Cl6(Roscoe, however, in 1874 obtained UCl5, like MoCl5and WCl5), and under those circumstances (the ignition of oxide of uranium mixed with charcoal, in a stream of chlorine), when the formation of this compound might be expected, it gives (U = 120) the chloride UCl2, which is characterised by its volatility; this is not a property, to such an extent, of any of the bichlorides, RCl2, of the iron group.The alteration or doubling of the atomic weight of uranium—i.e.the recognition of U = 240—was made for the first time in the first (Russian) edition of this work (1871), and in my memoir of the same year in Liebig'sAnnalen, on the ground that with an atomic weight 120, uranium could not be placed in the periodic system. I think it will not be superfluous to add the following remarks on this subject: (1) In the other groups (K—Rb—Cs, Ca—Sr—Ba, Cl—Br—I) the acid character of the oxides decreases and their basic character increases with the rise of atomic weight, and therefore we should expect to find the same in the group Cr—Mo—W—U, and if CrO3, MoO3, WO3be the anhydrides of acids then we indeed find a decrease in their acid character, and therefore uranium trioxide, UO3, should be a very feeble anhydride, but its basic properties should also be very feeble. Uranic oxide does indeed show these properties, as was pointed out above (Note10). (2) Chromium and its analogues, besides the oxides RO3, also form lower grades of oxidation RO2, R2O3, and the same is seen in uranium; it forms UO3, UO2, U2O3and their compounds. (3) Molybdenum and tungsten, in being reduced from RO3, easily and frequently give an intermediate oxide of a blue colour, and uranium shows the same property; giving the so-called green oxide which, according to present views, must be regarded as U3O8= UO22UO3, analogous to Mo3O8. (4) The higher chlorides, RCl6, possible for the elements of this group, are either unstable (WCl6) or do not exist at all (Cr); but there is one single lower volatile compound, which is decomposed by water, and liable to further reduction into a non-volatile chlorine product and the metal. The same is observed in uranium, which forms an easily volatile chloride, UCl4, decomposed by water. (5) The high sp. gr. of uranium (18·6) is explained by its analogy to tungsten (sp. gr. 19·1). (6) For uranium, as for chromium and tungsten, yellow tints predominate in the form RO3, whilst the lower forms are green and blue. (7) Zimmermann (1881) determined the vapour densities of uranous bromide, UBr4, and chloride, UCl4(19·4 and 13·2), and they were found to correspond to the formulæ given above—that is, they confirmed the higher atomic weight U = 240. Roscoe, a great authority on the metals of this group, was the first to accept the proposed atomic weight of uranium, U = 240, which since Zimmermann's work has been generally recognised.
[14]The atomic weight of uranium was formerly taken as half the present one, U = 120, and the oxides U2O3, suboxide UO, and green oxide U3O4, were of the same types as the oxides of iron. With a certain resemblance to the elements of the iron group, uranium presents many points of distinction which do not permit its being grouped with them. Thus uranium forms a very stable oxide, U2O3(U = 120), but does not give the corresponding chloride U2Cl6(Roscoe, however, in 1874 obtained UCl5, like MoCl5and WCl5), and under those circumstances (the ignition of oxide of uranium mixed with charcoal, in a stream of chlorine), when the formation of this compound might be expected, it gives (U = 120) the chloride UCl2, which is characterised by its volatility; this is not a property, to such an extent, of any of the bichlorides, RCl2, of the iron group.
The alteration or doubling of the atomic weight of uranium—i.e.the recognition of U = 240—was made for the first time in the first (Russian) edition of this work (1871), and in my memoir of the same year in Liebig'sAnnalen, on the ground that with an atomic weight 120, uranium could not be placed in the periodic system. I think it will not be superfluous to add the following remarks on this subject: (1) In the other groups (K—Rb—Cs, Ca—Sr—Ba, Cl—Br—I) the acid character of the oxides decreases and their basic character increases with the rise of atomic weight, and therefore we should expect to find the same in the group Cr—Mo—W—U, and if CrO3, MoO3, WO3be the anhydrides of acids then we indeed find a decrease in their acid character, and therefore uranium trioxide, UO3, should be a very feeble anhydride, but its basic properties should also be very feeble. Uranic oxide does indeed show these properties, as was pointed out above (Note10). (2) Chromium and its analogues, besides the oxides RO3, also form lower grades of oxidation RO2, R2O3, and the same is seen in uranium; it forms UO3, UO2, U2O3and their compounds. (3) Molybdenum and tungsten, in being reduced from RO3, easily and frequently give an intermediate oxide of a blue colour, and uranium shows the same property; giving the so-called green oxide which, according to present views, must be regarded as U3O8= UO22UO3, analogous to Mo3O8. (4) The higher chlorides, RCl6, possible for the elements of this group, are either unstable (WCl6) or do not exist at all (Cr); but there is one single lower volatile compound, which is decomposed by water, and liable to further reduction into a non-volatile chlorine product and the metal. The same is observed in uranium, which forms an easily volatile chloride, UCl4, decomposed by water. (5) The high sp. gr. of uranium (18·6) is explained by its analogy to tungsten (sp. gr. 19·1). (6) For uranium, as for chromium and tungsten, yellow tints predominate in the form RO3, whilst the lower forms are green and blue. (7) Zimmermann (1881) determined the vapour densities of uranous bromide, UBr4, and chloride, UCl4(19·4 and 13·2), and they were found to correspond to the formulæ given above—that is, they confirmed the higher atomic weight U = 240. Roscoe, a great authority on the metals of this group, was the first to accept the proposed atomic weight of uranium, U = 240, which since Zimmermann's work has been generally recognised.
