Sulphur, selenium, and tellurium belong to the uneven series of the sixth group. In the even series of this group there are knownchromium, molybdenum, tungsten, and uranium; these give acid oxides of the type RO3, like SO3. Their acid properties are less sharply defined than those of sulphur, selenium, and tellurium, as is the case with all elements of the even series as compared with those of the uneven series in the same group. But still the oxides CrO3, MoO3, WO3, and even UO3, have clearly defined acid properties, and form salts of the composition MO,nRO3with bases MO. In the case of the heavy elements, and especially of uranium, the type of oxide, UO3, is less acid and more basic, because in the even series of oxides the element with the highest atomic weight always acquires a more and more pronounced basic character. Hence UO3shows the properties of a base, and gives salts UO2X2. The basic properties of chromium, molybdenum, tungsten, and uranium are most clearly expressed in the lower oxides, which they all form. Thus chromic oxide, Cr2O3, is as distinct a base as alumina, Al2O3.
Of all these elementschromiumis the most widely distributed and the most frequently used. It gives chromic anhydride, CrO3, and chromic oxide, Cr2O3—two compounds whose relative amounts of oxygen stand in the ratio 2 : 1. Chromium is, although somewhat rarely, met with in nature as a compound of one or the other type. The red chromium ore of the Urals, lead chromate or crocoisite PbCrO4, was the source in which chromium was discovered by Vauquelin, who gave it this name (from the Greek word signifying colour) owing to the brilliant colours of its compounds; the chromates (salts of chromic anhydride) are red and yellow, and the chromic salts (from Cr2O3) green and violet. The red lead chromate is, however, a rare chromium ore found only in the Urals and in a few other localities. Chromic oxide, Cr2O3, is more frequently met with. In small quantities it forms the colouring matter of many minerals and rocks—for example,of some serpentines. The commonest ore, and the chief source of the chromium compounds, is thechrome iron oreor chromite, which occurs in the Urals[1]and Asia Minor, California, Australia, and other localities. This is magnetic iron ore, FeO,Fe2O3, in which the ferric oxide is replaced by chromic oxide, its composition being FeO,Cr2O3. Chrome iron ore crystallises in octahedra of sp. gr. 4·4; it has a feeble metallic lustre, is of a greyish-black colour, and gives a brown powder. It is very feebly acted on by acids, but when fused with potassium acid sulphate it gives a soluble mass, which contains a chromic salt, besides potassium sulphate and ferrous sulphate. In practice the treatment of chrome iron ore is mainly carried on for the preparation of chromates, and not of chromic salts, and therefore we will trace the history of the element by beginning with chromic acid, and especially with the working up of the chrome iron ore intopotassium dichromate, K2Cr2O7, as the most common salt of this acid. It must be remarked that chromic anhydride, CrO3, is only obtained in an anhydrous state, and is distinguished for its capacity for easily giving anhydro-salts with the alkalis, containing one, two, and even three equivalents of the anhydride to one equivalent of base. Thus among the potassium salts there is known the normal or yellow chromate, K2CrO4, which corresponds to, and is perfectly isomorphous with, potassium sulphate, easily forms isomorphous mixtures with it, and is not therefore suitable for a process in which it is necessary to separate the salt from a mixture containing sulphates. As in the presence of a certain excess of acid, the dichromate, K2Cr2O7= 2K2CrO4+ 2HX - 2KX - H2O, is easily formed from K2CrO4, the object of the manufacturer is to produce such a dichromate, the more so as it contains a larger proportion of the elements of chromic acid than the normal salt. Finely-ground chrome iron ore, when heated with an alkali, absorbs oxygen almost as easily (Chapter III., Note7) as a mixture of the oxides of manganese with an alkali. This absorption is due to the presence of chromic oxide, which is oxidised into the anhydride, and then combines with the alkali Cr2O3+ O3= 2CrO3. As the oxidation and formation of the chromate proceeds, the mass turnsyellow. The iron is also oxidised, but does not give ferric acid, because the capacity of the chromium for oxidation is incomparably greater than that of the iron.
