T H E SOLUBILITY OF FERROUS SULPHATE BY FRANK K. CAMERON
Son-aqueous Solvents. Ferrous sulphate is not soluble in ammonia,' carbon dioxide,2 alcoh01,~glacial actic acid, methyl acetateI4or ethyl acetate.j I t is slightly soluble in sulphuric acid,6 a saturated solution containing 0 . 2 2 percent FeS04 a t 30.2OC and 0.63 percent a t 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.' The limits of concentration of sulphuric acid between which the several solids are stable, have been determined, but not their solubilities. Varzous 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,gthe boiling points'O of aqueous solutions of varying composition, and the specific gravitiesll a t I 5OC. Agde and Barkholtl*have determined the specific gravitiesof saturated solutions from one degree to 8 o T . At this last temperature it is 1.367. At 54O, 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 a t one degree. The electrical conductivity of aqueous solutions at 25OC. has been determined by Wagner.I3 The dielectric constant14has been found to decrease and then risewith increasing content of ferrous sulphate. The surface tension of water is slightly increased by dissolving ferrous ~u1phate.I~ 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." . . Franklin: Am. Chem. J., 20, 828 (1898). Biichner: Z. physik. Chem., 54, 674 (190j). * Anthon: J. prakt. Chem., 14, I Z j (1838). Naumann: Rer., 42, 3790 (1909). 5 Saumann: Ber., 37, 3601 (1904). OKendall and Davidson: J. Am. Chem. Soc., 43, 979 (1921). Kenrick: J. Phys. Chem., 12, 704 (1908). *Rakechit: 2. Electrochemie, 31, 9; (1925); 32, 276 (1926). 8 Agde and Holtmann: 2. anorg. allgem. Chem., 158, 316 (1926). 'OGerlach: Z. anal. Chem., 26, 426 (1887). "Gerlach: Z. anal. Chem., 8, 2 8 ; (1869). I2Z. angew. Chem., 39, 851 (1926). 1%. physik. Chem., 71, 429 (1910). 14Hellman and Zahn: Ann. Physik, 81, 711 (1926). Wtocker: Z. physik. Chem., 94, 149 (1920). IeBelding: Trans. Am. Fish hssoc., 57, I I O (1927). "Brooke: 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.' 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 I n the order of the salting out of ions,3 Fe" lies between Mg and Zn. I n contact with zeolites or soil minerals, F e ' ' in aqueous solutions of ferrous sulphate is displaced4 by Ca. It is also displaced5 readily by Ba, but not by Be. Ozidation. I n aqueous solution ferrous sulphate is readily, and sometimes annoyingly, oxidized by air. Jilek6 finds no oxidation a t the end of fortyeight hours if sulphuric acid be present. Banerjee' 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 Machens find the oxidation to be slow, to fall off gradually, but to be positively catalyzed by pyrolusite (MnOp). This last fact is the basis of patents and commercial practice. Potassium permanganate, potassium dichromate, iodine chlorideg (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 Lehmann" find that in dilute solutions of ferrous sulphate one F e ' ' is equivalent to 3H202, probably Fe205being formed; while, in concentrated solutions, one Fe ' may be equivalent to as much as 24 HzOz. I n 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 pe~oxide.'~ 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. h 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. Druce: Pharm. J., 119, 333 (1927);Washburn: J. Phys. Chem., 31, 1246 (1927). 2Koenig: Chimie et Industrie, Special No. 187 (1926);Haczko: 2. anal. Chem., 73, 404 (1928);Kamienski: Bull. intern. Acad. Polonaise, 1928, 33. 3Randall and Failey: Chem. Reviews, 4,285 (1927). Magistad: Arizona Agr. Exp. Sta., Tech. Bull. 18, 445 (1928). Bodforss: 2. physik. Chem., 130, 82 (1927). BChem.Listy, 15, 105; 138 (1921). Proc. Asiatic SOC.Bengal, 18,No. 6,71 (1922);Z.anorg. allgem. Chem., 128,343 (1923). Ind. Eng. Chem., 15, 1271 (1923). PHeisig: J. Am.Chem. SOC.,50, 1687 (1928). 'Wohlman and Palmer: Eng. Kews Record, 100, 147 (1928). "Ann., 460, 179 (1928). I2Fricke and Morse: Am. Jour. Roentgenology and h d i u m Theraphy, 18,426 (1927;) Strahlentherapie, 26 749 (1927);Ber. ges. Physiol. expt. Pharmakol., 44, 336. 13Goldschmidt, Askenaay, and Pierros: Ber., 61,223 (1928). 1
'
694
FRAXK 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 os Ferrous Sulphate. At ordinary temperatures the heptahydrate, FeS04.7H20,is the stable solid, separating from an aqueous solution as deep green, monoclinic crystals. Rhombic crystals' 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 1.898a t 18.0C. In contact with its saturated aqueous solution it is stable from the cryohydrate point, -1.8z0C, to j6.6'C, according to F r a e n ~ k e l the , ~ latter being a transition point a t which the tetrahydrate becomes the stable form. Tilden6 found the melting point of the heptahydrate to be 64OC. It loses water readily. Heated in vacuo at 14oOC 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 8 2 . 5 % FeS04; and Pritzer and Jungkunz* found six molecules of water are removed when the heptahydrate is heated in xylene. Schumbg found that a t 2 5 O C . the dissociation pressure is 14.j6 mm Hg for the transformation FeS04.7H20 to FeS04.6H20. Cohen and Visserlo are quoted by Jorissenll as having found 1.91Kalories for the transformation: FeS04.4H20 3H2O = FeS04.7H20. The molecular volume of the salt and the hydrating water molecules were determined in the classical investigation of Thorp and Ratts,12 and recently by Moles and Cre~pi.~3who found 13.4 cm3for the first, 16.3 cm3 for the remaining water molecules. The hexahydrate, FeSO4.6H?O,is described by Lecoq de Boi~baudran,'~ and by Hensgen.'5 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 a t 5ooC, CoS04.6H20. The compound
+
Rammelsberg: Pogg. Ann., 91, 321 (1854); Volger: Jahrb. Mineralogie, 1855, 152. Schorlemmer: "Treatise on Chemistry", (1911). 3 Verslsg &ad. Wetenschapen, Amsterdam, 35, 1913; Proc. Acad. Sci. Amsterdam, 29, 1223 (1926). 'Z. physik. Chem., 3, 534 (1889). 6 Z. anorg. Chem., 55, 223 (1907). J. Chem. SOC.,45, 267 (1884). 'Pharm. J., 118, 106; Chemist and Druggist, 106, 141 (1927). 8Chem. Ztg., 50, 962 (1926). J. Am. Chem. SOC.,45, 364 (1923). 'OArch. nkrl., (2), 5 , 300 (1900). "Landolt and Bornstein: 3rd. Edition, 463, 1905; Z. physik. Chem., 74, 308 (1910). I2J. Chem. Soc., 37, 102 (1840). I3Z. physik. Chem., 130, 337 (1917). "Ann. Chim. Phys., (4), 18, 255 (1869). 'SBer., 11, 1776 (1878).
