Journal of
ChemicalVladislava
Technology
and Metallurgy,
52, Petkov,
2, 2017, Peter
333-339
Lachezar
Stamenov,
Stefanova,
Konstantin
Iliev
STUDY OF THE CRYSTALLIZATION PROCESS
OF FERRIC SULFATE HYDRATE FROM RICH OF Fe(III) WASTE SOLUTIONS
Lachezar Stamenov, Vladislava Stefanova, Konstantin Petkov, Peter Iliev
University of Chemical Technology and Metallurgy
8 Kl. Ohridski, 1756 Sofia, Bulgaria
E-mail: [email protected]
Received 14 July 2016
Accepted 12 December 2016
ABSTRACT
In present study the crystallization process of ferric sulfate hydrate - Fe2(SO4)3.xH2O from rich in Fe(III) sulfate
waste solutions was investigated. These solutions were obtained after autoclave oxidation of pyrite concentrate. They
are characterized with high concentrations (> 60 g l-1) of ferric ions and sulfuric acid.
Based on the ternary diagram of the Fe2(SO4)3-H2SO4-H2O system and the laboratory tests the necessary compositions and conditions for preparation of saturated solutions for ferric sulfate crystallization process were determined.
It was found that the crystallization process takes place with obtaining of bulky sludge containing following phases: FeH(SO4)2.4H2O (rhomboclase), Fe2(SO4)3.8H2O (ferric sulfate with eight molecules water)
and Fe4.67(SO4)6(OH)2.20H2O (ferric sulfate hydroxide hydrate). After detention of the sludge for seven days at
temperature 373 K two modifications of ferric sulfate hydrate were observed: ferric sulfate hydroxide hydrate
(Fe4.67(SO4)6(OH)2.8H2O) and paracoquimbite (Fe2(SO4)3.9H2O).
Keywords: waste solutions, crystallization process, ferric sulfate hydrate.
INTRODUCTION
Ferric sulfate is primary used in fresh/drinking water
and wastewater treatment. It is also used in pigments,
pickling baths for aluminum and steel. Ferric sulfate
is effective primary coagulant based on trivalent iron
(Fe3+), and is excellent for drinking water production,
wastewater treatment, phosphorus removal applications,
struvite control as well as sludge conditioning [1]. It is
also efficient in preventing odor and corrosion by controlling the formation of hydrogen sulfide.
Ferric sulfate can be produced in liquid or in solid
form. In patent US7387770 B2 [2] ferric sulfate solution
was obtained from finely-divided ferric oxide, sulfuric
acid and water in closed reaction vessel at temperatures
ranging from about 392 K to about 422 K and pressures
from about 172.4 kPa to about 482.6 kPa. It is usually
obtained from ferrous sulfate by direct oxidation using
strong oxidant such as H2O2 [3], KClO3 [4], NaClO [5],
HNO3 [6] or by catalytic oxidation of ferrous sulfate using NaNO2 or NaI as a catalyst in acid media. In Patent
US 5766566 A [6] crystals of ferric sulfate are produced
by solidification of ferric sulfate slurry by cooling the
solution. Also known are processes according to which
the ferric sulfate solution is solidified by granulation or
by other solidification methods, which always involve
evaporation of the water.
In this work we studied the crystallization process
of ferric sulfate hydrate from rich of Fe(III) hydrometallurgical waste solutions. Such solutions are obtained
from autoclave oxidation of pyrite concentrate. A special
feature of these solutions is the high concentration of
Fe(III) ions (> 60 g l-1) and sulfuric acid (> 61 g l-1). From
a practical point of view, there is an area for utilization of
ferric (III) ions contained in those solutions - for example
for crystallization of ferric sulfate hydrate.
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Journal of Chemical Technology and Metallurgy, 52, 2, 2017
Table 1. Chemical composition of the synthetic solution.