A mixture of lime (sometimes with potash) and chrome iron ore is heated in a reverberatory furnace, with free access of air and at ared heat for several hours, until the mass becomes yellow; it then contains normal calcium chromate, CaCr_O4, which is insoluble in water in the presence of an excess of lime.[1 bis]The resultant mass is ground up, and treated with water and sulphuric acid. The excess of lime forms gypsum, and the soluble calcium dichromate, CaCr2O7, together with a certain amount of iron, pass into solution. The solution is poured off, and chalk added to it; this precipitates the ferric oxide (the ferrous oxide is converted into ferric oxide in the furnace) and forms a fresh quantity of gypsum, while the chromic acid remains in solution—that is, it does not form the sparingly-soluble normal salt (1 part soluble in 240 parts of water). The solution then contains a fairly pure calcium dichromate, which by double decomposition gives other chromates; for example, with a solution of potassium sulphate it gives a precipitate of calcium sulphate and a solution of potassium dichromate, which crystallises when evaporated.[2]
Potassium dichromate, K2Cr2O7, easily crystallises from acid solutions in red, well-formed prismatic crystals, which fuse at a red heat and evolve oxygen at a very high temperature, leaving chromic oxide and the normal salt, which undergoes no further change: 2K2Cr2O7= 2K2CrO4+ Cr2O3+ O3. At the ordinary temperature 100 parts of water dissolve 10 parts of this salt, and the solubility increases as the temperature rises. It is most important to note that the dichromate does not contain water, it is K2CrO4+ CrO3; the acid salt corresponding to potassium acid sulphate, KHSO4, does not exist. It does not even evolve heat when dissolving in water, but on the contrary produces cold,i.e.it does not form a very stable compound with water. The solution and the salt itself are poisonous, and act as powerful oxidising agents, which is the character of chromic acid in general. When heated with sulphur or organic substances, with sulphurous anhydride, hydrogen sulphide, &c., this salt is deoxidised, yielding chromic compounds.[2 bis]Potassium dichromate[3]is used in the arts and in chemistry as a source for the preparation of all otherchromium compounds. It is converted into yellow pigments by means of double decomposition with salts of lead, barium, and zinc. When solutions of the salts of these metals are mixed with potassium dichromate (in dyeing generally mixed with soda, in order to obtain normal salts), they are precipitated as insoluble normal salts; for example, 2BaCl2+ K2Cr2O7+ H2O = 2BaCrO4+ 2KCl + 2HCl. It follows from this that these salts are insoluble in dilute acids, but the precipitation is not complete (as it would be with the normal salt). The barium and zinc salts are of a lemon yellow colour; the lead salt has a still more intense colour passing into orange. Yellow cotton prints are dyed with this pigment. The silver salt, Ag2CrO4, is of a bright red colour.
When potassium dichromate is mixed with potassium hydroxideor carbonate (carbonic anhydride being disengaged in the latter case) it forms thenormalsalt, K2CrO4, known asyellow chromate of potassium. Its specific gravity is 2·7, being almost the same as that of the dichromate. It absorbs heat in dissolving; one part of the salt dissolves in 1·75 part of water at the ordinary temperature, forming a yellow solution. When mixed even with such feeble acids as acetic, and more especially with the ordinary acids, it gives the dichromate, and Graham obtained a trichromate, K2Cr3O10= K2CrO4,2CrO3, by mixing a solution of the latter salt with an excess of nitric acid.
Chromic anhydrideis obtained by preparing a saturated solution of potassium dichromate at the ordinary temperature, and pouring it in a thin stream into an equal volume of pure sulphuric acid.[4]On mixing, the temperature naturally rises; when slowly cooled, the solution deposits chromic anhydride in needle-shaped crystals of a red colour sometimes several centimetres long. The crystals are freed from the mother liquor by placing them on a porous tile.[4 bis]It is very important at this point to call attention to the fact that a hydrate of chromic anhydride is never obtained in the decomposition of chromic compounds,but always theanhydride, CrO3. The corresponding hydrate, CrO4H2, or any other hydrate, is not even known. Nevertheless, it must be admitted that chromic acid is bibasic, because it forms salts isomorphous or perfectly analogous with the salts formed by sulphuric acid, which is the best example of a bibasic acid. A clear proof of the bibasicity of CrO3is seen in the fact that the anhydride and salts give (when heated with sodium chloride and sulphuric acid) a volatile chloranhydride, CrO2Cl2, containing two atoms of chlorine as a bibasic acid should.[5]Chromic anhydride is a red crystalline substance, which is converted into a black mass by heat; it fuses at 190°, and disengages oxygen above 250°, leaving a residue of chromium dioxide, CrO2,[6]and, on still further heating, chromic oxide, Cr2O3. Chromic anhydride is exceedingly soluble in water, and even attracts moisture from the air, but, as was mentioned above, it does not form any definite compound with water. The specific gravity of its crystals is 2·7, and when fused it has a specific gravity 2·6. The solution presents perfectly defined acid properties. It liberates carbonic anhydride from carbonates; gives insoluble precipitates of the chromates with salts of barium, lead, silver, and mercury.