[15]Uranium glass, obtained by the addition of the yellow salt K2U2O7to glass, has a green yellow fluorescence, and is sometimes employed for ornaments; it absorbs the violet rays, like the other salts of uranic oxide—that is, it possesses an absorption spectrum in which the violet rays are absent. The index of refraction of the absorbed rays is altered, and they are given out again as greenish-yellow rays; hence, compounds of uranic acid, when placed in the violet portion of the spectrum, emit a greenish-yellow light, and this forms one of the best examples (another is found in a solution of quinine sulphate) of the phenomenon of fluorescence. The rays of light which pass through uranic compounds do not contain the rays which excite the phenomena of fluorescence and of chemical transformation, as the researches of Stokes prove.
[15]Uranium glass, obtained by the addition of the yellow salt K2U2O7to glass, has a green yellow fluorescence, and is sometimes employed for ornaments; it absorbs the violet rays, like the other salts of uranic oxide—that is, it possesses an absorption spectrum in which the violet rays are absent. The index of refraction of the absorbed rays is altered, and they are given out again as greenish-yellow rays; hence, compounds of uranic acid, when placed in the violet portion of the spectrum, emit a greenish-yellow light, and this forms one of the best examples (another is found in a solution of quinine sulphate) of the phenomenon of fluorescence. The rays of light which pass through uranic compounds do not contain the rays which excite the phenomena of fluorescence and of chemical transformation, as the researches of Stokes prove.
[16]The comparison of potassium permanganate with potassium perchlorate, or of potassium manganate with potassium sulphate, shows directly that many of the physical and chemical properties of substances do not depend on the nature of the elements, but on the atomic types in which they appear, on the kind of movements, or on the positions in which the atoms forming the molecule occur.
[16]The comparison of potassium permanganate with potassium perchlorate, or of potassium manganate with potassium sulphate, shows directly that many of the physical and chemical properties of substances do not depend on the nature of the elements, but on the atomic types in which they appear, on the kind of movements, or on the positions in which the atoms forming the molecule occur.
[17]If, however, we compare the spectra (Vol. I. p.565) of chlorine, bromine, and iodine with that of manganese, a certain resemblance or analogy is to be found connecting manganese both to iron and to chlorine, bromine, and iodine.
[17]If, however, we compare the spectra (Vol. I. p.565) of chlorine, bromine, and iodine with that of manganese, a certain resemblance or analogy is to be found connecting manganese both to iron and to chlorine, bromine, and iodine.
[18]The name ‘peroxide’ should only be retained for thosehighestoxides (and MnO2stands between MnO and MnO3) which either by a direct method of double decomposition are able to give hydrogen peroxide or contain a larger proportion of oxygen than the base or the acid, just as hydrogen peroxide contains more oxygen than water. Their type will be H2O2, and they are exemplified by barium peroxide, BaO2, and sulphur peroxide, S2O7, &c. Such a dioxide as MnO2is, in all probability, a salt—that is, a manganous manganate, MnO3MnO, and also, as a basic salt of a feeble base, capable of combining with alkalis and acids. Hence the name of manganese peroxide should be abandoned, and replaced by manganese dioxide. PbO2is better termed lead dioxide than peroxide. Bisulphide of manganese, MnS2, corresponding to iron pyrites, FeS2, sometimes occurs in nature in fine octahedra (and cube combinations), for instance, in Sicily; it is called Hauerite.
[18]The name ‘peroxide’ should only be retained for thosehighestoxides (and MnO2stands between MnO and MnO3) which either by a direct method of double decomposition are able to give hydrogen peroxide or contain a larger proportion of oxygen than the base or the acid, just as hydrogen peroxide contains more oxygen than water. Their type will be H2O2, and they are exemplified by barium peroxide, BaO2, and sulphur peroxide, S2O7, &c. Such a dioxide as MnO2is, in all probability, a salt—that is, a manganous manganate, MnO3MnO, and also, as a basic salt of a feeble base, capable of combining with alkalis and acids. Hence the name of manganese peroxide should be abandoned, and replaced by manganese dioxide. PbO2is better termed lead dioxide than peroxide. Bisulphide of manganese, MnS2, corresponding to iron pyrites, FeS2, sometimes occurs in nature in fine octahedra (and cube combinations), for instance, in Sicily; it is called Hauerite.
[18 bis]On comparing the manganates with the permanganates—for example, K2MnO4with KMnO4—we find that they differ in composition by the abstraction of one equivalent of the metal. Such a relation in composition produced by oxidation is of frequent occurrence—for instance, K4Fe(CN)6in oxidising gives K3Fe(CN)6; H2SO4in oxidising gives persulphuric acid, HSO4, or H2S7O8; H2O forms HO or H2O2, &c.
[18 bis]On comparing the manganates with the permanganates—for example, K2MnO4with KMnO4—we find that they differ in composition by the abstraction of one equivalent of the metal. Such a relation in composition produced by oxidation is of frequent occurrence—for instance, K4Fe(CN)6in oxidising gives K3Fe(CN)6; H2SO4in oxidising gives persulphuric acid, HSO4, or H2S7O8; H2O forms HO or H2O2, &c.