Fetotal,
g
l-1
60.98
Fe(II),
Fe(III),
g l-1
g l-1
1.34
59.64
The most commonly applied methods for removal
of Fe(III) from waste process solutions are iron hydroxide precipitation [7 - 9] and precipitation of jarosite MeFe3(SO4)2(OH)6, where Me = Na, K, OH [10 - 14].
In both cases, the generated sludge is a waste product
and it is deposited in a special stores. In addition to
these methods “Geothite” [10, 15 - 17] and “Hematite“
[10, 17 - 19] processes are also applied for removal of
ferric ions from waste solutions. Both processes allows
purification of the solutions from iron (Fe < 1 g l-1 ) .
There is scarce number of studies in the literature
concerning the field of utilization of waste solutions for
obtaining ferric sulfate hydrate. A larger part of studies
have aimed the production of polyferric sulphate (PFS)
[Fe2(OH)n(SO4)3-n/2]m (n < 2, m > 10) which is used as
pre-polymerized inorganic coagulant with high cationic
charge [20].
Jiang J.-Q. et al. [21] have studied two methods
for preparation of PFS : sodium nitrite catalytic oxidation of waste ferrous sulfate solutions and industrial
wastes - TiO2 by-products (20 % H2SO4, 16 % Fe2+ and
23 % total iron ). The authors found that PFS obtained
from industrial wastes could not be applied in the
treatment of drinking water because of the low purity
of these by-products, but it could be used in the treatment of wastewater. The authors have investigated the
coagulation properties of the resulting PFS and did not
give a characterization of the obtained compound.
Zhongguo L. et al. [22] have studied a method
for disposing sodium-jarosite which can be used to
synthesize PFS. The method consists of two-step leaching experimental procedure. First step is pre-leaching
Cu,
g
H2SO4,
l-1
1.30
g l-1
61.75
process for removal of impurity metals and second step
- decomposition of sodium-jarosite to provide enough
ferric ions for synthesizing PFS. The second step was
carried out at temperature below 332 K and sulfuric
acid consumption 0.8 ml g-1 sodium-jarosite. The results
showed that the PFS synthesized from sodium-jarosite
had a poly-iron complex Fe4.67(SO4)6(OH)2.20H2O.
EXPERIMENTS
Materials and methods
Initially the experiments were conducted with
synthetic solution. Chemically pure reagents:
Fe2(SO4)3.9H2O, FeSO4.7H2O, CuSO4.5H2O, concentrated H2SO4 (98 %, d = 1.84 g cm-1) and distilled water
were used for the preparation. The chemical composition
of the syntetic solution is given in Table 1. It is characterized by a high concentration of Fe(III) ions and sulfuric
acid and is similar in composition to the waste solutions
obtained after autoclave pressure oxidation (POX) of
pyrite concentrate. The chemical composition of the
industrial solution is given in Table 2.
The concentration of Fe(III) and Fe(II) ions in the
industrial solution is respectively 58,96 g l-1 and 1,34 g
l-1. From all impurities in the solution the copper ions are
with the highest concentration (1.47 g l-1). The content
of the other impurities is comparatively low and should
not affect the crystallization process.
The evaporation was conducted in a water bath without stirring to avoid the formation of poorly soluble basic
sulfates. The crystallization process begins only once 70
% of the water from outgoing solution was evaporated.
Table 2. Chemical composition of the solution obtained after autoclave total pressure oxidation of pyrite concentrate.
Fetotal,
g l-1
60.3
334
H2SO4, Cu,
g l-1
g l-1
61.5 1.47
Na,
mg l-1
100
K,
mg l-1
100
Mg,
mg l-1
<10
Mn,
mg l-1
16.16
Ni,
mg l-1
76.64
Pb,
mg l-1
40.22
Al,
mg l-1
900
Ca,
mg l-1
200
Cr,
mg l-1
25.99
Lachezar Stamenov, Vladislava Stefanova, Konstantin Petkov, Peter Iliev
Table 3. Effect of the evaporated water amount on the crystallization process.
