THE SOLUBILITY OF FERROUS SULPHATE
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BY FRANK
K. CAMERON
Non-aqueous Solvents. Ferrous sulphate is not soluble in ammonia,1 carbon dioxide,2 alcohol,3 glacial actic acid, methyl acetate,4 or ethyl acetate.5 It is slightly soluble in sulphuric acid,6 a saturated solution containing 0.22 percent FeS04 at 3o.2°C and 0.63 percent at 63.8°C. The solid phase in contact with these solutions contains both ferrous and hydrogen sulphates but in undetermined proportions. Similar solids containing ferrous sulphate and hydrogen sulphate or ferrous sulphate, hydrogen sulphate, and water have been prepared.7 The limits of concentration of sulphuric acid between which the several solids are stable, have been determined, but not their solu-
bilities. Various Properties. Ferrous sulphate is quite soluble in water, and with noticeable contraction.8 Extensive tables have been prepared of the specific heats of its solutions in water and aqueous sulphuric acid,9 the boiling points10 of aqueous solutions of varying composition, and the specific gravities11 at i5°C. Agde and Barkholt12have determined the specific gravities of saturated solutions from one degree to 8o°C. At this last temperature it is 1.367. At 54°, a short way below the transition temperature of the heptahydrate to the tetrahydrate, the specific gravity of the saturated solution is 1.432 and it falls continuously to 1.114 at one degree. The electrical conductivity of aqueous solutions at 2S°C. has been determined by Wagner.13 The dielectric constant14 has been found to decrease and then rise with increasing content of ferrous sulphate. The surface tension of water is slightly increased by dissolving ferrous sulphate.15 The solutions are more or less toxic, and have been used as insecticides, fungicides, weed killers, etc. Ferrous sulphate is but slightly toxic to fish.16 The solutions are astringent and have been employed as coagulants and as a primer before painting resinous woods.17 1
2
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5 8
Franklin: Am. Chem. J., 20, 828 (1898). Büchner: Z. physik. Chem., 54, 674 (1905). Anthon: J. prakt. Chem., 14, 125 (1838). Naumann: Ber., 42, 3790 (1909). Naumann: Ber., 37, 3601 (1904). Kendall and Davidson: J. Am. Chem. Soc., 43, 979 (1921).
Kenrick: J. Phys. Chem., 12, 704 (1908). 8Rakschit: Z. Electrochemie, 31, 97 (1925); 32, 276 (1926). 9 Agde and Holtmann: Z. anorg. allgem. Chem., 158, 316 (1926). 7
10Gerlach: Z. anal. Chem., 26, 426 (1887). “Gerlach: Z. anal. Chem., 8, 287 (1869). 12Z. angew. Chem., 39, 851 (1926). 13Z. physik. Chem., 71, 429 (1910). “Heilman and Zahn: Ann. Physik, 81, 711 (1926). l5Stocker: Z. physik. Chem., 94, 149 (1920). 16Belding: Trans. Am. Fish Assoc., 57, no (1927). I7Brooke: Philipp. J. Sci., 30, 303 (1926).
SOLUBILITY
OF FERROUS SULPHATE
693
Solutions show the phenomena of creeping but to a slight extent as compared to those of many other salts.1 Ferrous sulphate heptahydrate effloresces. In water, ferrous sulphate hydrolyzes and the determination of the “free acidity” has received much attention in recent years, electrometric titration seeming to be favored.2 In the order of the salting out of ions,3 Fe' lies between Mg and Zn. In contact with zeolites or soil minerals, Fe' in aqueous solutions of ferrous sulphate is displaced4 *by Ca. It is also displaced8 readily by Ba, but not by Be. Oxidation. In aqueous solution ferrous sulphate is readily, and sometimes annoyingly, oxidized by air. Jilek6 finds no oxidation at the end of fortyeight hours if sulphuric acid be present. Banerjee7 finds the oxidation by air to be slow, an unimolecular reaction, approximately, hastened by the presence of potassium sulphate but retarded by all other sulphates, particularly sulphuric acid and copper sulphate. Reedy and Machen8 find the oxidation to be slow, to fall off gradually, but to be positively catalyzed by pyrolusite (Mn02). This last fact is the basis of patents and commercial practice. Potassium permanganate, potassium dichromate, iodine chloride9 (IC1) are readily reduced, and their solutions are mediums for the analytical estimation of ferrous sulphate. Chlorine10 is used commercially as is also sodium peroxide. The reaction with hydrogen peroxide is not well understood. Manchot and Lehmann11 find that in dilute solutions of ferrous sulphate one Fe' is equivalent to 3H202, probably Fe205 being formed; while, in concentrated solutions, one Fe' may be equivalent to as much as 24 H202. In acid solutions ferrous sulphate is oxidized by X-rays irrespective of the wave length.12 It induces the oxidation of other substances, and is important for various autoxidations as with glycolic acid by hydrogen peroxide.13 From the literature it appears that the best way to prevent oxidation of ferrous sulphate or its solutions is to keep them in contact with hydrogen. Some investigators have found an atmosphere of nitrogen satisfactory. Contact with iron wire or nails is unsatisfactory. A layer of nujol has been moderately successful for a few days, but not over a period of weeks. Satisfactory results have been attained by using water which has been long boiled for making solutions and keeping the solutions in contact with carbon dioxide. '
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Druce: Pharm. J., 119, 333 (1927); Washburn: J. Phys. Chem., 31, 1246 (1927). •Koenig: Chimie et Industrie, Special No. 187 (1926); Haczko: Z. anal. Chem., 73, 404 (1928); Kamienski: Bull, intern. Acad. Polonaise, 1928, 33. Randall and Failey: Chem. Reviews, 4, 285 (1927). Magistad: Arizona Agr. Exp. Sta., Tech. Bull. 18, 445 (1928). 1
3
4
Bodforss: Z. physik. Chem., 130, 82 (1927). •Chem. Listy, 15, 105; 138 (1921). 7 Proc. Asiatic Soc. Bengal, 18, No. 6, 71 (1922); Z. anorg. allgem. Chem., 128, 343 (1923). 8 Ind. Eng. Chem., 15, 1271 (1923). •Heisig: J. Am. Chem. Soc., 50, 1687 (1928). 10Mohlman and Palmer: Eng. News Record, 100, 147 (1928). 6
“Ann., 460,
179 (1928).