* Roscoe and
SOLUBILITY OF FERROUS SULPHATE
69;
is metastable, however, and is soon transformed to the heptahydrate. I t is pale green in color. Crystals with faces of several mms. dimensions were obtained. Hensgen obtained it as green needles, by treating the heptahydrate with concentrated hydrochloric acid. The pentahydrate, FeS0,.5H20, is said by Roscoe and Schorlemmer to form when a solution of ferrous sulphate in aqueous sulphuric acid is evaporated in vacuum, the heptahydrat,e first separating, then the pentahydrate, and finally the tetrahydrate, isomorphous with the corresponding manganese salt. It is probable that this sequence does not take place, as will be shown later in this paper. Lecoq de Boisbaudran' found that seeding with cupric sulphate pentahydrate will induce the separation of ferrous sulphate pentahydrate from an aqueous solution as a metastable form quickly transforming to the heptahydrate. He states that it is difficult to obtain and only from solutions much supersatured with respect to the heptahydrate. I t was not observed by Agde and Barkholt,z nor by Cameron and Crockford3 and the evidence for the possible existence of this hydrate needs confirmation. The tetrahydrate, FeS04.4H20, was found by FraenckeP to be the stable form in contact with its aqueous solutions between j6.6"C and 64.4'C and he describes it as bright green in color. I t will be shown presently that it is stable in contact with a saturated aqueous solution to 67.4'C and that it is stable a t 65OC in solutions containing as much as 2 . 5 percent sulphuric acid, The crystals obtained were quite small and faintly green. When dried, the crystals remained clear, apparently quite stable, and unaffected by exposure to the air for weeks. The trihydrate, FeS04.3H20,has been reported by Kanes and by Kuhne.6 Kane obtained it by crystallization from a solution saturated with hydrochloric acid, as bright-green, hard, transparent crystals, but could not. determine their form. The water found on analysis agreed with that calculabed for the trihydrate. The dihydrate, FeS04.zH20,has been reported by von Bonsdorff,7 in a reference not' accessible to the writer. As will be shown presently it is stable in contact with its saturated aqueous solution above 67.4'C. It is stable at 65'C, in contact wit'h aqueous solutions containing more than about 2 . 5 percent sulphuric acid. It was obtained as a very finely divided white crystalline powder quite stable in the air when dry, dissolving at ordinary temperah r e s more slowly than the other hydrates obtained in this investigation. It packs on a filter to a dense mass, which may be washed only with difficulty. The monohydrate, FeS01.H20, was found by Fraenckel to be the stable solid in contact with the saturated aqueous solutions above 64.4OC; but l Loc. cit. See also Narignac: Ann. d. Mines, ( 5 ) , 9, 9 (1856); Lecoq de Boisbaudran, Liebig and Kopp: Jahresber., 1867, I 52. Z. angew. Chem., 39, 851 (1926). a J . Phys. Chem., 33, 7wj (1929). ' Z. anorg. Chem., 5 5 , 223 (1907). sAnn., 19, 7 (1836). Schweigger's Journal, 61, 235 (1831). ' Bericht uber d. Versammlung deutsch. Naturforscher und &zte in Prag, 1837, 124.
