1911 Encyclopædia Britannica/Manganese
MANGANESE [symbol Mn; atomic weight, 54.93 (O = 16)], a metallic chemical element. Its dioxide (pyrolusite) has been known from very early times, and was at first mistaken for a magnetic oxide of iron. In 1740 J. H. Pott showed that it did not contain iron and that it yielded a definite series of salts, whilst in 1774 C. Scheele proved that it was the oxide of a distinctive metal. Manganese is found widely distributed in nature, being generally found to a greater or less extent associated with the carbonates and silicates of iron, calcium and magnesium, and also as the minerals braunite, hausmannite, psilomelane, manganite, manganese spar and hauerite. It has also been recognized in the atmosphere of the sun (A. Cornu, Comptes rendus, 1878, 86, pp. 315, 530), in sea water, and in many mineral waters.
The metal was isolated by J. G. Gahn in 1774, and in 1807 J. F. John (Gehlen’s Jour. chem. phys., 1807, 3, p. 452) obtained an impure metal by reducing the carbonate at a high temperature with charcoal, mixed with a small quantity of oil. R. Bunsen prepared the metal by electrolysing manganese chloride in a porous cell surrounded by a carbon crucible containing hydrochloric acid. Various reduction methods have been employed for the isolation of the metal. C. Brunner (Pogg. Ann., 1857, 101, p. 264) reduced the fluoride by metallic sodium, and E. Glatzel (Ber., 1889, 22, p. 2857) the chloride by magnesium, H. Moissan (Ann. Chim. Phys., 1896 (7) 9, p. 286) reduced the oxide with carbon in the electric furnace; and H. Goldschmidt has prepared the metal from the oxide by means of his “thermite” process (see Chromium). W. H. Green and W. H. Wahl [German patent 70773 (1893)] prepare a 97% manganese from pyrolusite by heating it with 30% sulphuric acid, the product being then converted into manganous oxide by heating in a current of reducing gas at a dull red heat, cooled in a reducing atmosphere, and finally reduced by heating with granulated aluminium in a magnesia crucible with lime and fluorspar as a flux. A purer metal is obtained by reducing manganese amalgam by hydrogen (O. Prelinger, Monats., 1894, 14, p. 353).
Prelinger’s manganese has a specific gravity of 7.42, and the variety obtained by distilling pure manganese amalgam in vacuo is pyrophoric (A. Guntz, Bull. Soc. [3], 7, 275), and burns when heated in a current of sulphur dioxide. The pure metal readily evolves hydrogen when acted upon by sulphuric and hydrochloric acids, and is readily attacked by dilute nitric acid. It precipitates many metals from solutions of their salts. It is employed commercially in the manufacture of special steels. (See Iron and Steel.)
Compounds
Manganese forms several oxides, the most important of which are manganous oxide, MnO, trimanganese tetroxide, Mn3O4, manganese sesquioxide, Mn2O3, manganese dioxide, MnO2, manganese trioxide, MnO3, and manganese heptoxide, Mn3O7.