Test ID
Evaporation
of water, %
F1
F2
F3
F4
F5
F6
F7
F8
F9
0
50
50
60
60
70
70
73
75
Density of
solution
g cm -3
1226
1520
1520
1647
1652
1788
1788
1874
1925
Mass composition of solution before
crystallization in %
Fe2(SO4)3
H2SO4
H 2O
17.76
5.04
77.20
28.00
6.56
65.44
27.25
9.72
63.03
31.44
8.26
60.30
28.95
10.38
60.67
38.08
9.18
52.74
37.86
5.81
56.33
42.95
2.44
54.60
42.95
2.44
54.60
The crystals obtained after the crystallization process were separated from the mother liquor by filtration with a Buchner funnel and BOECO R300 vacuum
pump. The crystals were washed, dried and weighed
with analytical scales.
The complexometric method was used for determination of the concentration of Fetotal and Cu2+ ions
in the mother liquor. The concentration of Fe(II) was
determined by the bichromate method and the concentration of Fe(III) ions was determined by the difference in
concentrations between Fetotal and Fe(II) ions. For evaluation of the impurities content the ААА and ICP-OES
Crystals
mass
g
5.78
15.40
48.80
48.80
analyzes were used. This involves the use of PERKINELMER 5000 atomic-absorption spectrophotometer and
Jeledyne Leeman Lab devices.
The phase composition of the crystals was determined by means of X-ray diffraction (XRD). For this
purpose PANalytical Empyrean device, equipped with
a multi-channel detector (Pixel 3D) using the (Cu-K𝛼
45 kV-40mA) radiation in the range of 2𝜃 between 20
– 115o, with a step of scanning of 0.01o for 20 s is used.
The morphology of the resulting crystals was determined microscopically. The stereo microscope Leica
80M with digital color camera was used.
Fig. 1. Ternary diagram of the system Fe2(SO4)3–H2SO4–H2O [23].
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Journal of Chemical Technology and Metallurgy, 52, 2, 2017
The concentration of free sulfuric acid was determined by titration with 0.1 N NaOH solution in the
presence of an indicator methyl orange after oxidation
of Fe(II) ions to Fe(III) ions with hydrogen peroxide.
RESULTS AND DISCUSSION
In our study we have used a ternary diagram of the
system Fe2(SO4)3-H2SO4-H2O [23] built at 297 K, as a
starting point for the study of the process of crystallization of ferric sulfate hydrate.
The initial composition (by mass) of the synthetic
solution was 17.76 % Fe2(SO4)3, 5.04 % H2SO4 and
77.2 % H2O. It is very clearly seen from Fig. 1 that at
this composition the crystallization process could not
be performed.
In order to fall into the area of compositions where
crystallization process is possible, we undertook a series
of experiments with evaporation of different amount of
water. The composition of the solutions after evaporation
is shown in Table 3 and in Fig. 2.
The analysis of the results shows that the process
of crystallization starts after evaporation of 70 % water,
i.e., at a density of the solutions over 1788 g l-1. In the
experiments F8 and F9 all mother liquor is converted to
sludge which is extremely undesirable because all the
iron present in the solution crystallized along with all
solution impurities.
The appearance of the resulting crystals can be seen
on Fig.2. The photographs were made with a stereo
microscope Leica 80M equipped with a color digital
camera. On the same figure the appearance of the sample, taken from commercial product - ferric sulphate,
is shown.
To carry out the process of internal crystallization,
the crystalline precipitate is maintained for 7 days in an
ampoule at 372.15 K.
In Figs. 3a and 3b diffraction patterns of the crystals
obtained from experiment F6 and F9 are presented. They
were obtained after 70 % and 75 % evaporation of the
water from the solution.
The main phases identified in the sample F6 were:
rhomboclase - FeH(SO4)2.4H2O, ferric sulfate with eight
molecules water - Fe2(SO4)3.8H2O and poly ferric sulfate
- Fe4.67(SO4)6(OH)2.20H2O.