The action of hydrogen peroxide on a solution of chromic acid or of potassium dichromate gives a blue solution, which very quickly becomes colourless with the disengagement of oxygen. Barreswill showed that this is due to the formation of aperchromic anhydride, Cr2O7, corresponding with sulphur peroxide. This peroxide is remarkable from the fact that it very easily dissolves in ether and is much more stable in this solution, so that, by shaking up hydrogen peroxide mixed with a small quantity of chromic acid, with ether, it is possible to transfer all the blue substance formed to the ether.[6 bis]
With oxygen acids, chromic acid evolves oxygen; for example, withsulphuric acid the following reaction takes place: 2CrO3+ 3H2SO4= Cr2(SO4)3+ O3+ 3H2O. It will be readily understood from this thata mixture of chromic acidorof its salts with sulphuric acidforms an excellentoxidising agent, which is frequently employed in chemical laboratories and even for technical purposes as a means of oxidation. Thus hydrogen sulphide and sulphurous anhydride are converted into sulphuric acid by this means. Chromic acid is able to act as a powerful oxidising agent because it passes into chromic oxide, and in so doing disengages half of the oxygen contained in it: 2CrO3= Cr2O3+ O3. Thus chromic anhydride itself is a powerful oxidising agent, and is therefore employed instead of nitric acid in galvanic batteries (as a depolariser), the hydrogen evolved at the carbon being then oxidised, and the chromic acid converted into a non-volatile product of deoxidation, instead of yielding, as nitric acid does, volatile lower oxides of offensive odour. Organic substances are more or less perfectly oxidised by means of chromic anhydride, although this generally requires the aid of heat, and does not proceed in the presence of alkalis, but generallyin the presence of acids. In acting on a solution of potassium iodide, chromic acid, like many oxidising agents, liberates iodine; the reaction proceeds in proportion to the amount of CrO3present, and may serve for determining the amount of CrO3, since the amount of iodine liberated can be accurately determined by the iodometric method (Chapter XX., Note42). If chromic anhydride be ignited in a stream of ammonia, it gives chromic oxide, water, and nitrogen. In all cases when chromic acid acts as an oxidising agent in the presence of acids and under the action of heat, the product of its deoxidation is a chromic salt, CrX3, which is characterised by the green colour of its solution, so that theredor yellowsolutionof a salt of chromic acid is then transformed into agreen solutionof a chromic salt, derived from chromic oxide, Cr2O3, which is closely analogous to Al2O3, Fe2O3, and other bases of the composition R2O3. This analogy is seen in the insolubility of the anhydrous oxide, in the gelatinous form of the colloidal hydrate, in the formation of alums,[7]of a volatile chloride of chromium, &c.[7 bis]
Chromic oxide, Cr2O3, rarely found, and in small quantities, in chrome ochre, is formed by the oxidation of chromium and its lower oxides, by the reduction of chromates (for example, of ammonium or mercuric chromate) and by the decomposition (splitting up) of the saline compounds of the oxide itself, CrX3or Cr2X6, like alumina, which it resembles in forming a feeble base easily giving double and basic salts, which are either green or violet.
The reduction of chromic oxide—for instance, in a solution by zinc and sulphuric acid—leads to the formation of chromous oxide, CrO, and its salts, CrX2, of a blue colour (seeNotes7and7bis). The further reduction[8]of oxide of chromium and its corresponding compounds givesmetallic chromium. Deville obtained it (probably containing carbon) by reducing chromic oxide with carbon, at a temperature near the melting point of platinum, about 1750°, but the metal itself does not fuse at this temperature. Chromium has a steel-grey colour and is very hard (sp. gr. 5·9), takes a good polish, and dissolves in hydrochloric acid, but cold dilute sulphuric and nitric acids have no action upon it. Bunsen obtained metallic chromium by decomposing a solution of chromic chloride, Cr2Cl6, by a galvanic current, as scales of a grey colour (sp. gr. 7·3). Wöhler obtained crystalline chromium by igniting a mixture of the anhydrous chromic chloride Cr2Cl6(seeNote7 bis) with finely-divided zinc, and sodium and potassium chlorides, at the boiling-point of zinc. When the resultant mass has cooled the zinc maybe dissolved in dilute nitric acid, and grey crystalline chromium (sp. gr. 6·81) is left behind. Frémy also prepared crystalline chromium by the action of the vapour of sodium on anhydrous chromic chloride in a stream of hydrogen, using the apparatus shown in the accompanying drawing, and placing the sodium and the chromic chloride in separate porcelain boats. The tube containing these boats is only heated when it is quite full of dry hydrogen. The crystals of metallic chromium obtained in the tube are grey cubes having a considerable hardness and withstanding the action of powerful acids, and even of aqua regia. The chromium obtained by Wöhler by the action of a galvanic current is, on the contrary, acted on under these conditions. The reason of this difference must be looked for in the presence of impurities, and in the crystalline structure. But in any case, among the properties of metallic chromium, the following may be considered established: it is white in colour, with a specific gravity of about 6·7, is extremely hard in a crystalline form, is not oxidised by air at the ordinary temperature, and with carbon it forms alloys like cast iron and steel.
see captionFig.92.—Apparatus for the preparation of metallic chromium by igniting chromic chloride and sodium in a stream of hydrogen.
Fig.92.—Apparatus for the preparation of metallic chromium by igniting chromic chloride and sodium in a stream of hydrogen.