696
FRANK K. CAMERON
he is mistaken, as will be shown presently. It will be shown also that it is the stable compound below this temperature in contact with solutions containing higher concentrations of sulphuric acid. It is a snow white, crystalline, fine powder. When dry it is stable in the air, without appreciable oxidation on long standing, or other obvious change. The anhydrous salt, FeS04, is not stable in contact with aqueous solutions under any conditions, so far as now known. I t is readily prepared by heating a hydrate in a neutral atmosphere, or better in hydrogen. But it is also decomposed rather readily on heating, the thermal decomposition having interested a number of investigators,’ and recently Greulich,* who finds it decomposes according to the equation: 2FeS04 = Fez03 SO3 SO*.At atmospheric pressure the dissociation temperature is probably 680°C and the heat of dissociation 189.5 calories. T h e E f e c t of Temperature. For comparison the data of Etard? Fraencke1,l Agde and Barkholt: and the International Critical Tables,6 have been recalculated to the basis of percent FeS04 in solution. The data of Brandes and Firnhaber,’ Tobler,s and M ~ l d e rare , ~ not included as being of historical interest only. The results are assembled in Table I. Below 6o0C the figures of Etard and Agde and Barkholt agree very well with the International Critical Table. Etard’s figures for 130°C and 1 5 2 O C lie fairly near to extrapolated values from the International Critical Table, but his figures below 60°C and for 13oOC do not. I n good agreement with the International Critical Table is Agde and Barkholt’s figure a t 80°C, and Schreinmakers”O 24.89 percent at 3oOC. Too high are Wirth’s” figure of 2 2 . 8 percent a t 2 5 O C , Occleshaw’sl? 22.98 percent at 25OC and Schiff’s1337.2 percent a t 20°C, while H a u e r ’ ~ ’17.02 ~ percent at rs°C is too low. Fraenckel determined two transition temperatures approximately, but all his solubility data are too high. It has been assumed hitherto that the stable solid in contact with its saturated solution in water, above 64O, is the monohydrate. It is so stated in the International Critical Tables. Fraenckel thought he had proved it. He used the van Bijlert’s method with sodium chloride as “tell-tale” to
+
+
Marchal: J. Chim. phys., 22, 325 (1925); Keppeler and D’Ans: Z. physik. Chem., 62, 89 (1908). 2 Z. anorg. allgem. Chem., 168, 19; (192;). Ann. Chim. Phys., (;), 2, 553 (1894). 4 2.anorg. Chem., 55, 223 (IF;). 5 Z . angew. Chem., 39, 851 (1926). 6 “International Critical Tables”, 4, 224 (1928). 7 Archiv. Apothekerverein in nordl. Deutschland, VII, 83 Ann., 95, 193 (1855). 9 “Scheikundige Verhandelingen,” 111, 3, 141 (1864). 1°Z. physik. Chem., 71, I I O (1910). 1’Z. anorg. Chem., 79, 364 (1913). I*J. Chem. SOC.,127, 2598 (1925). IsAnn., 118, 362 (1861). L‘J. prakt. Chem., 103, 114 (1868). l6Z. physik. Chem., 8, 343 (1891).
SOLUBILITY O F FERROUS SULPHATE
697
determine the composition of the solid phases a t so", 64" and 80°C. At this last temperature he found FeO to be 42.55 per cent, calculated 42.30, corresponding to the monohydrate FeSO,.H,O. At all three temperatures he found good agreement with calculated figures in spite of the recognized difficulties and uncertainties of the analytical procedure. But the results to be detailed later on the solubility of ferrous sulphate in aqueous solutions of sulphuric acid have made necessary a re-examination of the solubility in water alone, the transition point where the tetrahydrate is in stable contact with a lower hydrate, and the composition of the lower hydrate. To this end the heptahydrate was recrystallized from dilute alcohol, filtered from the mother liquor, washed with alcohol, then with ether, and dried by pressing between bibulous paper. This product was then added gradually to hot distilled water which had just been boiled vigorously for an hour, a current of hydrogen gas being bubbled continually through the mass, until there was, approximately, 50 grams of solid in contact with IOO ccs. of saturated solution. The containing bottle was tightly closed with a new rubber stopper, and alternately heated and cooled for 36 hours. No signs of oxidation appearing, it was placed in a thermostat at 75"C, being momentarily withdrawn from time to time for vigorous shaking. After being in the thermostat for 24 hours, it was seeded with about 5 gms. of monohydrate. After another 24 hours, a sample of the solid phase was freed from mother liquor by suction, washed with alcohol and ether, and dried by pressing between filter paper. Calculated from its iron content, it contained 3.8 moles water per mole ferrous sulphate. The mother liquor contained 35.03 per cent ferrous sulphate. The mass was then seeded with about 2 gms. of the dihydrate. After another 24 hours, 5.3808 gms. of the washed and dried solid phase contained 4.1744 gms. FeSO,. It contained 2.4 mols water per mole ferrous sulphate and the mother liquor contained 32.97 per cent ferrous sulphate. After another 24 hour period, 5.6315 gms. of the solid contained 4.4386 gms. FeS04, and the water content was therefore 2 . 2 moles per mole FeS04, while the mother liquor contained 31.58 per cent FeS0,. A day later the mother liquor contained 31.02 per cent FeSO,, and two days later two separate samples of the mother liquor gave respectively 31.06 and 31.04 per cent FeSO,. It seems safe to conclude, therefore, that the stable solid phase a t 7 5 " in contact with a saturated water solution is the dihydrate and the solubility at this temperature is 3 1.04 gms. FeS04 per IOO gms. solution. A large sample, about 2 0 0 gms., of heptahydrate, crystallized from a dilute alcohol solution, washed with alcohol and ether, and dried between bibulous paper, was suspended in freshly boiled water t o which a small amount of alcohol had been added. The containing flask was fitted with a reflux condenser and the mass boiled for about 5 hours a t a temperature of 96' rt. The solid was then separated quickly from the mother liquor on a Buchner funnel, and, after washing with alcohol, then ether, and dried, 4.0725 gms. were found, through an iron determination, to contain 3.2761
698
FRANK K . CAMERON
gin-. FeS04, corresponding to a water content of 2 05 moles per mole ferrous sulphate. Fracnckel found the temperature for the transition heptahydrate to tetrahydrate, to be ~ 6 . ~ 5C 6by a dilatometer. JTe have found it to be not higher than j6O.68C by a Bremer-Frowein differential tensimeter.’ Fraenckel also found, with a dilatometer, a transition point at 64.O70 C, and cites in support a “Iinicke” in the solubility curve of RIulder a t 63.O5, and one in the solubility curve of Etard at 65”, as well as the melting point of 64’ for the heptahydrate as found by Tilden. If the melting point of the heptahydrate be 64.OC as found by Tilden, it cannot be the transition temperature for the tetrahydrate to the dihydrate or to the monohydrate as assumed by Fraenckel and the International Critical Tables.