[19]In the preparation of oxygen from the dioxide by means of H2SO4, MnSO4is formed; in the preparation of chlorine from HCl and MnO2, MnCl2is obtained. These two manganous salts may be taken as examples of compounds MnX2. Manganous sulphate generally contains various impurities, and also a large amount of iron salt (from the native MnO2), from which it cannot be freed by crystallisation. Their removal may, however, be effected by mixing a portion of the liquid with a solution of sodium carbonate; a precipitate of manganous carbonate is then formed. This precipitate is collected and washed, and then added to the remaining mass of the impure solution of manganous sulphate; on heating the solution with this precipitate, the whole of the iron is precipitated as oxide. This is due to the fact that in the solution of the manganese dioxide in sulphuric acid the whole of the iron is converted into the ferric state (because the dioxide acts as an oxidising agent), which, as an exceedingly feeble base precipitated by calcium carbonate and other kindred salts, is also precipitated by manganous carbonate. After being treated in this manner, the solution of manganous sulphate is further purified by crystallisation. If it be a bright red colour, it is due to the presence of higher grades of oxidation of manganese; they may be destroyed by boiling the solution, when the oxygen from the oxides of manganese is evolved and a very faintly coloured solution of manganous sulphate is obtained. This salt is remarkable for the facility with which it gives various combinations with water. By evaporating the almost colourless solution ofmanganous sulphateat very low temperatures, and by cooling the saturated solution at about 0°, crystals are obtained containing 7 atoms of water of crystallisation, MnSO4,7H2O, which are isomorphous with cobaltous and ferrous sulphates. These crystals, even at 10°, lose 5 p.c. of water, and completely effloresce at 15°, losing about 20 p.c. of water. By evaporating a solution of the salt at the ordinary temperature, but not above 20°, crystals are obtained containing 5 mol. H2O, which are isomorphous with copper sulphate; whilst if the crystallisation be carried on between 20° and 30°, large transparent prismatic crystals are formed containing 4 mol. H2O (see Nickel). A boiling solution also deposits these crystals together with crystals containing 3 mol. H2O, whilst the first salt, when fused and boiled with alcohol, gives crystals containing 2 mol. H2O. Graham obtained a monohydrated salt by drying the salt at about 200°. The last atom of water is eliminated with difficulty, as is the case with all salts like MnSO4nH2O. The crystals containing a considerable amount of water are rose-coloured, and the anhydrous crystals are colourless. The solubility of MnSO4,4H2O (Chapter I., Note24) per 100 parts of water is: at 10°, 127 parts; at 37°·5, 149 parts; at 75°, 145 parts; and at 101°, 92 parts. Whence it is seen that at the boiling-point this salt is less soluble than at lower temperatures, and therefore a solution saturated at the ordinary temperature becomes turbid when boiled. Manganous sulphate, being analogous to magnesium sulphate, is decomposed, like the latter, when ignited, but it does not then leave manganous oxide, but the intermediate oxide, Mn3O4. It gives double salts with the alkali sulphates. With aluminium sulphate it forms fine radiated crystals, whose composition resembles that of the alums—namely, MnAl2(SO4)4,24H2O. This salt is easily soluble in water, and occurs in nature.Manganous chloride, MCl2, crystallises with 4 mol. H2O, like the ferrous salt, and not with 6 mol. H2O like many kindred salts—for example, those of cobalt, calcium, and magnesium; 100 parts of water dissolve 38 parts of the anhydrous salt at 10° and 55 parts at 62°. Alcohol also dissolves manganous chloride, and the alcoholic solution burns with a red flame. This salt, like magnesium chloride, readily forms double salts. A solution of borax gives a dirty rose-coloured precipitate having the composition MnH4(BO3)2H2O, which is used as a drier in paint-making. Potassium cyanide produces a yellowish-grey precipitate, MnC2N2, with manganous salts, soluble in an excess of the reagent, a double salt, K4MnC6N6, corresponding with potassium ferrocyanide, being formed. On evaporation of this solution, a portion of the manganese is oxidised and precipitated, whilst a salt corresponding to Gmelin's red salt, K3,MnC6N6(seeChapterXXII.), remains in solution. Sulphuretted hydrogen does not precipitate salts of manganese, not even the acetate, but ammonium sulphide gives a flesh-coloured precipitate, MnS; at 320° this sulphide of manganese passes into a green variety (Antony). Oxalic acid in strong solutions of manganous salts gives a white precipitate of the oxalate, MnC2O4. This precipitate is insoluble in water, and is used for the preparation of manganous oxide itself because it decomposes like oxalic acid when ignited (in a tube without access of air), with the formation of carbonic anhydride, carbonic oxide, and manganous oxide.Manganous oxidethus obtained is a green powder, which however oxidises with such facility that it burns in air when brought into contact with an incandescent substance, and passes into the red intermediate oxide Mn3O4. In solutions of manganous salts, alkalis produce a precipitate of the hydroxide MnH2O2, which rapidly absorbs oxygen in the presence of air and gives the brown intermediate oxide, or, more correctly speaking, its hydrate.Manganous oxide, besides being obtained by the above-described method from manganous oxalate, may also be obtained by igniting the higher oxides in a stream of hydrogen, and also from manganese carbonate. The manganous oxide ignited in the presence of hydrogen acquires a great density, and is no longer so easily oxidised. It may also be obtained in a crystalline form, if during the ignition of the carbonate or higher oxide a trace of dry hydrochloric acid gas be passed into the current of hydrogen. It is thus obtained in the form of transparent emerald green crystals of the regular system, and in this state is easily soluble in acids.Manganous oxide in oxidising gives thered oxide of manganese, Mn5O4. This is the most stable of all the oxides of manganese; it is not only stable at the ordinary but also at a high temperature—that is, it does not absorb or disengage oxygen spontaneously. When ignited, all the higher oxides of manganese pass into it by losing oxygen, and manganous