The two analogues of chromium,molybdenumandtungsten(or wolfram), are of still rarer occurrence in nature, and form acid oxides, RO3, which are still less energetic than CrO3. Tungsten occurs in the somewhat rare minerals,scheelite, CaWO4, andwolfram; the latter being an isomorphous mixture of the normal tungstates of iron and manganese, (MnFe)WO4. Molybdenum is most frequently met with asmolybdenite, MoS2, which presents a certain resemblance to graphite in its physical properties and softness. It also occurs, but much more rarely, as a yellow lead ore, PbMoO4. In both these forms molybdenum occurs in the primary rocks, in granites, gneiss, &c., and in iron and copper oresin Saxony, Sweden, and Finland. Tungsten ores are sometimes met with in considerable masses in the primary rocks of Bohemia and Saxony, and also in England, America, and the Urals. The preliminary treatment of the ore is very simple; for example, the sulphide, MoS2, is roasted, and thus converted into sulphurous anhydride and molybdic anhydride, MoO3, which is then dissolved in alkalis, generally in ammonia. The ammonium molybdate is then treated with acids, when the sparingly soluble molybdic acid is precipitated. Wolfram is treated in a different manner. Most frequently the finely-ground ore is repeatedly boiled with hydrochloric and nitric acids, and the resultant solutions (of salts of manganese and iron) poured off, until the dark brown mass of ore disappears, whilst the tungstic acid remains, mixed with silica, as an insoluble residue; it is treated also with ammonia, and is thus converted into soluble ammonium tungstate, which passes into solution and yields tungstic acid when treated with acids. This hydrate is then ignited, and leaves tungstic anhydride. The general character of molybdic and tungstic anhydrides is analogous to that of chromic anhydride; they are anhydrides of a feebly acid character, which easily give polyacid salts and colloid solutions.[8 bis]
Hydrogen (which does not directly form compounds with Cr, Mo, and W) reduces molybdic and tungstic anhydride at a red heat; and this forms the means of obtaining metallic molybdenum and tungsten.Both metalsare infusible, and both under the action of heat form compounds with carbon and iron (the addition of tungsten to steel renders the latter ductile and hard).[9]Molybdenum forms a grey powder, which scarcely aggregates under a most powerful heat, and has a specific gravity of 8·5. It is not acted on by the air at the ordinary temperature, but when ignited it is first converted into a brown, and then into a blue oxide, and lastly into molybdic anhydride. Acids do not act on it—that is, it does not liberate hydrogen from them, not even from hydrochloric acid—but strong sulphuric acid disengages sulphurous anhydride, forming a brown mass, containing a lower oxide of molybdenum. Alkalis in solution do not act on molybdenum, but when fusedwith it hydrogen is given off, which shows, as does its whole character, the acid properties of the metal. The properties of tungsten are almost identical; it is infusible, has an iron-grey colour, is exceedingly hard, so that it even scratches glass. Its specific gravity is 19·1 (according to Roscoe), so that, like uranium, platinum, &c., it is one of the heaviest metals.[9 bis]Just as sulphur and chromium have their corresponding persulphuric and perchromic acids, H2S2O8and H2CrO8, having the properties of peroxides, and corresponding to peroxide of hydrogen, so also molybdenum and tungsten are known to givepermolybdicandpertungsticacids, H2Mo2O8and H2W2O8, which have the properties of true peroxides,i.e.easily disengage iodine from KI and chlorine from HCl, easily part with their oxygen, and are formed by the action of peroxide of hydrogen, into which they are readily reconverted (hence they may be regarded as compounds of H2O2with 2MoO3and 2WO3), &c. Their formation (Boerwald 1884, Kemmerer 1891) is at once seen in the coloration (not destroyed by boiling), which is obtained on mixing a solution of the salts with peroxide of hydrogen, and on treating, for instance, molybdic acid with a solution of peroxide of hydrogen (Péchard 1892). The acid then forms an orange-coloured solution, which after evaporation in vacuo leaves Mo2H2O84H2O as a crystalline powder, and loses 4H2O at 100°, beyond which it decomposes with the evolution of oxygen.[9 tri]
Uranium, U = 240, has the highest atomic weight of all the analogues of chromium, and indeed of all the elements yet known. Itshighest salt-forming oxide, UO3, shows very feeble acid properties. Although it gives sparingly-soluble yellow compounds with alkalis, which fully correspond with the dichromates—for example, Na2U2O7= Na2O,2UO3,[10]—yet it more frequently and easily reacts with acids, HX,forming fluorescent yellowish-green salts of the composition UO2X2, and in this respect uranic trioxide, UO3, differs from chromic anhydride, CrO3, although the latter is able to give the oxychloride, CrO2Cl2. In molybdenum and tungsten, however, we see a clear transition from chromium to uranium. Thus, for example, chromyl chloride, CrO2Cl2, is a brown liquid which volatilises without change, and is completely decomposed by water; molybdenum oxychloride, MoO2Cl2, is a crystalline substance of a yellow colour, which