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 metals are 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 fused with 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, H2S2O8 and H2CrO8, having the properties of peroxides, and corresponding to peroxide of hydrogen, so also molybdenum and tungsten are known to give permolybdic and pertungstic acids, H2Mo2O8 and 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 H2O2 with 2MoO3 and 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. Its highest 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 the salts UO2X2. The property of uranic oxide, UO3, of forming salts UO2X2 is 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 UO2X2 and dioxide UO2, to which the salts UX4 correspond, is rarely met with in nature. Uranite or the double orthophosphate of uranic oxide, 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 as pitchblende (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 the uranium 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 volatile uranium tetrachloride, UCl4,[13] which, when heated with sodium, gave metallic uranium as 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 RO3 does 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 describe manganese, 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 the halogens 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 HClO4 and HMnO4 are 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 Mn2O7 manganese 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, Mn2O3 and 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 in nature, 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″CO3 in 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, an alkaline solution 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. Thus the faculty of oxidising in the presence of alkalis forms an essential character of manganese. The higher grades of oxidation containing Mn2O7 and MnO3 are quite unknown in nature, and even MnO2 is 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-called peroxide, MnO2, which is known in mineralogy as pyrolusite. Manganese also occurs as an oxide corresponding with magnetic iron ore, MnO,Mn2O3 = Mn3O4, forming the mineral known as hausmannite. The oxide Mn2O3 also occurs in nature as the anhydrous mineral braunite, and in a hydrated form, Mn2O3,H2O, called manganite. 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 Sharapansky district 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.