oxide by absorbing oxygen. This oxide does not give any distinct salts, but it dissolves in sulphuric acid, forming a dark red solution, which contains both manganous and manganic (of theoxide, Mn2O3) sulphates. The latter with potassium sulphate gives a manganese alum, in which the alumina is replaced by its isomorphous oxide of manganese. But this alum, like the solution of the intermediate oxide in sulphuric acid, evolves oxygen and leaves a manganous salt when slightly heated.Manganese dioxideis still less basic than the oxide, and disengages oxygen or a halogen in the presence of acids, forming manganous salts, like the oxide. However, if it be suspended in ether, and hydrochloric acid gas passed into the mixture, which is kept cool, the ether acquires a green colour, owing to the formation of tetrachloride of manganese, MnCl4, corresponding with the dioxide which passes into solution. It is however very unstable, being exceedingly easily decomposed with the evolution of chlorine. The corresponding fluoride, MnF4, obtained by Nicklés is much more stable. At all events, manganese dioxide does not exhibit any well-defined basic character, but has rather an acid character, which is particularly shown in the compounds MnF4and MnCl4just mentioned, and in the property of manganese dioxide of combining with alkalis. If the higher grades of oxidation of manganese be deoxidised in the presence of alkalis, they frequently give the dioxide combined with the alkali—for example, in the presence of potash a compound is formed which contains K2O,5MnO2, which shows the weak acid character of this oxide. When ignited in the presence of sodium compounds manganese dioxide frequently forms Na2O,8MnO2and Na2O,12MnO2, and lime when heated with MnO2gives from CaO,3MnO2to (CaO)2,MnO2(Rousseau) according to the temperature. Besides which, perhaps, MnO2is a saline compound, containing MnOMnO3or (MnO)3Mn2O7, and there are reactions which support such a view (Spring, Richards, Traube, and others); for instance it is known that manganous chloride and potassium permanganate give the dioxide in the presence of alkalis.Manganese dioxide may be obtained from manganous salts by the action of oxidising agents. If manganous hydroxide or carbonate be shaken up in water through which chlorine is passed, the hypochlorite of the metal is not formed, as is the case with certain other oxides, but manganese dioxide is precipitated: 2MnO2H2+ Cl2= MnCl2+ MnO2,H2O + H2O. Owing to this fact, hypochlorites in the presence of alkalis and acetic acid when added to a solution of manganous salts give hydrated manganese dioxide, as was mentioned above. Manganous nitrate also leaves manganese dioxide when heated to 200°. It is also obtained from manganous and manganic salts of the alkalis, when they are decomposed in the presence of a small amount of acid; the practical method of converting the salts MnX2into the higher grades of oxidation is given in Chapter II., Note6.
[19]In the preparation of oxygen from the dioxide by means of H2SO4, MnSO4is formed; in the preparation of chlorine from HCl and MnO2, MnCl2is obtained. These two manganous salts may be taken as examples of compounds MnX2. Manganous sulphate generally contains various impurities, and also a large amount of iron salt (from the native MnO2), from which it cannot be freed by crystallisation. Their removal may, however, be effected by mixing a portion of the liquid with a solution of sodium carbonate; a precipitate of manganous carbonate is then formed. This precipitate is collected and washed, and then added to the remaining mass of the impure solution of manganous sulphate; on heating the solution with this precipitate, the whole of the iron is precipitated as oxide. This is due to the fact that in the solution of the manganese dioxide in sulphuric acid the whole of the iron is converted into the ferric state (because the dioxide acts as an oxidising agent), which, as an exceedingly feeble base precipitated by calcium carbonate and other kindred salts, is also precipitated by manganous carbonate. After being treated in this manner, the solution of manganous sulphate is further purified by crystallisation. If it be a bright red colour, it is due to the presence of higher grades of oxidation of manganese; they may be destroyed by boiling the solution, when the oxygen from the oxides of manganese is evolved and a very faintly coloured solution of manganous sulphate is obtained. This salt is remarkable for the facility with which it gives various combinations with water. By evaporating the almost colourless solution ofmanganous sulphateat very low temperatures, and by cooling the saturated solution at about 0°, crystals are obtained containing 7 atoms of water of crystallisation, MnSO4,7H2O, which are isomorphous with cobaltous and ferrous sulphates. These crystals, even at 10°, lose 5 p.c. of water, and completely effloresce at 15°, losing about 20 p.c. of water. By evaporating a solution of the salt at the ordinary temperature, but not above 20°, crystals are obtained containing 5 mol. H2O, which are isomorphous with copper sulphate; whilst if the crystallisation be carried on between 20° and 30°, large transparent prismatic crystals are formed containing 4 mol. H2O (see Nickel). A boiling solution also deposits these crystals together with crystals containing 3 mol. H2O, whilst the first salt, when fused and boiled with alcohol, gives crystals containing 2 mol. H2O. Graham obtained a monohydrated salt by drying the salt at about 200°. The last atom of water is eliminated with difficulty, as is the case with all salts like MnSO4nH2O. The crystals containing a considerable amount of water are rose-coloured, and the anhydrous crystals are colourless. The solubility of MnSO4,4H2O (Chapter I., Note24) per 100 parts of water is: at 10°, 127 parts; at 37°·5, 149 parts; at 75°, 145 parts; and at 101°, 92 parts. Whence it is seen that at the boiling-point this salt is less soluble than at lower temperatures, and therefore a solution saturated at the ordinary temperature becomes turbid when boiled. Manganous sulphate, being analogous to magnesium sulphate, is decomposed, like the latter, when ignited, but it does not then leave manganous oxide, but the intermediate oxide, Mn3O4. It gives double salts with the alkali sulphates. With aluminium sulphate it forms fine radiated crystals, whose composition resembles that of the alums—namely, MnAl2(SO4)4,24H2O. This salt is easily soluble in water, and occurs in nature.