is volatile and soluble in water (Blomstrand), like many salts. Tungsten oxychloride, WO2Cl2, stands still nearer to uranyl chloride in its properties; it forms yellow scales on which water and alkalis act, as they do on many salts (zinc chloride, ferric chloride, aluminium chloride, stannic chloride, &c.), and perfectly corresponds with the difficultly volatile salt, UO2Cl2(obtained by Peligot by the action of chlorine on ignited uranium dioxide, UO2), which is also yellow and gives a yellow solution with water, like all thesalts UO2X2. The property of uranic oxide, UO3, of forming salts UO2X2is shown in the fact that the hydrated oxide of uranium, UO2(HO)2, which is obtained from the nitrate, carbonate, and other salts by the loss of the elements of the acid, is easily soluble in acids, as well as in the fact that the lower grades of oxidation of uranium are able, when treated with nitric acid, to form an easily crystallisable uranyl nitrate, UO2(NO3)2,6H2O; this is the most commonly occurring uranium salt.[11]
Uranium, which gives an oxide, UO3, and the corresponding salt UO2X2and dioxide UO2, to which the salts UX4correspond, is rarely met with in nature. Uranite or the double orthophosphate of uranicoxide, R(UO2)H2P2O8,7H2O, where R = Cu or Ca, uranium-vitriol U(SO4)2,H2O, samarakite, and æschynite, are very rarely found, and then only in small quantities. Of more frequent and abundant occurrence is the non-crystalline, earthy brown uranium ore known aspitchblende(sp. gr. 7·2), which is mainly composed of the intermediate oxide, U3O8= UO2,2UO3. This ore is found at Joachimsthal in Bohemia and in Cornwall. It usually contains a number of different impurities, chiefly sulphides and arsenides of lead and iron, as well as lime and silica compounds. In order to expel the arsenic and sulphur it is roasted, ground, washed with dilute hydrochloric acid, which does not dissolve the uranoso-uranic oxide, U3O8, and the residue is dissolved in nitric acid, which transforms the uranium oxide into the nitrate, UO2(NO3)2.
It must be observed that the oxide of uranium, first distinguished by Klaproth (1789), was for a long time regarded as able to give metallic uranium under the action of charcoal and other reducing agents (with the aid of heat). But the substance thus obtained was only theuranium dioxide, UO2. The compound nature of this dioxide,[12]or the presence of oxygen in it, was demonstrated by Peligot (1841), by igniting it with charcoal in a stream of chlorine. He thus obtained a volatileuranium tetrachloride, UCl4,[13]which, when heated with sodium, gavemetallic uraniumas a grey metal, having a specific gravity of 18·7, and liberating hydrogen from acids, with the formation of green uranous salts, UX4, which act as powerful reducing agents.[14]
As the salts of uranic oxide are reduced in the absence of organic matter by the action of light, and as they impart a characteristic coloration to glass,[15]they find a certain application in photography and glass work.
If we compare together the highly acid elements, sulphur, selenium, and tellurium, of the uneven series, with chromium, molybdenum, tungsten, and uranium of the even series, we find that the resemblance of the properties of the higher form RO3does not extend to the lower forms, and even entirely disappears in the elements, for there is not the smallest resemblance between sulphur and chromium and their analogues in a free state. In other words, this means that the small periods, like Na, Mg, Al, Si, P, S, Cl, containing seven elements, do not contain any near analogues of chromium, molybdenum, &c., and therefore their true position among the other elements must be looked for only in those large periods which contain two small periods, and whose type is seen in the period containing: K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se, Br. These large periods contain Ca and Zn, giving RO, Sc, and Ga of the third group, Ti and Ge giving RO2, V and As forming R2O5, Cr and Se of the sixth group, Mn and Br of the seventh group, and the remaining elements, Fe, Co, Ni, form connective members of the intermediate eighth group, to the description of the representatives of which we shall turn in the following chapters. We will now proceed to describemanganese, Mn = 55, as an element of the seventh group of the even series, directly following after Cr = 52, which corresponds with Br = 80 to the same degree that Cr does with Se = 79. For chromium, selenium, and bromine very close analogues are known, but for manganese as yet none have been obtained—that is, it is the only representative of the even series in the seventh group. In placing manganese with thehalogens in one group, the periodic system of the elements only requires that it should bear an analogy to the halogens in the higher type of oxidation—i.e.in the salts and acids—whilst it requires that as great a difference should be expected in the lower types and elements as there exists between chromium or molybdenum and sulphur or selenium. And this is actually the case. The elements of the seventh group form a higher salt-forming oxide, R2O7, and its corresponding hydrate, HRO4, and salts—for example, KClO4. Manganese in the form of potassium permanganate, KMnO4, actually presents a great analogy in