56.8
6,O
6.1
but d&8
67d
3
ZO
Ttm\ c r s L . ; c FIG.I Solubility of Ferrous Sulphate in Water, showing temperatures a t Transition Points. Solid lines stable, broken lines metastable conditions.
A Bremer-Frowein tensimeter was arranged with heptahydrate in one arm and tetrahydrat,e, with a few drops of water, in t,he other. On raising the temperature a tenth of a degree about every 1 2 hours, equal vapor pressure in both arms was attained a t about 56.8”C. A second instrument was prepared with tetrahydrate in one arm and dihydrate in the other. A small addition of heptahydrate was made to each before evacuating and sealing. This instrument showed equal vapor tensions in both arms a t about 68’. Both instruments, after cooling, were very slowly heated in a large water bath, readings of the difference in level of the oil gauge and the temperature being made at minute intervals. Similar readings were made with the temperature falling. The response in vapor pressure to change in temperature was fairly prompt although there was an appreciable lag. Five transition points or temperature a t which the vapor pressures in the two arms closely approached equality, were discovered. The procedure was repeated a number of times, until the instruments indicated but two transitions, the one a t Z. physik. Chem., 1, 5 ; 362 (1888); ’7,
260
(1891); 17,
52
(1895).
SOLUBILITY O F FERROUS SULPHATE
699
about 57' and the other at about 68'. The average of the determinations, five sets on a rising thermometer with two sets on a falling thermometer, gave transitions at about jj', 61", 64.4', 65.8' and 67.8'. Of these the one a t 64.4' seemed to be the most sharply and clearly determined, the one a t 6j.8" the least. These data seem t o be most reasonably interpreted as in Fig. I .
TABLE I Comparison of Data. The Effect of Temperature on the Solubility of Ferrous Sulphate, FeS04, in Water. Stated in Percentages of Saturated Solution. T. Etard Fraenckel Agde and International Solid Phase
- IO'
Barkholt
-1.82' I'
I3 .79
20.0
17.02
20.90
33.0 26.3
24.70 28.70 30.04 32.50 34.50
54.58 j5.02
35.48 35.73
37.7
70.
77'
33.79 37.8
80.
86 90. 94. 102. 112,
130. 152.
43 .o
30.20
30.43
37.8 27.20
36.7 34.7 28.0 '7.3 2 . 5
,, 1
26.56
50,
67.
1'
22.98
36.4
ff
1,
22.7
25.
54. 56.6 60. 64.
,, 1)
21.30
21.
,, , ,, >>
I j , I O
20.85 26.42
10.0
30. 34. 40. 43 '
'3.58
15.1
9.6'
24.
? FeS04.7H20
14.98 15.62
0'
5O
Critical Tables
13 . o
),
,, ff
,, FeSOa.7Hn0 FeS04.4H20
7 00
FRANK K. CAMERON
The line AB is a part of the solubility curve of the heptahydrate based on the figures in the International Critical Tables. It is continued to D since all investigators agree that there is at ransition at 64'.4C. Assuming the real existence of this point and the correctness of my solubility determination a t 75OC (circle) a t G, the line DEFG is drawn. Assuming the correctness of the International Critical Table's solubility a t 60" and a transition at 56.8'C, the line B F is drawn, showing a transition a t 67.4OC instead of 67.8' as found in the tensimeter measurements. The line C E is the solubility curve of a hydrate of undetermined composition metastable over the range of temperature indicated. I am indebted to Mr. A. E. Hughes for the tensimeter readings and t o Mr. R. H. Munch for constructing the apparatus; and more particularly for building and maintaining the thermostats necessary for the experiments to be described presently. T h e effect of other electrolytes on the solubility of ferrous sulphate in water is generally quite marked. Electrolytes without a common ion increase the solubility; but qualitative data only are available. Systems containing another sulphate as well as ferrous sulphate and water are often mentioned in the literature. For instance, FeS04.7H20,is isomorphous with CoSO4. 7Hz0, MgS04.7Hz0, etc. FeS04.(NH4)2S04is isomorphous with ZnS04. (NH4)zS04.Many similar cases are cited; but the attention generally, has been concentrated upon the solid phases, with scanty or no data of value upon the accompanying liquid phases. I n the following paragraphs are summarized the principal studies pertinent to this investigation. T h e System: Ferrous Sulphate-Sodium Sulphate-Water has been studied by Koppel.' The cryohydrate temperature for ferrous sulphate and water is -2'C, while that for ferrous sulphate, sodium sulphate decahydrate, and water is -3'C. With both salts present as solid phases, the solubility of each seems to be slightly greater than when present by itself up to 18".5-18'.8C. At this temperature a double salt, FeS04.Na2S04.4H20,becomes a stable solid phase, The liquid phase at this point contains 18.3 per cent FeS04 and 13.8 per cent sodium sulphate. Continued addition of ferrous sulphate heptahydrate with rise of temperature is accompanied by an increase in concentration of ferrous sulphate on a smooth curve from the cryohydrate point to 40'C. But the solubility of the sodium sulphate decreases, the concentration falling on a straight line. When sodium sulphate decahydrate is added in excess, the solubility of the sodium sulphate increases with rising temperature, and the ferrous sulphate decreases, until 31.4'C is reached, when the two stable solids are the double salt and the anhydrous sodium sulphate. Curiously, further rise in temperature produces no change in composition of the liquid phase. Also, there is no change in composition of the liquid phase in contact with the double salt alone from z0.5'C (below which the latter is not stable by itself) up to 40%. Z. physik. Chem., 52, 405 (1905).