Manganous chloride, MCl2, crystallises with 4 mol. H2O, like the ferrous salt, and not with 6 mol. H2O like many kindred salts—for example, those of cobalt, calcium, and magnesium; 100 parts of water dissolve 38 parts of the anhydrous salt at 10° and 55 parts at 62°. Alcohol also dissolves manganous chloride, and the alcoholic solution burns with a red flame. This salt, like magnesium chloride, readily forms double salts. A solution of borax gives a dirty rose-coloured precipitate having the composition MnH4(BO3)2H2O, which is used as a drier in paint-making. Potassium cyanide produces a yellowish-grey precipitate, MnC2N2, with manganous salts, soluble in an excess of the reagent, a double salt, K4MnC6N6, corresponding with potassium ferrocyanide, being formed. On evaporation of this solution, a portion of the manganese is oxidised and precipitated, whilst a salt corresponding to Gmelin's red salt, K3,MnC6N6(seeChapterXXII.), remains in solution. Sulphuretted hydrogen does not precipitate salts of manganese, not even the acetate, but ammonium sulphide gives a flesh-coloured precipitate, MnS; at 320° this sulphide of manganese passes into a green variety (Antony). Oxalic acid in strong solutions of manganous salts gives a white precipitate of the oxalate, MnC2O4. This precipitate is insoluble in water, and is used for the preparation of manganous oxide itself because it decomposes like oxalic acid when ignited (in a tube without access of air), with the formation of carbonic anhydride, carbonic oxide, and manganous oxide.Manganous oxidethus obtained is a green powder, which however oxidises with such facility that it burns in air when brought into contact with an incandescent substance, and passes into the red intermediate oxide Mn3O4. In solutions of manganous salts, alkalis produce a precipitate of the hydroxide MnH2O2, which rapidly absorbs oxygen in the presence of air and gives the brown intermediate oxide, or, more correctly speaking, its hydrate.
Manganous oxide, besides being obtained by the above-described method from manganous oxalate, may also be obtained by igniting the higher oxides in a stream of hydrogen, and also from manganese carbonate. The manganous oxide ignited in the presence of hydrogen acquires a great density, and is no longer so easily oxidised. It may also be obtained in a crystalline form, if during the ignition of the carbonate or higher oxide a trace of dry hydrochloric acid gas be passed into the current of hydrogen. It is thus obtained in the form of transparent emerald green crystals of the regular system, and in this state is easily soluble in acids.
Manganous oxide in oxidising gives thered oxide of manganese, Mn5O4. This is the most stable of all the oxides of manganese; it is not only stable at the ordinary but also at a high temperature—that is, it does not absorb or disengage oxygen spontaneously. When ignited, all the higher oxides of manganese pass into it by losing oxygen, and manganous oxide by absorbing oxygen. This oxide does not give any distinct salts, but it dissolves in sulphuric acid, forming a dark red solution, which contains both manganous and manganic (of theoxide, Mn2O3) sulphates. The latter with potassium sulphate gives a manganese alum, in which the alumina is replaced by its isomorphous oxide of manganese. But this alum, like the solution of the intermediate oxide in sulphuric acid, evolves oxygen and leaves a manganous salt when slightly heated.
Manganese dioxideis still less basic than the oxide, and disengages oxygen or a halogen in the presence of acids, forming manganous salts, like the oxide. However, if it be suspended in ether, and hydrochloric acid gas passed into the mixture, which is kept cool, the ether acquires a green colour, owing to the formation of tetrachloride of manganese, MnCl4, corresponding with the dioxide which passes into solution. It is however very unstable, being exceedingly easily decomposed with the evolution of chlorine. The corresponding fluoride, MnF4, obtained by Nicklés is much more stable. At all events, manganese dioxide does not exhibit any well-defined basic character, but has rather an acid character, which is particularly shown in the compounds MnF4and MnCl4just mentioned, and in the property of manganese dioxide of combining with alkalis. If the higher grades of oxidation of manganese be deoxidised in the presence of alkalis, they frequently give the dioxide combined with the alkali—for example, in the presence of potash a compound is formed which contains K2O,5MnO2, which shows the weak acid character of this oxide. When ignited in the presence of sodium compounds manganese dioxide frequently forms Na2O,8MnO2and Na2O,12MnO2, and lime when heated with MnO2gives from CaO,3MnO2to (CaO)2,MnO2(Rousseau) according to the temperature. Besides which, perhaps, MnO2is a saline compound, containing MnOMnO3or (MnO)3Mn2O7, and there are reactions which support such a view (Spring, Richards, Traube, and others); for instance it is known that manganous chloride and potassium permanganate give the dioxide in the presence of alkalis.
Manganese dioxide may be obtained from manganous salts by the action of oxidising agents. If manganous hydroxide or carbonate be shaken up in water through which chlorine is passed, the hypochlorite of the metal is not formed, as is the case with certain other oxides, but manganese dioxide is precipitated: 2MnO2H2+ Cl2= MnCl2+ MnO2,H2O + H2O. Owing to this fact, hypochlorites in the presence of alkalis and acetic acid when added to a solution of manganous salts give hydrated manganese dioxide, as was mentioned above. Manganous nitrate also leaves manganese dioxide when heated to 200°. It is also obtained from manganous and manganic salts of the alkalis, when they are decomposed in the presence of a small amount of acid; the practical method of converting the salts MnX2into the higher grades of oxidation is given in Chapter II., Note6.