Manganous oxide, MnO, is obtained by heating a mixture of anhydrous manganese chloride and sodium carbonate with a small quantity of ammonium chloride (J. v. Liebig and F. Wöhler, Pogg. Ann., 1830, 21, p. 584); or by reducing the higher oxides with hydrogen or carbon monoxide. It is a dark coloured powder of specific gravity 5.09. Manganous hydroxide, Mn(OH)2, is obtained as a white precipitate on adding a solution of a caustic alkali to a manganous salt. For the preparation of the crystalline variety identical with the mineral pyrochroite (see A. de Schulten, Comptes rendus, 1887, 105, p. 1265). It rapidly oxidizes on exposure to air and turns brown, going ultimately to the sesquioxide. Trimanganese tetroxide, Mn3O4, is produced more or less pure when the other oxides are heated. It may be obtained crystalline by heating manganese sulphate and potassium sulphate to a bright red heat (H. Debray, Comptes rendus, 1861, 52, p. 985). It is a reddish-brown powder, which when heated with hydrochloric acid yields chlorine. Manganese sesquioxide, Mn2O3, found native as the mineral braunite, may be obtained by igniting the other oxides in a mixture of nitrogen and oxygen, containing not more than 26% of the latter gas (W. Dittmar, Jour. Chem. Soc., 1864, 17, p. 294). The hydrated form, found native as the mineral manganite, is produced by the spontaneous oxidation of manganous hydroxide. In the hydrated condition it is a dark brown powder which readily loses water at above 100° C., it dissolves in hot nitric acid, giving manganous nitrate and manganese dioxide: 2MnO(OH) + 2HNO3 = Mn(NO3)2 + MnO2 + 2H2O. Manganese dioxide, or pyrolusite (q.v.), MnO2, the most important oxide, may be prepared by heating crystallized manganous nitrate until red fumes are given off, decanting the clear liquid, and heating to 150° to 160° C. for 40 to 60 hours (A. Gorgen, Bull. Soc., 1890 [3], 4, p. 16), or by heating manganese carbonate to 260° C. in the presence of air and washing the residue with very dilute cold hydrochloric acid. It is a hard black solid which readily loses oxygen when strongly heated, leaving a residue of Mn3O4. When heated with concentrated hydrochloric acid it yields chlorine, and with concentrated sulphuric acid it yields oxygen. It is reduced to the monoxide when heated in a current of hydrogen. It is a strong oxidizing agent. It dissolves in cold concentrated hydrochloric acid, forming a dark brown solution which probably contains manganic chloride (see R. J. Meyer, Zeit. anorg. Chem., 1899, 22, p. 169; G. Neumann, Monats., 1894, 15, p. 489). It is almost impossible to prepare a pure hydrated manganese dioxide owing to the readiness with which it loses oxygen, leaving residues of the type xMnO·yMnO2. Such mixtures are obtained by the action of alkaline hypochlorites on manganous salts, or by suspending manganous carbonate in water and passing chlorine through the mixture. The solid matter is filtered off, washed with water, and warmed with 10% nitric acid (A. Gorgen). It is a dark brown powder, which reddens litmus. Manganese dioxide combines with other basic oxides to form manganites, and on this property is based the Weldon process for the recovery of manganese from the waste liquors of the chlorine stills (see Chlorine). The manganites are amorphous brown solids, insoluble in water, and decomposed by hydrochloric acid with the evolution of chlorine. Manganese trioxide, MnO3, is obtained in small quantity as an unstable deliquescent red solid by dropping a solution of potassium permanganate in sulphuric acid on to dry sodium carbonate (B. Franke, Jour. prak. Chem., 1887 [2], 36, p. 31). Above 50° C. it decomposes into the dioxide and oxygen. It dissolves in water forming manganic acid, H2MnO4. Manganese heptoxide, Mn2O7, prepared by adding pure potassium permanganate to well cooled, concentrated sulphuric acid, when the oxide separates as a dark oil (H. Aschoff, Pogg. Ann., 1860, 111, p. 217), is very unstable, continually giving off oxygen. It decomposes violently on heating, and explodes in contact with hydrogen, sulphur, phosphorus, &c. It dissolves in water to form a deep red solution which contains permanganic acid, HMnO4. This acid is also formed by decomposing barium or lead permanganate with dilute sulphuric acid. It is only known in aqueous solution. This solution is of a deep violet-red colour, and is somewhat fluorescent; it decomposes on exposure to light, or when heated. It is a monobasic acid, and a very powerful oxidizing agent (M. M. P. Muir, Jour. Chem. Soc., 1907, 91, p. 1485).
Manganous Salts.—The anhydrous chloride, MnCl2, is obtained as a rose-red crystalline solid by passing hydrochloric acid gas over manganese carbonate, first in the cold and afterwards at a moderate red heat. The hydrated chloride, MnCl2·4H2O, is obtained in rose-red crystals by dissolving the metal or its carbonate in aqueous hydrochloric acid and concentrating the solution. It may be obtained in at least two different forms, one isomorphous with NaCl·2H2O, by concentrating the solution between 15° C. and 20°C.; the other, isomorphous with FeCl2·4H2O, by slow evaporation of the mother liquors from the former. It forms double salts with the chlorides of the alkali metals. The bromide MnBr2·4H2O, iodide, MnI2, and fluoride, MnF2, are known.