SOLUBILITY O F FERROUS SULPHATE
701
The double salt can be prepared by melting together the components, but better by bringing them together in solution and adding an excess of sulphuric acid. Washed with alcohol and ether and dried, the white double salt is quite stable in the air. The publication does not give data from which isotherms can be plotted. The System: Ferrous Sulphate-Lithium Sulphate-Water, a t 3ooC, has been studied by Schreinemakers with Reind1er.l The two salts mutually depress one anothers solubility in water. There is a transition point nrith a liquid phase containing about 16.1 per cent ferrous sulphate (FeS04) and 16.j per cent lithium sulphate, LilS04. Solutions richer in ferrous sulphate are in contact with ferrous sulphate heptahydrate, FeS04.7H20, as stable solid phase. Solutions richer in lithium sulphate are in contact with lithium sulphate monohydrate, Li,S04.H20,as the stable solid. The System: Ferrous Sulphate-Ammonium Sulphate-Water, at 3ooC, has been studied by Schreinemakers with van Meurs.2 The two salts mutually depress each other's solubility in water. The double salt, FeS04.(NH4)2 S04.6H20, is the stable solid phase in contact with solutions between the limits 0.79 per cent FeS04 - 43.88 per cent (XH4)2S04.and 25.24 per cent FeS04, - 5.91 per cent (XH4)2S04. Solutions richer in ferrous sulphate are in contact with ferrous sulphate heptahydrate, FeS04.7H20. Those richer in ammonium sulphate are in contact with the anhydrous salt, (NH&SO4, as stable solid phase. The Quartenary System: Ferrous Sulphate-Lithzum Sulphate-Ammonium Sulphate-Water, at 30°C,was also studied by Schreinemakers, and found to have three "constant solutions" in contact with three solid phases. The composition of the solutions and the formulas of the accompanying solid phases a t the several transition points are assembled in Table 11.
TABLE I1 Percentage Composition of Liquid Phases and Formulas of Accompanying Solids in the System, FeS04-Li2S04- (NH4)2S04-H20 at 30°C FeS04
LizS04
16.1 16.85
16.5 15.62
percent
2 j .22
percent
(NHI)2SOI percent 0.
4.82 5.93
0.00 0.00
43.86
6.23
40.48
0.00
6.59
39.55
4.15
20.03
12.32
.79 .61
' Z . physik. Chem., 71, * Z . physik. Chem., 71,
IIO
(1910).
I I I (1910).
Solid Phases
FcS04.7H20 - Li2S04.H20 FeS04. 7H20- Li2S04.H20-FeS04 (NH~)zSO~.~HZO FeS04.7H20 - FeS04.(NH4)zS04.6H20 FeS04.(NH4)2S04.6H20 - (NH4),SOa FeS04.(NH4)zS04.6H20 - (NH4),S04 - (SH4)2SO4.Li2SO4 (NH4)&04- (NH4)2S04.LiZS04 FeS04. (NH4)?S04.6H?O- ("a)zSO4. Li2S04-Li2S04.H20
702
FRANK K. CAMERON
T h e System: Ferrous Sulphate-Magnesium Sulphate-Water has been studied by Kammelsberg,’ Retgers,2and Barker3; but quantitative data for the liquid phase are yet lacking. The two salts mutually depress one anothers solubility. Two series of solid solutions, or mixed crystals are found. I n one, containing from zero to j 7 per cent MgS0,.7H20, the crystals are monoclinic. In the other, containing from 7j to I O O per cent JIgS04.7H20, the crystals are rhombic. Retgers worked mainly at 2oo-23OC. He determined the specific gravities, from which he computed the molecular volumes, the purpose of the investigation being to obtain an insight into the factors involved in isomorphism. T h e System: Ferrous Sulphate-Alu?7iznuni Sulphate-Water, a t 2 goC, has been studied by Occleshaw.4 The two salts mutually depress each other’s solubility. There are three solubility curves. Hydrated aluminum sulphate, A12(S04)3.181120, is the stable solid phase in contact with solutions containing less than 4.13 per cent ferrous sulphate, FeS04. At this point the solution contains 25.4 per cent aluminum sulphate, X1,(S04)3. From this point, until a liquid phase is reached containing 10.17 per cent FeS04 and 20.16 per cent A12(S04)a,the stable solid in contact with the solutions is a double salt of the composition, FeS04.A12(S04)3.24H20,corresponding to the alums. With higher concentrations of ferrous sulphate the solutions are in contact with ferrous sulphate heptahydrate, FeS04.iH20,as the stable solid. Occleshaw determined the solid phase by analyzing residues in contact with the several liquid phases and plotting the results on the triangular diagram by the Schreinemakers-Bancroft method. There is a congruent point on the curve corresponding to the double salt. This was confirmed experimentally by crystallizing the double salt from a solution containing ferrous sulphate and aluminum sulphate in equimolecular proportions. The salt so obtained consisted of white needles which matted on the filter to an asbestos-like mass. T h e System: Ferrous Sulphate, Thallous Sulphate, and Water. Benrathj finds that a yellowish-green, soluble salt is formed as a crystalline turbid mass if a t least five times excess of ferrous sulphate be added to a saturated solution of thallous sulphate. The salt has the composition FeSO4.T1S04.6H20 and has been noted previously by Werther.6 Benrath does not furnish figures, but a diagram of the system, from which the data for the transition point can be read approximately. T h e System: Ferrous Sulphate, Cupric Sulphate, and Water. Agde and Barkholt’ made a large number of cooling curves from which they have plotted the isotherms for IO’, 2 jo,30°, 40°, and 56°C. “Krystallographische Chemie”, 1, 434. Z. physik. Chem., 3, j34 (1889). a J. SOC.Chem. Ind., 44, 20 (192jj. J. Chem. S O C . ,127, 2j98 (1925). Z.anorg. allgem. Chem., 151, 23 (1926’ J. prakt. Chem., 92, 134 (1864). 7Z. angew Chem., 39,8 j 1 (1926).