[20]Other chemists have obtained manganese by different methods, and attributed different properties to it. This difference probably depends on the presence of carbon in different proportions. Deville obtained manganese by subjecting the pure dioxide, mixed with pure charcoal (from burnt sugar), to a strong heat in a lime crucible until the resultant metal fused. The metal obtained had a rose tint, like bismuth, and like it was very brittle, although exceedingly hard. It decomposed water at the ordinary temperature. Brunner obtained manganese having a specific gravity of about 7·2, which decomposed water very feebly at the ordinary temperature, did not oxidise in air, and was capable of taking a bright polish, like steel; it had the grey colour of cast iron, was very brittle, and hard enough to scratch steel and glass, like a diamond. Brunner's method was as follows: He decomposed the manganese fluoride (obtained as a soluble compound by the action of hydrofluoric acid on manganese carbonate) with sodium, by mixing these substances together in a crucible and covering the mixture with a layer of salt and fluor spar; after which the crucible was first gradually heated until the reaction began, and then strongly heated in order to fuse the metal separated. Glatzel (1889) obtained 25 grms. of manganese, having a grey colour and sp. gr. 7·39, by heating a mixture of 100 grms. of MnCl2with 200 grms. KCl and 15 grms. Mg to a bright white heat. Moissan and others, by heating the oxides of manganese with carbon in the electric furnace, obtained carbides of manganese—for example, Mn3C—and remarked that the metal volatilised in the heat of the voltaic arc. Metallic manganese is, however, not prepared on a large scale, but only its alloys with carbon (they readily and rapidly oxidise) andferro-manganeseor a coarsely crystalline alloy of iron, manganese and carbon, which is smelted in blast-furnaces like pig-iron (seeChapterXXII.) This ferro-manganese is employed in the manufacture of steel by Bessemer's and other processes (see ChapterXXII.) and for the manufacture of manganese bronze. However, in America, Green and Wahl (1895) obtained almost pure metallic manganese on a large scale. They first treat the ore of MnO2with 30 p.c. sulphuric acid (which extracts all the oxides of iron present in the ore), and then heat it in a reducing flame to convert it into MnO, which they mix with a powder of Al, lime and CaF2(as a flux), and heat the mixture in a crucible lined with magnesia; a reaction immediately takes place at a certain temperature, and a metal of specific gravity 7·3 is obtained, which only contains a small trace of iron.Manganese gives two compounds withnitrogen, Mn5N2and Mn3N2. They were obtained by Prelinger (1894) from the amalgam of manganese Mn2Hg5(obtained on a mercury anode by the action of an electric current upon a solution of MnCl2); the mercury may be removed from this amalgam by heating it in an atmosphere of hydrogen, and then metallic manganese is obtained as a grey porous mass of specific gravity 7·42. If this amalgam be heated in dry nitrogen it gives Mn5N2(grey powder, sp. gr. 6·58), but if heated in an atmosphere of NH3it gives (as also does Mn5N2) Mn3N2, (a dark mass with a metallic lustre, sp. gr. 6·21), which, when heated in nitrogen is converted into Mn5N2, and if heated in hydrogen evolves NH3and disengages hydrogen from a solution of NH4Cl. At all events, manganese is a metal which decomposes water more easily than iron, nickel, and cobalt.
[20]Other chemists have obtained manganese by different methods, and attributed different properties to it. This difference probably depends on the presence of carbon in different proportions. Deville obtained manganese by subjecting the pure dioxide, mixed with pure charcoal (from burnt sugar), to a strong heat in a lime crucible until the resultant metal fused. The metal obtained had a rose tint, like bismuth, and like it was very brittle, although exceedingly hard. It decomposed water at the ordinary temperature. Brunner obtained manganese having a specific gravity of about 7·2, which decomposed water very feebly at the ordinary temperature, did not oxidise in air, and was capable of taking a bright polish, like steel; it had the grey colour of cast iron, was very brittle, and hard enough to scratch steel and glass, like a diamond. Brunner's method was as follows: He decomposed the manganese fluoride (obtained as a soluble compound by the action of hydrofluoric acid on manganese carbonate) with sodium, by mixing these substances together in a crucible and covering the mixture with a layer of salt and fluor spar; after which the crucible was first gradually heated until the reaction began, and then strongly heated in order to fuse the metal separated. Glatzel (1889) obtained 25 grms. of manganese, having a grey colour and sp. gr. 7·39, by heating a mixture of 100 grms. of MnCl2with 200 grms. KCl and 15 grms. Mg to a bright white heat. Moissan and others, by heating the oxides of manganese with carbon in the electric furnace, obtained carbides of manganese—for example, Mn3C—and remarked that the metal volatilised in the heat of the voltaic arc. Metallic manganese is, however, not prepared on a large scale, but only its alloys with carbon (they readily and rapidly oxidise) andferro-manganeseor a coarsely crystalline alloy of iron, manganese and carbon, which is smelted in blast-furnaces like pig-iron (seeChapterXXII.) This ferro-manganese is employed in the manufacture of steel by Bessemer's and other processes (see ChapterXXII.) and for the manufacture of manganese bronze. However, in America, Green and Wahl (1895) obtained almost pure metallic manganese on a large scale. They first treat the ore of MnO2with 30 p.c. sulphuric acid (which extracts all the oxides of iron present in the ore), and then heat it in a reducing flame to convert it into MnO, which they mix with a powder of Al, lime and CaF2(as a flux), and heat the mixture in a crucible lined with magnesia; a reaction immediately takes place at a certain temperature, and a metal of specific gravity 7·3 is obtained, which only contains a small trace of iron.
Manganese gives two compounds withnitrogen, Mn5N2and Mn3N2. They were obtained by Prelinger (1894) from the amalgam of manganese Mn2Hg5(obtained on a mercury anode by the action of an electric current upon a solution of MnCl2); the mercury may be removed from this amalgam by heating it in an atmosphere of hydrogen, and then metallic manganese is obtained as a grey porous mass of specific gravity 7·42. If this amalgam be heated in dry nitrogen it gives Mn5N2(grey powder, sp. gr. 6·58), but if heated in an atmosphere of NH3it gives (as also does Mn5N2) Mn3N2, (a dark mass with a metallic lustre, sp. gr. 6·21), which, when heated in nitrogen is converted into Mn5N2, and if heated in hydrogen evolves NH3and disengages hydrogen from a solution of NH4Cl. At all events, manganese is a metal which decomposes water more easily than iron, nickel, and cobalt.
[21]Volume I. p. 157, Note7.
[21]Volume I. p. 157, Note7.
