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.