SOLCBILITY O F FERROUS SULPHATE
i 03
Hydrolysis of the salts was found not to be important. At lower concentrations of the liquid phase with respect to ferrous sulphate, the solutions are in contact with cupric sulphate pentahydrate, CuSOd.jHZ0, as stable solid phase. The solubility of the cupric sulphate is, practically, not changed by increasing concentration of ferrous sulphate until a transition point is reached. Beyond the transition point the two salts markedly decrease each others solubility and the solid phases in contact with these solutions are members of a series of solid solutions or isoniorphous mixtures of FeS04.7H20and CUSOI. 7H20 previously dcscribed by Retgers.' The limiting member of this series of solid solutions, a t the transition point, probably has the composition, zFeSOa.7H20.3CuS04.7H~0.Retgers described two series of mixed crystals of the heptahydrates of cupric sulphate and ferrous sulphate; a monoclinic series, ferrous sulphate heptahydrate varying from 47 to roo per cent; and a triclinic series, ferrous sulphate heptahydrate varying from zero to j per cent. The monoclinic double salt, ~FeS04.7H20.3CuS04.~H20, described by Pisani2as a natural mineral occurring in rather large masses as stalactites in a Turkish mine, was not observed. Agde and Barkholt suggest that it is metastable only and changes to the triclinic limiting member of the series of solid solutions. Agde and Barkholt's results are remarkably consistent and it would appear to be easy to duplicate them. Cameron and C r ~ c k f o r d made ,~ a series of solubility determinations a t 3ooC which confirmed the general nature and slopes of the Bgde and Barkholt isotherms, but the location of the transition point is not consistent with the several location found by Agde and Barkholt. Agde and Barkholt found equilibrium to be reached easily and quickly while Cameron and Crockford had exactly opposite experiences. The composition of the solutions at the transition point are assembled in Table 111.
TABLE I11 Composition of Aqueous Solutions of Cupric Sulphate and Ferrous Sulphate a t the Transition Point, CuS04.5H20- (Cu, Fe)S04.iH20. Calculated from the Results of Agde and Barkholt. Temp.
cuso4
percent
FeSOI percent
14.34 17.70
8.82
40'
21.57
1 2 ,I 2
56O
25.34
1.5.64
IO.OC 2 jo
5,57
From Cameron and Crockford's figures the composition of the solution a t the transition point at 3ooC is, approximately I j per cent CuSO, and 12.6 per cent FeSO?. '2. physik. Chem., 15, jjj (1894). * Compt. rena., 48, 80; (18j9). 3 J. Phys. Chem., 33, 709 (1929).
7 04
FRANK K . CAMERON
T h e Quaternary System: Ferrous Sulphate, Cupric Sulphate, Sulphuric A c i d , and Water at 3oOC. Cameron and Crockford’ have charted two isotherms for the system finding that not only do the two salts depress each other’s solubility, but that the presence of sulphuric acid increases the depression in every case. KO transition points were found on either isotherm. The nature of the solid phases was the particular object of the investigation. These were found to be, probably, in every case, cupric sulphate pentahydrate, CuSO4.5H20and a series of solid solutions each member of which contained ferrous sulphate, sulphuric acid, and water. T h e System: Ferrous Sulphate, Sulphuric Acid, and Water. Kenrick? cites Damme? for references to the older literature describing a number of hydrates of ferrous sulphate and double ferrous hydrogen sulphates, some of doubtful validity. Kenrick found that, a t room temperature, ferrous sulphate heptahydrate, FeS04.7H20,is the stable solid phase in contact with solutions up to 43.9 percent sulphuric acid. Tn the present investigation at
TABLE IV Solubility of Ferrous Sulphate in Aqueous Solutions of Sulphuric Acid Series No.
Solution FeSO, per cent per cent
Residue FeSO, per cent per cent
Solid Phase
HzS04
&SO4
Solubility a t o°C I 2
3 4 5 6 7 8 9 IO
I1 I2
13
I4 I5
16 I7
I8 I9
.81 4.10 8.45 15.56 17.76 25.98 26.84 32.50 34.48 36 .os 36.33 37.79 38.62 I
41.80 47 .82 53.25 56.76 60.21 63.60
14.1
3.23
28.91
It
11.10
8.93 7.67 4.80 447 3.99 3.67 3.64 3.68 3.46 3.38 2.34 .91 .55 .46 .37 .28
FeS04.7H20 I1
13.18 10.67
2 7 .OI
11
It
,I I, I, Jl
16.24
33.10
,! I, 13
F~S04.7H20 +FeS04.H20 FeS04.H20 I1
It
25.42
49.50
25.62
40.00
J. Phys. Chem., 33, 709 (1929).
* J. Phys. Chem., 12, 693 ( 1 9 8 ) .
“Handbuch anorg. Chemie”, 3, 329-337 (1893).
I)
,I ,f
SOLUBILITY OF FERROUS SCLPHATE
705
TABLE IV (Continued) Series S o .