“Fricke and Morse: Am. Jour. Roentgenology and Radium Theraphy, 18, 426 (1927;) Strahlentherapie, 26 749 (1927); Ber. ges. Physiol, expt. Pharmakol., 44, 336. ‘•Goldschmidt, Askenasy, and Fierros: Ber., 61, 223 (1928).
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FRANK K. CAMERON
At higher temperatures
an atmosphere of water vapor alone has proved quite sufficient to prevent noticeable oxidation over periods of several weeks, and at lower temperatures the presence of a few percent of alcohol has proved effective. Hydrates of Ferrous Sulphate. At ordinary temperatures the heptahydrate, FeSCh.yHiO, is the stable solid, separating from an aqueous solution as deep green, monoclinic crystals. Rhombic crystals1 have been observed and can be induced by seeding the mother liquor with corresponding heptahydrates of other bases which crystallize in the rhombic system.2 Westerbrink3 by studying spectrograms found it to be monoclinic. It has a specific gravity of 1.889 according to Roscoe and Schorlemmer, quoting Joule and Playfair. Retgers4 found it to be 1.898 at 18.0C. In contact with its saturated aqueous solution it is stable from the cryohydrate point, i.82°C, to s6.6°C, according to Fraenckel,5 the latter being a transition point at which the tetrahydrate becomes the stable form. Tilden6 found the melting point of the heptahydrate to be 64°C. It loses water readily. Heated in vacuo at i4o°C it is transformed to the monohydrate and on further gentle heating out of contact with the air, the anhydrous salt is formed. Liversidge7 found heating in a water oven for 90 minutes left a residue containing 82.5% FeS04; and Pritzer and Jungkunz8 found six molecules of water are removed when the heptahydrate is heated in xylene. Schumb9 found that at 25°C. the dissociation pressure is 14.56 mm Hg for the transformation FeS04.7H20 to FeS04.6H20. Cohen and Visser10 are quoted by Jorissen11 as having found 1.91 Kalories for the transformation: FeS04.4H20 + sH20 FeS04.7H20. The molecular volume of the salt and the hydrating water molecules were determined in the classical investigation of Thorp and Watts,12 and recently by Moles and Crespi,13 who found 13.4 cm3for the first, 16.3 cm3 for the remaining water molecules. The hexahydrate, FeS04.6H>0, is described by Lecoq de Boisbaudran,14 and by Hensgen.15 The former obtained it by seeding a solution of ferrous sulphate, slightly under-saturated with respect to FeS04.7H20, with a crystal of cobalt sulphate crystallized at 5o°C, CoS04.6H20. The compound —
=
Rammelsberg: Fogg. Ann., 91, 32: (1854); Volger: Jahrb. Mineralogie, 1855, 152. Roscoe and Schorlemmer: “Treatise on Chemistry”, (1911). 3 Verslag Akad. Wetenschapen, Amsterdam, 35, 1913; Proc. Acad. Sci. Amsterdam, 29, 1223 (1926). 4 Z. physik. Chem., 3, 534 (1889). 6 Z. anorg. Chem., 55, 223 (1907). “ J. Chem. Soc., 45, 267 (1884). 7 Pharm. J., 118, 106; Chemist and Druggist, 106, 141 (1927). 8 Chem. Ztg., 50, 962 (1926). 9 J. Am. Chem. Soc., 45, 364 (1923). “Arch, néerl., (2), 5, 300 (1900). "Landolt and Bornstein: 3rd. Edition, 463, 1905; Z. physik. Chem., 74, 308 (1910). 12J. Chem. Soc., 37, 102 (1840). 13Z. physik. Chem., 130, 337 (1927). 14Ann. Chim. Phys., (4), 18, 255 1869). “Ber., 11, 1776 (1878). 1
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