Manganous Sulphate, MnSO4, is prepared by strongly heating a paste of pyrolusite and concentrated sulphuric acid until acid fumes cease to be evolved. The ferric and aluminium sulphates present are thus converted into insoluble basic salts, and the residue yields manganous sulphate when extracted with water. The salt crystallizes with varying quantities of water, according to the temperature at which crystallization is effected: between −4° C. and +6° C. with 7H2O, between 15° C. and 20° C. with 5H2O, and between 25° C. and 31° C. with 4H2O. It crystallizes in large pink crystals, the colour of which is probably due to the presence of a small quantity of manganic sulphate or of a cobalt sulphate. It combines with the sulphates of the alkali metals to form double salts.
Manganous Nitrate, Mn(NO3)2·6H2O, obtained by dissolving the carbonate in nitric acid and concentrating the solution, crystallizes from nitric acid solutions in long colourless needles, which melt at 25.8° C. and boil at 129.5° C. with some decomposition.
Manganous Carbonate, MnCO3, found native as manganese spar, may be prepared as an amorphous powder by heating manganese chloride with sodium carbonate in a sealed tube to 150° C., or in the hydrated form as a white flocculent precipitate by adding sodium carbonate to a manganous salt. In the moist condition it rapidly turns brown on exposure to air.
Manganous Sulphide, MnS, found native as manganese glance, may be obtained by heating the monoxide or carbonate in a porcelain tube in a current of carbon bisulphide vapour. R. Schneider (Pogg. Ann., 1874, 151, 449) obtained a crystalline variety by melting sulphur with anhydrous manganous sulphate and dry potassium carbonate, extracting the residue and drying it in a current of hydrogen. Four sulphides are known; the red and green are anhydrous, a grey variety contains much water, whilst the pink is a mixture of the grey and red (J. C. Olsen and W. S. Rapalje, Jour. Amer. Chem. Soc., 1904, 26, p. 1615). Ammonium sulphide alone gives incomplete precipitation of the sulphide. In the presence of ammonium salts the precipitate is dirty white in colour, whilst in the presence of free ammonia it is a buff colour. This form of the sulphide is readily oxidized when exposed in the moist condition, and is easily decomposed by dilute mineral acids.
Manganese Disulphide, MnS2, found native as hauerite, is formed as a red coloured powder by heating manganous sulphate with potassium polysulphide in a sealed tube at 160°–170° C. (H. v. Senarmont, Jour. prak. Chem., 1850, 51, p. 385).
Manganic Salts.—The sulphate, Mn2(SO4)3, is prepared by gradually heating at 138° C. a mixture of concentrated sulphuric and manganese dioxide until the whole becomes of a dark green colour. The excess of acid is removed by spreading the mass on a porous plate, the residue stirred for some hours with nitric acid, again spread on a porous plate, and finally dried quickly at about 130° C. It is a dark green deliquescent powder which decomposes on heating or on exposure to moist air. It is readily decomposed by dilute acids. With potassium sulphate in the presence of sulphuric acid it forms potassium manganese alum, K2SO4·Mn2(SO4)2·24H2O. A. Piccini (Zeit. anorg. Chem. 1898, 17, p. 355) has also obtained a manganese caesium alum. Manganic Fluoride, MnF3, a solid obtained by the action of fluorine on manganous chloride, is decomposed by heat into manganous fluoride and fluorine. By suspending the dioxide in carbon tetrachloride and passing in hydrochloric acid gas, W. B. Holmes (Abst. J.C.S., 1907, ii., p. 873) obtained a black trichloride and a reddish-brown tetrachloride.