Solution
H,SOd
FeSOl
1.13 3.41 6.32 9.37 13 .oo 14.68 17.34 24.54 26.17 27.78
22.88 20.64 18.67 16.79 15.56 14.34 13.2j 11.23
Residue
H2SOd
Solid Phase
FeSOr
Solubility a t 2 5°C I 2
3 4 5 6 7 8 9 IO
I1 I2
13 14 15
16 I7
18 I9 20 21
31 .oo 34.52 35.66 36.17 41.47 45.70 54.71 57.15 60.23 61.92 64,35
FeS04.7H20 I .OI
33.41
f1
>>
7.42
32.51
1
.os IO. 7 0
12
.50
34,6I
,,
17
.OI
48.61
f J f)
1J
18.12
56.21
.97 '76 .56
.55
,, FeS04.7 I b O +FeS04.H20 FeS04.H20
8.50
.75
,, ,, ,,
I1
6.26 5.89 4.99 3.07
JJ
f,
,, ,, J )
29.52
46.98
JJ
tJ
.40 Solubility at 55OC
1 .74 2.42 3.87 5.93 6.45 7.73 11.44
33.48 32.76 31.91 29.20 28.62 26.87 24.34
15.43
20.45
I2
22.26 31.30 38.31 45.37
13
51.02
15.42 9.25 5.39 3.03 I .64
14
56.49 64.03 68.12 69.20
I 2
3 4
5 6 7 8 9 IO I1
1.5
16 =7
FeS04.7H20 Jl
JJ
3 ,04
47.95
)I
f J f J
,, 2,
26.00
22.1
t,
3,
JJ
,, ,,
I .OI
.86 .76 .61
Fe SO4.H20
JJ 77
53 ' 2 9
22.90
,J
j06
FRANK K . CAMERON
TABLE I V (Continued) Series No.
I 2
3 4
5 6
Solution H2S04 FeSO,
.82 .61 3.29 I
34.24
I
34.66 32.57 25.1' 20.48 10.38
10.21
16.32 29.46
Residue
H2S04 Solubility at 65OC
Solid Phase
FeSO4.4Hz0 f )
FeS04.z H 2 0 t) fr
,, Solubilit,y at 7 j"C
I 2
3 4 5 6 7 8
0.43 3.45 5.60 8.71 10.78 21.90
31.46 28.00 25.58 22.60 21.29
27 .72
11 .26
34,72
7 .oj
14.40
2 5 O C the limit has been found to be 45.6 percent. Kenrick found ferrous sulphate tetrahydrate to be metastable below, but near this concentration, which has been confirmed by several observations in the present study. From 43.9 percent to 8 2 . 2 percent sulphuric acid, the stable solid in contact with the solutions is the monohydrate, FeS04.H20,white granular crystals. From 8 2 . 2 to 87.7 percent sulphuric acid in solution the stable solid is white, small, thin hexagons, with the composition, zFeS04.HzSOa. The compound FeS04.H2S04in irregular groups of fine crystals was found to be stable in contact with solutions of 87.7 to 94.1 percent sulphuric acid, and the compound FeS04.3H2S04in fine needles was stable in more concentrated solutions of sulphuric acid. Wirthl quotes Scharitze? as having noted that the presence of sufficient sulphuric acid induces a transformation of the heptahydrate to the monohydrate of ferrous sulphate. Wirth determined the zj°C isotherm for the system finding the transition point a t about a 40 percent or a 1 2 . 2 normal solution of sulphuric acid. His data are in terms of grams ferrous sulphate dissolved by various volumes of aqueous sulphuric acid of stated normalities. Although a direct comparison is not possible, qualitatively, Wirth's results are confirmed by those to be given presently. Cameron and Crockford3 have reported a few figures, which can now be regarded as of qualitative significance only, The system has been investigated by &loser and Hertzner,* who found the various ratios of ferrous sulphate and sulphuric acid which dissolves a definite amount of nitric oxide, or of sulphur dioxide.
'Z. anorg. Chem., 79,364 (1913). 2 Z . Min. Krist., 30, 209 (1899); 37, 549 (1903). J. Phys. Chem., 33, 709 (1929). Z.anal. Chem., 64, 81 (1924).
707
SOLUBILITY O F FERROUS SULPHATE
Solubility curves have been determined at o°C and 25OC by Ethel Ruth N-ard.1 Eight-ounce nursing bottles were employed as containers each solution being about 150 cc in volume in contact with a solid phase of 30 to 50 grams. So large a mass, especially of solid, made the approach to final equilibrium rather slow. But, equilibrium once attained, it was less likely to be disturbed by extraneous influences and it was an advantage to be able to command a large volume of clear solution from which to draw samples without the intervention of filters. Usually, but not always, the sample was of about ten cc. volume and was quickly transferred by an ordinary pipette to a glass-stoppered weighing bottle. The samples were made up to 500 cc. volumes and suitable aliquots taken for analysis. Since the iron must be first removed before the sulphate determination, iron was determined gravimetrically as ferric oxide. To prevent occlusion or adsorption of sulphates, the ferric hydrate was always dissolved with hydrochloric acid and reprecipitated. Although somewhat more labor was involved, this procedure appeared preferable to the usual volumetric estimation with standard potassium permanganate solution particularly as a rather wide range of concentrations was involved in each series. Sulphates were determined as the barium salt, with the usual precautions. The precipitations were made from a relatively large volume to avoid as far as possible adsorption of excess barium. The criteria that equilibria had finally been attained were that successive analyses of the contents a t intervals of a week or more, should give prac-
TABLE V Moles Water of Crystallization found in Solid Phase in Contact with Aqueous Solutions of Ferrous Sulphate and Sulphuric Acid Series KO.