[22]It was known to the alchemists by this name, but the true explanation of the change in colour is due to the researches of Chevillot, Edwards, Mitscherlich, and Forchhammer. The change in colour of potassium manganate is due to its instability and to its splitting up into two other manganese compounds, a higher and a lower: 3MnO3= Mn2O7+ MnO2. Manganese trioxide is really decomposed in this manner by the action of water (see later): 3MnO3+ H2O = 2MnHO4+ MnO2(Franke, Thorpe, and Humbly). The instability of the salt is proved by the fact of its being deoxidised by organic matter, with the formation of manganese dioxide and alkali, so that, for instance, a solution of this salt cannot be filtered through paper. The presence of an excess of alkali increases the stability of the salt; when heated it breaks up in the presence of water, with the evolution of oxygen.The method of preparingpotassium permanganatewill be understood from the above. There are many recipes for preparing this substance, as it is now used in considerable quantities both for technical and laboratory purposes. But in all cases the essence of the methods is one and the same: a mixture of alkali with any oxide of manganese (even manganous hydroxide, which may be obtained from manganous chloride) is first heated in the presence of air or of an oxidising substance (for the sake of rapidity, with potassium chlorate); the resultant mass is then treated with water and heated, when manganese dioxide is precipitated and potassium permanganate remains in solution. This solution may be boiled, as the liquid will contain free alkali; but the solution cannot be evaporated to dryness, because a strong solution, as well as the solid salt, is decomposed by heat.By adding a dilute solution of manganous sulphate to a boiling mixture of lead dioxide and dilute nitric acid, the whole of the manganese may be converted into permanganic acid (Crum).
[22]It was known to the alchemists by this name, but the true explanation of the change in colour is due to the researches of Chevillot, Edwards, Mitscherlich, and Forchhammer. The change in colour of potassium manganate is due to its instability and to its splitting up into two other manganese compounds, a higher and a lower: 3MnO3= Mn2O7+ MnO2. Manganese trioxide is really decomposed in this manner by the action of water (see later): 3MnO3+ H2O = 2MnHO4+ MnO2(Franke, Thorpe, and Humbly). The instability of the salt is proved by the fact of its being deoxidised by organic matter, with the formation of manganese dioxide and alkali, so that, for instance, a solution of this salt cannot be filtered through paper. The presence of an excess of alkali increases the stability of the salt; when heated it breaks up in the presence of water, with the evolution of oxygen.
The method of preparingpotassium permanganatewill be understood from the above. There are many recipes for preparing this substance, as it is now used in considerable quantities both for technical and laboratory purposes. But in all cases the essence of the methods is one and the same: a mixture of alkali with any oxide of manganese (even manganous hydroxide, which may be obtained from manganous chloride) is first heated in the presence of air or of an oxidising substance (for the sake of rapidity, with potassium chlorate); the resultant mass is then treated with water and heated, when manganese dioxide is precipitated and potassium permanganate remains in solution. This solution may be boiled, as the liquid will contain free alkali; but the solution cannot be evaporated to dryness, because a strong solution, as well as the solid salt, is decomposed by heat.
By adding a dilute solution of manganous sulphate to a boiling mixture of lead dioxide and dilute nitric acid, the whole of the manganese may be converted into permanganic acid (Crum).
[22 bis]The solution of this salt with an excess of impure commercial alkali generally acquires a green tint.
[22 bis]The solution of this salt with an excess of impure commercial alkali generally acquires a green tint.
[23]A solution of potassium permanganate gives a beautiful absorption spectrum (ChapterXIII.) If the light in passing through this solution loses a portion of its rays in it (if one may so account for it), this is partially explained by the increased oxidising power which the solution then acquires. We may here also remark that a dilute solution of permanganate of potassium forms a colourless solution with nickel salts, because the green colour of the solution of nickel salts is complementary to the red. Such a decolorised solution, containing a large proportion of nickel and a small proportion of manganese, decomposes after a time, throws down a precipitate, and re-acquires the green colour proper to the nickel salts. The addition of a solution of a cobalt salt (rose-red) to the nickel salt also destroys the colour of both salts.
[23]A solution of potassium permanganate gives a beautiful absorption spectrum (ChapterXIII.) If the light in passing through this solution loses a portion of its rays in it (if one may so account for it), this is partially explained by the increased oxidising power which the solution then acquires. We may here also remark that a dilute solution of permanganate of potassium forms a colourless solution with nickel salts, because the green colour of the solution of nickel salts is complementary to the red. Such a decolorised solution, containing a large proportion of nickel and a small proportion of manganese, decomposes after a time, throws down a precipitate, and re-acquires the green colour proper to the nickel salts. The addition of a solution of a cobalt salt (rose-red) to the nickel salt also destroys the colour of both salts.