The diffraction patterns in the sample obtained after
experiment F9 before heating are: Fe2(OH)2(SO4)2.7H2O,
336
Fe2(SO4)3.H2SO4.2H2O, Fe2(SO4)3.H2SO4.8H2O and
Fe2(SO4)3.9H2O.
It is obvious that the resulting sludge is a mixture of
crystalline phases that do not match a commercial product Fe2(SO4)3.9H2O and require conducting of internal
recrystallization.
To determine the conditions of the recrystallization
process we performed two parallel experiments with
synthetic solution and with solution obtained during
autoclave oxidation of pyrite concentrate. After evapora-
a)
b)
c)
Fig. 2. Appearance of the crystals obtained after evaporation of: а) 70 % ; b) 73 % and c) 75 % of water.
Lachezar Stamenov, Vladislava Stefanova, Konstantin Petkov, Peter Iliev
Table 4. Content of impurities in the crystals obtained after recrystallization.
Cu,
%
Na,
%
0.03 0.02
K,
%
0.01
Mg,
%
Mn,
%
Ni,
%
Pb,
%
0.01 0.0041 0.00096 0.001
tion of 70 % of the volume of the solution, it is placed in
a bottle weight and left for 2 days at a temperature 389.15
K. After cooling, the resulting crystalline precipitates
were weighed on an analytical balance, analyzed for the
content of iron, sulfate sulfur and impurities. The results
a)
Al,
%
Ca,
%
Cr,
%
0.2
0.06 0.0003
obtained are summarized in Table 4.
The content of impurities in the precipitate obtained
from the industrial solution is minimal in respect of heavy
non-ferrous metals. From this it can be concluded that
it may be used as a coagulant for wastewater treatment.
The appearance of the crystals obtained from industrial solution after recrystallization at 388.15 K and
crystals of commercial ferric sulfate are compared on
Fig. 4. It is seen that the obtained crystals are bigger than
the commercial, but with the same pale purple color.
The diffraction pattern of the crystals obtained from
industrial solution is presented in Fig. 5. It can be seen
that the main phases after recrystallization at a high temperature are Fe2(SO4)3.H2SO4.8H2O and Fe2(SO4)3.9H2O.
a)
b)
b)
Fig. 3. Diffractogram of the crystals obtained after experiments F6 and F9.
Fig. 4. Appearance of the crystals: a) obtained after recrystallization; b) crystals of commercial ferric sulfate.
337
Journal of Chemical Technology and Metallurgy, 52, 2, 2017
Fig. 5. Diffractogram of crystals obtained from the industrial solution after recrystallization.
CONCLUSIONS
Based on the experimental studies the possibility
of crystallization of ferric sulphate (Fe2(SO4)3) from
solution obtained upon pressure oxidation of pyrite concentrate with high concentration of ferric ions (over 60 g
l-1) has been demonstrated. The following conditions for
conducting of both the processes of crystallization: 70
% evaporation of the water, till reaching density of the
solution and of recrystallization: temperature 388.15 K,
absence of air, process duration 2 days, are determined.
The obtained precipitate under these conditions is coarse
crystalline, with minimal impurities content and is suitable as a coagulant for industrial wastewaters treatment.
REFERENCES
1. http://www.kemira.com/en/industries-applications/
pages/ferric-sulfate.aspx
2. Patent US7387770 B2, Method for ferric sulfate manufacturing, M. Wilkinson, J. Hurd, D. Stone, 2005.
3. Patent US4707349 A, Process of preparing a preferred
ferric sulfate solution and product, B. Norman, 1987.
4. Patent CN 103626274 A, Method for preparing floc-
338
culant agent, 王传虎, 秦英月, 江彬彬, 会路, 董传
明, 王孝虎, 王志远, 2014.