O0
I
7.03
2
3 4 5 6 7 8 IO
74 16 I7 I8 I9 20 1
25'
7.03 6.99
55" 6.96 7.03 6.82
6.99 6.68
6j"
75"
4.25 4.50 2.04 2 .oo
1.73
2 .OI 2
7 .oo 6.99
2.02
.oo 2.16 2.80 2.08 2
I
.98
I .22
I .02
I
.16
2.04
.17 I
.94 .61
1.12
.67
Master's Thesis, Cniversity of North Carolina, (1929).
.40
708
FRANK E. CAMERON
tically identical results, and that these results should, when plotted, fall on smooth curves. The approach to equilibria was very irregular with the individual cases; and, apparently false equilibrium is a common phenomenon with this system. The solutions were generally %ceded" with both the heptahydrate and the monohydrate during the approach to the final state. Composition of the solid phases was established in two ways. I n some cases residues of the solid and adhering mother liquor were analyzed, the data being given in Table IV. The composition of the solid was then found by plotting on the equilateral triangle by the method of SchreinemakersBancroft, or on the isosceles triangle, or by algebraical computations. In other cases, the solid was quickly drained of the mother liquor, washed successively with 98 percent alcohol, then with ether, and dried by pressing between bibulous paper. The results obtained by this procedure are assembled in Table V. At zero degree, there is a transition point, with a “constant solution” containing 38.62 percent HzS04 and 3.38 percent FeS04. Extrapolation gives a value of about 15.3 percent FeSOa in water alone, a figure somewhat low in comparison with Fraenckel’s results but about 12 percent too high in comparison with the data of the International Critical Tables. It would appear that the results here given for o°C may be somewhat high up to about 2 5 percent sulphuric acid. The number and character of the determinations in the vicinity of the transition point leave little doubt that they are substantially correct. I n contact with solutions of less sulphuric acid content than that of the transition point, ferrous sulphate heptahydrate is the solid phase. With more concentrated solutions of sulphuric acid, ferrous sulphate monohydrate is the stable solid. At zj°C the composition of the solution a t the transition point is 2 7 . 7 8 percent H2S04and 10.70 percent. FeS04. Extrapolation shows approxiinately 24 percent FeSOa in water alone, in good agreement with the data of the International Critical Tables. The stable solid phases are, again, the heptahydrate and monohydrate of ferrous sulphate. The writer determined the solubilities a t 55’ and 6 f C . At 55OC, efforts to realize the solution a t the transition point failed. It is certainly near tc solution No. 3. Plotting the results obtained on a large scale and interpolating, the figures 4 percent HzS04 and 31.6 percent FeSOa are obtained Extrapolation gives 34.9 percent FeS04 as the solubility in water alone, in good agreement with the International Critical Tables. The solid phases are the heptahydrate and the monohydrate. At 65’C two curves were realized. The stable solid phase in contact with the more dilute solutions of sulphuric acid was the tetrahydrate, while that in contact with the more concentrated solutions of acid was the dihydrate. By interpolation the concentrations at the transition point were found to be 2 . 5 percent H2S04 and 33.8 percent FeS04. The results obtained a t this temperature were unexpected as it was anticipated that but one curve would be obtained and but one solid phase, the monohydrate. All the solutions had been seeded, generously, with mono-
SOLUBILITY OF FERROUS SULPHATE
709
hydrate &s well as heptahydrate, and there could be no doubt as to the stability of the solid phases as found. The thermometer was then checked. It registered +0.03 on standing in a mush of ice and water, 99.9' in steam with a barometer reading of 759.35 mm Hg and 32.43' in melting sodium sulphate decahydrate.
FIQ.2 Solubility of Ferrous Sulphate in Aqueous Solutions of Sulphuric Acid. I
FlO.
3
Solubility of Ferrous Sulphate in Aqueous Solutions of Other Sulphatea.
I am indebted to Mr. A. T. Clifford for a series run at 75' C. The stable solid phase in contact with solutions from less than a half percent to nearly 35 percent sulphuric acid was found to be the dihydrate. It became necessary then to reinvestigate the solubility in water alone as described under the
FRANK K . CAMERON
710
heading Effect of Temperature. There can be no doubt that the stable solid at 7 5 O C is the dihydrate throughout the whole range of acid solutions from zero to 35 percent. It would appear, however, that Mr. Clifford's determinations of FeS04 a t the lower concentrations of H2S04may be slightly high. Comparison of the E f e c t s of Other Sulphates o n the Solubility of Ferrous Sulphate has been attempted by computing the available data on a common basis of moles per 1000 grams of solution and charting the results. Some of the results are shown in Fig. 3 . It would be needlessly confusing to include all. Temperatures and solid phases are indicated. Beyond the general conclusion that all other sulphates depress the solubility of ferrous sulphate, this method of attack does not appear helpful.
Summary I. Ferrous sulphate is insoluble or very slightly soluble in solvents other than water. 2. The general properties of aqueous solutions of ferrous sulphate have been examined. 3. The conditions necessary to the existence of the several hydrates of ferrous sulphate have been examined. 4. The data regarding the effect of temperature on the solubility of ferrous sulphate in water have been compared. Corrections have been made and new data added. 5. The effects of other sulphates on the solubility of ferrous sulphate have been compared. 6. The solubilities of ferrous sulphate in aqueous solutions of sulphuric acid have been found a t oo, zs0, sso, 65", and 75°C. The composition of the solid phases in contact with the liquid solutions, and the composition of the solutions a t the transition points have been determined.
University of iyorth Carolina, Chapel Hill, North Carolina.