many respects to potassium perchlorate, KClO4. The analogy of the crystalline form of both salts was shown by Mitscherlich. The salts of permanganic acid are also nearly all soluble in water, like those of perchloric acid, and if the silver salt of the latter, AgClO4, be sparingly soluble in water, so also is silver permanganate, AgMnO4. The specific volume of potassium perchlorate is equal to 55, because its specific gravity = 2·54; the specific volume of potassium permanganate is equal to 58, because its specific gravity = 2·71. So that the volumes of equivalent quantities are in this instance approximately the same whilst the atomic volumes of chlorine (35·5/1·3 = 27) and manganese (55/7·5) are in the ratio 4 : 1. In a free state the higher acids HClO4and HMnO4are both soluble in water and volatile, both are powerful oxidisers—in a word, their analogy is still closer than that of chromic and sulphuric acids, and those points of distinction which they present also appear among the nearest analogues—for example, in sulphuric and telluric acids, in hydrochloric and hydriodic acids, &c. Besides Mn2O7manganese gives a lower grade of oxidation, MnO3, analogous to sulphuric and chromic trioxides, and with it corresponds potassium manganate, K2MnO4, isomorphous with potassium sulphate.[16]In the still lower grades of oxidation, Mn2O3and MnO, there is hardly any similarity to chlorine, whilst every point of resemblance disappears when we come to the elements themselves—i.e.to manganese and chlorine—for manganese is a metal, like iron, which combines directly with chlorine to form a saline compound, MnCl2, analogous to magnesium chloride.[17]
Manganese belongs to the number of metals widely distributed innature, especially in those localities where iron occurs, whose ores frequently contain compounds of manganous oxide, MnO, which presents a resemblance to ferrous oxide, FeO, and to magnesia. In many minerals magnesia and the oxides allied to it are replaced by manganous oxide; calcspars and magnesites—i.e.R″CO3in general—are frequently met with containing manganous carbonate, which also occurs in a separate state, although but rarely. The soil also and the ash of plants generally contain a small quantity of manganese. In the analysis of minerals it is generally found that manganese occurs together with magnesia, because, like it, manganous oxide remains in solution in the presence of ammoniacal salts, not being precipitated by reagents. The property of this manganous oxide, MnO, of passing into the higher grades of oxidation under the influence of heat, alkalis, and air, gives an easy means not only of discovering the presence of manganese in admixture with magnesia, but also of separating these two analogous bases. Magnesia is not able to give higher grades of oxidation, whilst manganese gives them with great facility. Thus, for instance, analkalinesolution of sodium hypochlorite produces a precipitate of manganese dioxide in a solution of a manganous salt: MnCl2+ NaClO + 2NaHO = MnO2+ H2O + 3NaCl; whilst magnesia is not changed under these circumstances, and remains in the form of MgCl2. If the magnesia be precipitated owing to the presence of alkali, it may be dissolved in acetic acid, in which manganese dioxide is insoluble. The presence of small quantities of manganese may also be recognised by the green coloration which alkalis acquire when heated with manganese compounds in the air. This green coloration depends on the property of manganese of giving a green alkaline manganate: MnCl2+ 4KHO + O2= K2MnO4+ 2KCl + 2H2O. Thusthe faculty of oxidising in the presence of alkalisforms an essential character of manganese. The higher grades of oxidation containing Mn2O7and MnO3are quite unknown in nature, and even MnO2is not so widely spread in nature as the ores composed of manganous compounds which are met with nearly everywhere. The most important ore of manganese is its dioxide, or so-calledperoxide, MnO2, which is known in mineralogy aspyrolusite. Manganese also occurs as an oxide corresponding with magnetic iron ore, MnO,Mn2O3= Mn3O4, forming the mineral known ashausmannite. The oxide Mn2O3also occurs in nature as the anhydrous mineralbraunite, and in a hydrated form, Mn2O3,H2O, calledmanganite. Both of these often occur as an admixture in pyrolusite. Besides which, manganese is met with in nature as a rose-coloured mineral,rhodonite, or silicate, MnSiO3. Very fine and rich deposits of manganese ores have been found in the Caucasus, the Urals, and along the Dnieper. Those at the Sharapanskydistrict of the Government of Kutais and at Nicopol on the Dnieper are particularly rich. A large quantity of the ore (as much as 100,000 tons yearly) is exported from these localities.
Thus manganese gives oxides of the following forms: MnO, manganous oxide, and manganous salts, MnX2, corresponding with the base, which resembles magnesia and ferrous oxide in many respects; Mn2O3, a very feeble base, giving salts, MnX3, analogous to the aluminium and ferric salts, easily reduced to MnX2; MnO2, dioxide, generally called peroxide, an almost indifferent oxide, or feebly acid;[18]MnO3, manganic anhydride, which forms salts resembling potassium sulphate;[18 bis]Mn2O7, permanganic anhydride, giving salts analogous to the perchlorates.