[24]If sulphuric acid is allowed to act on potassium permanganate without any special precautions, a large amount of oxygen is evolved (it may even explode and inflame), and a violet spray of the decomposing permanganic acid is given off. But if the pure salt (i.e.free from chlorine) be dissolved in pure well-cooled sulphuric acid, without any rise in temperature, a green-coloured liquid settles at the bottom of the vessel. This liquid does not contain any sulphuric acid, and consists of permanganic anhydride, Mn2O7(Aschoff, Terreil). It is impossible to prepare any considerable quantity of the anhydride by this method, as it decomposes with an explosion as it collects, evolving oxygen and leaving red oxide of manganese.Permanganic anhydride, Mn2O7, in dissolving in sulphuric acid, gives a green solution, which (according to Franke, 1887) contains a compound Mn2SO10= (MnO3)2SO4—that is, sulphuric acid in which both hydrogens are replaced by the group MnO3, which is combined with OK in permanganate of potassium. This mixture with a small quantity of water gives Mn2O7, according to the equation: (MnO3)2SO4+ H2O = H2SO4+ Mn2O7, and when heated to 30° it givesmanganese trioxide, (MnO3)2SO4+ H2O = 2MnO2+ H2SO4+ O. Pure manganese trioxide is obtained if the solution of (MnO3)2SO4be poured in drops on to sodium carbonate. Then, together with carbonic anhydride, a spray of manganese trioxide passes over, which may be collected in a well-cooled receiver, and this shows that the reaction proceeds according to the equation: (MnO3)2SO4+ Na2CO3= Na2SO4+ 2MnO3+ CO2+ O (Thorpe). The trioxide is decomposed by water, forming manganese dioxide and a solution ofpermanganic acid: 3MnO3+ H2O = MnO2+ 2HMnO4. The same acid is obtained by dissolving permanganic anhydride in water.Barium permanganate when treated with sulphuric acid gives the same acid. This barium salt may be prepared by the action of barium chloride on the difficultly soluble silver permanganate, AgMnO4, which is precipitated on mixing a strong solution of the potassium salt with silver nitrate. The solution of permanganic acid forms a bright red liquid which reflects a dark violet tint. A dilute solution has exactly the same colour as that of the potassium salt. It deposits manganese dioxide when exposed to the action of light, and also when heated above 60°, and this proceeds the more rapidly the more dilute the solution. It shows its oxidising properties in many cases, as already mentioned. Even hydrogen gas is absorbed by a solution of permanganic acid; and charcoal and sulphur are also oxidised by it, as they are by potassium permanganate. This may be taken advantage of in analysing gunpowder, because when it is treated with a solution of potassium permanganate, all the sulphur is converted into sulphuric acid and all the charcoal into carbonic anhydride. Finely-divided platinum immediately decomposes permanganic acid. With potassium iodide it liberates iodine (which may afterwards be oxidised into iodic acid) (Mitscherlich, Fromherz, Aschoff, and others). Ammonia does not form a corresponding salt with free permanganic acid, because it is oxidised with evolution of nitrogen. The oxidising action of permanganic acid in a strong solution may be accompanied by flame and the formation of violet fumes of permanganic acid; thus a strong solution of it takes fire when brought into contact with paper, alcohol, alkaline sulphides, fats, &c.We may add that, according to Franke, 1 part of potassium permanganate with 13 parts of sulphuric acid at 100° gives brown crystals of the salt Mn2(SO4)3,H2SO4,4H2O, which gives a precipitate of hydrated manganese dioxide, H2MnO3= MnO2H2O, when treated with water.Spring, by precipitating potassium permanganate with sodium sulphite and washing the precipitate by decantation, obtained a soluble colloidal manganese oxide, whose composition was the mean between Mn2O3and MnO2—namely, Mn2O3,4(MnO2H2O).
[24]If sulphuric acid is allowed to act on potassium permanganate without any special precautions, a large amount of oxygen is evolved (it may even explode and inflame), and a violet spray of the decomposing permanganic acid is given off. But if the pure salt (i.e.free from chlorine) be dissolved in pure well-cooled sulphuric acid, without any rise in temperature, a green-coloured liquid settles at the bottom of the vessel. This liquid does not contain any sulphuric acid, and consists of permanganic anhydride, Mn2O7(Aschoff, Terreil). It is impossible to prepare any considerable quantity of the anhydride by this method, as it decomposes with an explosion as it collects, evolving oxygen and leaving red oxide of manganese.Permanganic anhydride, Mn2O7, in dissolving in sulphuric acid, gives a green solution, which (according to Franke, 1887) contains a compound Mn2SO10= (MnO3)2SO4—that is, sulphuric acid in which both hydrogens are replaced by the group MnO3, which is combined with OK in permanganate of potassium. This mixture with a small quantity of water gives Mn2O7, according to the equation: (MnO3)2SO4+ H2O = H2SO4+ Mn2O7, and when heated to 30° it givesmanganese trioxide, (MnO3)2SO4+ H2O = 2MnO2+ H2SO4+ O. Pure manganese trioxide is obtained if the solution of (MnO3)2SO4be poured in drops on to sodium carbonate. Then, together with carbonic anhydride, a spray of manganese trioxide passes over, which may be collected in a well-cooled receiver, and this shows that the reaction proceeds according to the equation: (MnO3)2SO4+ Na2CO3= Na2SO4+ 2MnO3+ CO2+ O (Thorpe). The trioxide is decomposed by water, forming manganese dioxide and a solution ofpermanganic acid: 3MnO3+ H2O = MnO2+ 2HMnO4. The same acid is obtained by dissolving permanganic anhydride in water.
Barium permanganate when treated with sulphuric acid gives the same acid. This barium salt may be prepared by the action of barium chloride on the difficultly soluble silver permanganate, AgMnO4, which is precipitated on mixing a strong solution of the potassium salt with silver nitrate. The solution of permanganic acid forms a bright red liquid which reflects a dark violet tint. A dilute solution has exactly the same colour as that of the potassium salt. It deposits manganese dioxide when exposed to the action of light, and also when heated above 60°, and this proceeds the more rapidly the more dilute the solution. It shows its oxidising properties in many cases, as already mentioned. Even hydrogen gas is absorbed by a solution of permanganic acid; and charcoal and sulphur are also oxidised by it, as they are by potassium permanganate. This may be taken advantage of in analysing gunpowder, because when it is treated with a solution of potassium permanganate, all the sulphur is converted into sulphuric acid and all the charcoal into carbonic anhydride. Finely-divided platinum immediately decomposes permanganic acid. With potassium iodide it liberates iodine (which may afterwards be oxidised into iodic acid) (Mitscherlich, Fromherz, Aschoff, and others). Ammonia does not form a corresponding salt with free permanganic acid, because it is oxidised with evolution of nitrogen. The oxidising action of permanganic acid in a strong solution may be accompanied by flame and the formation of violet fumes of permanganic acid; thus a strong solution of it takes fire when brought into contact with paper, alcohol, alkaline sulphides, fats, &c.
We may add that, according to Franke, 1 part of potassium permanganate with 13 parts of sulphuric acid at 100° gives brown crystals of the salt Mn2(SO4)3,H2SO4,4H2O, which gives a precipitate of hydrated manganese dioxide, H2MnO3= MnO2H2O, when treated with water.
Spring, by precipitating potassium permanganate with sodium sulphite and washing the precipitate by decantation, obtained a soluble colloidal manganese oxide, whose composition was the mean between Mn2O3and MnO2—namely, Mn2O3,4(MnO2H2O).