5. Patent CN 104402145 A, Production method for preparing ferric hydroxide by utilizing ferrite-containing
waste water, 郭必裕, 张胜超, 丁宇航, 佟正宇, 蔺鸿
珺, 赵爱婧, 2015.
6. Patent US 5766566 A, H. Mattila, T. Kenakkala, O.
Konstari, Process for preparing ferric sulfate, 1998.
7. N.J. Bolin, J.E. Sundkvist, Two stage precipitation
process of iron and arsenic from acidic leaching solutions, Transactions of Nonferrous Metals Society of
China, 18, 6, 2008, 1513-1517.
8. M. Balintova, A. Petrilakova, Study of pH Influence
on Selective Precipitation of Heavy Metals from Acid
Mine Drainage, Chemical Engineering Transactions,
25, 2011, 345-350.
9. G. Biedermann, P. Schindler, On the Solubility
Product of Precipitated Iron (III) Hydroxide, ACTA
Chemica Scandinavica, 11, 1957, 731-740.
10. C.K. Gupta, T.K. Mukherjee, Hydrometallurgy in
Extraction processes, Bombay-India, 1990.
11. J. Babcan, Synthesis of jarosite, Geol. Zb., 22, 2,
1971, 299-304.
12. M. Danovska, M. Golomeova, D. Karanfilov, A. Zendel-
Lachezar Stamenov, Vladislava Stefanova, Konstantin Petkov, Peter Iliev
ska, Treatment of Fe(III) ions from leaching solution
with neutralization and precipitation, 5-th Balkan
Mining Congress, Ohrid, Macedonia, 2013, p.533-538.
13. V.B. Sefton, G.M. Swinkels, C.R. Kirby, Process
for the precipitation of iron as jarosite, 1970, Pat.
4193970.
14. G.A. Norton, R.G. Richardson, Precipitation of jarosite compounds as a method for removing impurities from acidic wastes from chemical coal cleaning,
Environ. Sci. Technol., 25, 3, 1991, 449-455.
15. P.T. Davey, T.R. Scott, Removal of iron from leach
liquors by the “Goethite” process. Hydrometallurgy,
2, 1,1976, 25-33.
16. R.S. Stahl, D.S. Fanning, B.R. James, Goethite and
Jarosite Precipitation from Ferrous Sulfate Solutions, ACSESS DL, 57 , 1,1993, 280-282.
17. R.A. Silva, C.D. Castro, C.O. Petter, Production of
Iron Pigments (Goethite and Haematite) from Acid
Mine Drainage. Mine Water - Managing the Challenges, Aachen Germany, 2011, p. 469-474.
18. M. Rodriguez, B.J. Wedderburn, Hematite precipitation, Pat. EP 1971697 B1, 2012.
19. S.J. Hock, Precipitation of Hematite and Recovery
of Hydrochloric Acid from Aqueous Iron (II, III)
Chloride Solutions by Hydrothermal Processing,
Abstract of Thesis, McGill University, Montreal,
Canada, 2009.
20. A. Zouboulis, P. Moussas, Polyferric sulfate: Preparation, characterization and application in coagulation experiments, Journal of Hazardous Materials,
155, 3, 2008, 459-468.
21. Y. He, F. Li, J.-Q. Jiang , H.Wang, Preparation and
application of polyferric sulfate in drinking water
treatment, 12th International conference on Environmental Science and Technology, Rhodes, Greece,
2011, 714-720.
22. L. Zhongguo, Y. Wenyi, Industrial waste utilization
method: producing poly-ferric sulfate (PFS) from
sodium-jarosite residue, Front. Environmental Science Engineering, 9, 4, 2015, 731-737.
23. P. Kobylin, T. Kaskiala, J. Salminen, Modeling of
H2SO4-FeSO4-H2O and H2SO4-Fe2(SO4)3-H2O Systems for Metallurgical Aplications, Ind. Eng. Chem.
Res., 46, 2007, 2601-2608.
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