All the oxides of manganese when heated with acids give salts, MnX2, corresponding with the lower grade of oxidation,manganous oxide, MnO. Manganic oxide, Mn2O3, is a feebly energetic base; it is true that it dissolves in hydrochloric acid and gives a dark solution containing the salt MnCl3, but the latter when heated evolves chlorine and gives a salt corresponding with manganous oxide MnCl2—i.e.at first: Mn2O3+ 6HCl = 3H2O + Mn2Cl6, and then the Mn2Cl6decomposes into 2MnCl2+ Cl2. None of the remaining higher grades of oxidation have a basic character, butact as oxidising agents in the presence of acids, disengaging oxygen and passing into salts of the lower grade of oxidation of manganese, MnO. Owing to this circumstance,the manganous saltsare often obtained; they are, for instance, left in the residue when the dioxide is used for the preparation of oxygen and chlorine.[19]
As the salts of manganous oxide MnX2closely resemble (and are isomorphous with) the salts of magnesia MgX2in many respects (with the exception of the fact that MnX2are rose coloured and are easily oxidised in the presence of alkalis), we will not dwell upon them, butlimit ourselves to illustrating the chemical character of manganese by describing the metal and its corresponding acids. The fact alone that the oxides of manganese are not reduced to the metal when ignited in hydrogen (whilst the oxides of iron give metallic iron under these circumstances), but only to manganous oxide, MnO, shows that manganese has a considerable affinity for oxygen—that is, it is difficult to reduce. This may be effected, however, by means of charcoal or sodium at a very high temperature. A mixture of one of the oxides of manganese with charcoal or organic matter gives fusedmetallic manganeseunder the powerful heat developed by coke with an artificial draught. The metal was obtained for the first time in this manner by Gahn, after Pott, and more especially Scheele, had in the last century shown the difference between the compounds of iron and manganese (they were previously regarded as being the same). Manganese is prepared by mixing one of its oxides in a finely-divided state with oil and soot; the resultant mass is then first ignited in order to decompose the organic matter, and afterwards strongly heated in a charcoal crucible. The manganese thus obtained, however, contains, as a rule, a considerable amount of silicon and other impurities. Its specific gravity varies between 7·2 and 8·0. It has a light grey colour, a feebly metallic lustre, and although it is very hard it can be scratched by a file. It rapidly oxidises in air, being converted into a black oxide; water acts on it with the evolution of hydrogen—this decomposition proceeds very rapidly with boiling water, and if the metal contain carbon.[20]
It has been shown above that if manganese dioxide, or any lower oxide of manganese, be heated with an alkali in the presence of air, the mixture absorbs oxygen,[21]and forms an alkaline manganate of a green colour: 2KHO + MnO2+ O = K2MnO4+ H2O. Steam is disengaged during the ignition of the mixture, and if this does not take place there is no absorption of oxygen. The oxidation proceeds much more rapidly if, before igniting in air, potassium chlorate or nitre be added to the mixture, and this is the method of preparingpotassium manganate, K2MnO4. The resultant mass dissolved in a small quantity of water gives a dark green solution, which, when evaporated under the receiver of an air-pump over sulphuric acid, deposits green crystals of exactly the same form as potassium sulphate—namely, six-sided prisms and pyramids. The composition of the product is not changed by being re-dissolved, if perfectly pure water free from air and carbonic acid be taken. But in the presence of even very feeble acids the solution of this salt changes its colour and becomes red, and deposits manganese dioxide. The same decomposition takes place when the salt is heated with water, but when diluted with a large quantity of unboiled water manganese dioxide does not separate, although the solution turns red. This change of colour depends on the fact that potassium manganate, K2MnO4, whose solution is green, is transformed into potassium permanganate, KMnO4, whose solution is of a red colour. The reaction proceeding under the influence of acids and a large quantity of wateris expressed in the following manner: 3K2MnO4+ 2H2O = 2KMnO4+ MnO2+ 4KHO. If there is a large proportion of acid and the decomposition is aided by heat, the manganese dioxide and potassium permanganate are also decomposed, with formation of manganous salt. Exactly the same decomposition as takes place under the action of acids is also accomplished by magnesium sulphate, which reacts in many cases like an acid. When water holding atmospheric oxygen in solution acts on a solution of potassium manganate, the oxygen combines directly with the manganate and forms potassium permanganate, without precipitating manganese dioxide, 2K2MnO4+ O + H2O = 2KMnO4+ 2KHO. Thus a solution of potassium manganate undergoes a very characteristic change in colour and passes from green to red; hence this salt received the name ofchameleon mineral.[22]
Potassium permanganate, KMnO4, crystallises in well-formed, long red prisms with a bright green metallic lustre. In the arts the potash is frequently replaced by soda, and by other alkaline bases, but no salt of permanganic acid crystallises so well as the potassium salt, and therefore this salt is exclusively used in chemical laboratories. One part of the crystalline salt dissolves in 15 parts of water at the ordinary temperature. The solution is of a very deepred colour, which is so intense that it is still clearly observable after being highly diluted with water. In a solid state it is decomposed by heat, with evolution ofoxygen, a residue consisting of the lower oxides of manganese and potassium oxide being left.[22 bis]A mixture of permanganate of potassium, phosphorous and sulphur takes fire when struck or rubbed, a mixture of the permanganate with carbon only takes fire when heated, not when struck. The instability of the salt is also seen in the fact that its solution is decomposed by peroxide of hydrogen, which at the same time it decomposes. A number of substances reduce potassium permanganate to manganese dioxide (in which case the red solution becomes colourless).[23]Many organic substances (although far from all, even when boiled in a solution of permanganate) act in this manner, being oxidised at the expense of a portion of its oxygen. Thus, a solution of sugar decomposes a cold solution of potassium permanganate. In the presence of an excess of alkali, with a small quantity of sugar, the reduction leads to the formation of potassium manganate, because 2KMnO4+ 2KHO = O + 2K2MnO4+ H2O. With a considerable amount of sugar and a more prolonged action, the solution turns brown and precipitates manganese dioxide or even oxide. In the oxidation of many organic bodies by an alkaline solution of KMnO4generally three-eighths of the oxygen in the salt are utilised for oxidation: 2KMnO4= K2O + 2MnO2+ O3. A portion of the alkali liberated is retained by the manganese dioxide, and the other portion generally combines with the substance oxidised, because the latter most frequently gives an acid with an excess of alkali. A solution of potassium iodide acts in a similar manner, being converted into potassium iodate at the expense of the three atoms of oxygen disengaged by two molecules of potassium permanganate.