Manganese Carbide, Mn3C, is prepared by heating manganous oxide with sugar charcoal in an electric furnace, or by fusing manganese chloride and calcium carbide. Water decomposes it, giving methane and hydrogen (H. Moissan); Mn3C + 6H2O = 3Mn(OH)2 + CH4 + H2.
Manganates.—These salts are derived from manganic acid H2MnO4. Those of the alkali metals are prepared by fusing manganese dioxide with sodium or potassium hydroxide in the presence of air or of some oxidizing agent (nitre, potassium chlorate, &c.); MnO2 + 2KHO + O = K2MnO4 + H2O. In the absence of air the reaction proceeds slightly differently, some manganese sesquioxide being formed; 3MnO2 + 2KHO = K2MnO4 + Mn2O3 + H2O. The fused mass has a dark olive-green colour, and dissolves in a small quantity of cold water to a green solution, which is, however, only stable in the presence of an excess of alkali. The green solution is readily converted into a pink one of permanganate by a large dilution with water, or by passing carbon dioxide through it: 3K2MnO4 + 2CO2 = 2K2CO3 + 2KMnO4 + MnO2.
Permanganates are the salts of permanganic acid, HMnO4. The potassium salt, KMnO4, may be prepared by passing chlorine or carbon dioxide through an aqueous solution of potassium manganate, or by the electrolytic oxidation of the manganate at the anode [German patent 101710 (1898)]. It crystallizes in dark purple-red prisms, isomorphous with potassium perchlorate. It acts as a powerful oxidizing agent, both in acid and alkaline solution; in the first case two molecules yield five atoms of available oxygen and in the second, three atoms:
| 2KMnO4 + 3H2SO4 | = K2SO4 + 2MnSO4 + 3H2O + 5O; |
| 2KMnO4 + 3H2O | = 2MnO2·H2O + 2KHO + 3O. |
It completely decomposes hydrogen peroxide in sulphuric acid solution—
It decomposes when heated to
and when warmed with hydrochloric acid it yields chlorine:
Sodium Permanganate, NaMnO4.3H2O (?), may be prepared in a similar manner, or by precipitating the silver salt with sodium chloride. It crystallizes with great difficulty. A solution of the crude salt is used as a disinfectant under the name of “Condy’s fluid.”
Ammonium Permanganate, NH4·MnO4, explodes violently on rubbing, and its aqueous solution decomposes on boiling (W. Muthmann, Ber., 1893, 26, p. 1018); NH4·MnO4 = MnO2 + N2 + 2H2O.
Barium Permanganate, BaMn2O8, crystallizes in almost black needles, and is formed by passing carbon dioxide through water containing suspended barium manganate.
Detection.—Manganese salts can be detected by the amethyst colour they impart to a borax-bead when heated in the Bunsen flame, and by the green mass formed when they are fused with a mixture of sodium carbonate and potassium nitrate. Manganese may be estimated quantitatively by precipitation as carbonate, this salt being then converted into the oxide, Mn3O4 by ignition; or by precipitation as hydrated dioxide by means of ammonia and bromine water, followed by ignition to Mn3O4. The valuation of pyrolusite is generally carried out by means of a distillation with hydrochloric acid, the liberated chlorine passing through a solution of potassium iodide, and the amount of iodine liberated being ascertained by means of a standard solution of sodium thiosulphate.
The atomic weight of manganese has been frequently determined. J. Berzelius, by analysis of the chloride, obtained the value 54.86; K. v. Hauer (Sitzb. Akad. Wien., 1857, 25, p. 132), by conversion of the sulphate into sulphide, obtained the value 54.78; J. Dewar and A. Scott (Chem. News, 1883, 47, p. 98), by analysis of silver permanganate, obtained the value 55.038; J. M. Weeren (Stahl. u. Eisen, 1893, 13, p. 559), by conversion of manganous oxide into the sulphate obtained the value 54.883, and of the sulphate into sulphide the value 54.876 (H = 1), and finally G. P. Baxter and Hines (Jour. Amer. Chem. Soc., 1906, 28, p. 1360), by analyses of the chloride and bromide, obtained 54.96 (O = 16).