In the presence of acids, potassium permanganate acts as an oxidising agentwith still greater energy than in the presence of alkalis. At any rate, a greater proportion of oxygen is then available for oxidation, namely, not ⅜, as in the presence of alkalis, but ⅝, because in the first instance manganese dioxide is formed, and in the second case manganous oxide, or rather the salt, MnX2, corresponding with it. Thus, forinstance, in the presence of an excess of sulphuric acid, the decomposition is accomplished in the following manner: 2KMnO4+ 3H2SO4= K2SO4+ 2MnSO4+ 3H2O + 5O. This decomposition, however, does not proceed directly on mixing a solution of the salt with sulphuric acid, and crystals of the salt even dissolve in oil of vitriol without the evolution of oxygen, and this solution only decomposes by degrees after a certain time. This is due to the fact that sulphuric acid liberates free permanganic acid from the permanganate,[24]which acid is stable in solution. But if, in the presence of acids and a permanganate, thereis a substance capable of absorbing oxygen—for instance, capable of passing into a higher grade of oxidation—then the reduction of the permanganic acid into manganous oxides sometimes proceeds directly at the ordinary temperature. This reduction is very clearly seen, because the solutions of potassium permanganate are red whilst the manganous salts are almost colourless. Thus, for instance, nitrous acid and its salts are converted into nitric acid and decolorise the acid solution of the permanganate. Sulphurous anhydride and its salts immediately decolorise potassium permanganate, forming sulphuric acid. Ferrous salts, and in general salts of lower grades of oxidation capable of being oxidised in solution, act in exactly the same manner. Sulphuretted hydrogen is also oxidised to sulphuric acid; even mercury is oxidised at the expense of permanganic acid, and decolorises its solution, being converted into mercuric oxide. Moreover, the end point of these reactions may easily be seen, and therefore, having first determined the amount of active oxygen in one volume of a solution of potassium permanganate, and knowing how many volumes are required to effect a given oxidation, it is easy to determine the amount of an oxidisable substance in a solution from the amount of permanganate expended (Marguerite's method).
The oxidising action of KMnO4, like all other chemical reactions, is not accomplished instantaneously, but only gradually. And, as the course of the reaction is here easily followed by determining the amount of salt unchanged in a sample taken at a given moment,[25]the oxidising reaction of potassium permanganate, in an acid liquid, was employed by Harcourt and Esson (1865) as one of the first cases for the investigation of the laws of therate of chemical change[26]as a subject of great importance in chemical mechanics. In their experiments they took oxalic acid,C2H2O_4, which in oxidising gives carbonic anhydride, whilst, with an excess of sulphuric acid, the potassium permanganate is converted into manganous sulphate, MnSO4, so that the ultimate oxidation will be expressed by the equation: 5C2H2O4+ 2MnKO4+ 3H2SO4= 10CO2+ K2SO4+ 2MnSO4+ 8H2O. The influence of the relative amount of sulphuric acid is seen from the annexed table, which gives the measure of reactionpper 100 parts of potassium permanganate, taken four minutes after mixing, using n molecules of sulphuric acid, H2SO4, per 2KMnO4+ 5C2H2O4:
showing that in a given time (4 minutes) the oxidation is the more perfect the greater the amount of sulphuric acid taken for given amounts of KMnO4and C2H2O4. It is obvious also that the temperature and relative amount of every one of the acting and resulting substances should show its influence on the relative velocity of reaction; thus, for instance, direct experiment showed the influence of the admixture of manganous sulphate. When a large proportion of oxalic acid (108 molecules) was taken to a large mass of water and to 2 molecules of permanganate 14 molecules of manganous sulphate were added, the quantity x of the potassium permanganate acted on (in percentages of the potassium permanganate taken) in t minutes (at 16°) was as follows:
These figures show that the rate of reaction—that is, the quantity of permanganate changed in one minute—decreases proportionally to the decrease in the amount of unchanged potassium permanganate. At the commencement, about 2·6 per cent. of the salt taken was decomposed in the course of one minute, whilst after an hour the rate was about 0·5 per cent. The same phenomena are observed in every case which has been investigated, and this branch of theoretical or physical chemistry, now studied by many,[27]promises to explain the course of chemical transformations from a fresh point of view, which is closely allied to the doctrine of affinity, because the rate of reaction, without doubt, is connected with the magnitude of the affinities acting between the reacting substances.