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Materials Research Bulletin, Vol. 34, No. 6, pp. 905–914, 1999
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UNIFORM PARTICLES WITH A LARGE SURFACE AREA FORMED BY
HYDROLYSIS OF Fe2(SO4)3 WITH UREA
Jan S̆ubrt, Jaroslav Bohác̆ek, Václav S̆tengl, Tomás̆ Grygar*, and Petr Bezdic̆ka
Institute of Inorganic Chemistry, Academy of Sciences of the Czech Republic,
250 68 R̆ez̆, Czech Republic
(Refereed)
(Received May 20, 1998; August 12, 1998)
ABSTRACT
Synthesis of poorly crystalline iron(III) hydrous oxides and basic salts (ferrihydrite, schwertmannite, and jarosite) consisting of uniform spherical particles of characteristic shape with a diameter of ⬃10 ␮m is described. The
preparation procedure is based on homogeneous precipitation of aqueous
solutions of Fe2(SO4)3 with urea in the temperature range 60 –100°C. Thermal
dehydration of iron(III) hydroxide leads to an amorphous ferric oxide with a
high surface area (⬃200 m2/g) that is stable up to about 500°C. © 1999 Elsevier
Science Ltd
KEYWORDS: A. amorphous materials, A. microporous materials, C. chemical synthesis
INTRODUCTION
Poorly crystalline precipitates formed by hydrolysis of ferric salts using various alkaline
reagents are usually called amorphous ferric oxide– hydroxide or ferrihydrite [1,2]. Recently,
a similar sulfate-containing solid phase called schwertmannite was identified [3–7]. Synthetic
precipitates of ferrihydrite are often used as sorbents [8 –12] as well as starting materials for
the synthesis of catalysts, pigments, and magnetic recording materials [2,12–15]. Schwertmannite and the products of its transformation have also been proposed as prospective
materials for practical utilization [7].
*To whom correspondence should be addressed.
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Vol. 34, No. 6
Schwertmannite can be prepared by precipitation of Fe3⫹ or oxidative precipitation of
Fe in the presence of sulfates and at pH ⬃2 [3–7]. Ferrihydrite is conventionally prepared
by the neutralization of aqueous solutions of ferric salts with ammonium or alkali metal
hydroxides [2,6,15–19]. As-formed precipitates usually consist of highly irregular aggregates
of very small particles (mostly 4 – 6 nm) [1,2]. Schwertmannite particles are usually spherical
aggregates of a characteristic shape resembling a pincushion with a diameter well below 1
␮m [6].
The use of urea as a neutralizing agent for the homogeneous hydrolysis of aqueous
solutions of some metal salts has already been reported [20 –27]. According to these results,
the rate of hydrolysis of metal ions with urea is significantly lower than that with NH4OH or
alkali metal hydroxides. The lower hydrolysis rate entailed a better control of the precipitation conditions resulting in a higher reproducibility of the properties of the precipitates.
Surprisingly, the particles of ferric oxides obtained by the homogeneous precipitation were
reported to be of uniform shape and size [26,27]. Moreover, urea has been reported to
stabilize ferric hydrosols [28].
We hereafter describe the products of the precipitation of ferric salts with urea. The
formation of the recently described solid phase, schwertmannite, of very unusual particle size
is reported. All the products obtained by hydrolysis of ferric sulfate with urea possess a
remarkable morphology, which facilitates their filtration and drying and also enables their
transformation to amorphous powders with a large specific surface area by heating at
200 – 400°C.
2⫹
EXPERIMENTAL
Hydrolysis Methods
Method A. A mass of 250 g Fe2(SO4)3•5H2O was dissolved in 2 L of distilled water in a
5 L beaker. The solution was brought to boil and then a solution of 500 g of urea in 700 ml
of distilled water was added dropwise for 2– 4 h. After adding approximately 3/4 of the urea
solution, the color of the reaction mixture turned to orange. At this stage the reaction mixture
started to effervesce as a result of the CO2 evolution. After discoloration of the reaction
mixture, traces of NH3 escaped from the boiling solution. Then an additional 100 g of urea
in 200 ml of distilled water was added and the mixture was boiled for 5 h. The changes of
solution pH are given in Table 1.
Method B. The solution of ferric sulfate with urea, at different final ratios urea/Fe, was
stirred at some selected temperature between 60 and 90°C (see Table 2). The 5 M urea
solution was added at the rate of 2 L per hour to 2 L of Fe2(SO4)3 solution preheated to 60
and 95°C. The actual concentration of Fe3⫹ is given in Table 2. Due to a slow decomposition
of urea, pH of the reacting mixture continuously increased. Sample B7 was prepared in a
similar way but from ferrous sulfate. Unlike method A, method B allows truly homogeneous
precipitation without local oversaturation and vigorous decomposition of urea.
Method C. A volume of 0.5 L of 0.183 M solution of ferric sulfate was preheated to 80°C.
Then a solution of hexamethylene tetraamine (HMTA) was added dropwise within a half
hour. The suspension was stirred for another 8 h.
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TABLE 1
Properties of the Intermediates and Product of Hydrolysis Using Method A
(Boiling Solution)
Sample
A2
A3
A4
A5
A6
Time (min)
pH
Color and phase compositiona
␴(m2/g)b
115
170
220
350
500
2.8
3.1
3.3
5.6
5.7
light ochre, Jt
orange, Jt
red
dark violet, Fh ⬎⬎ Gt
dark violet, Fh ⬎⬎ Gt
20
5
1
280
20
a
Phase composition: Jt: jarosite; Fh: ferrihydrite; Gt: goethite. Particle appearance is shown
in Figure 1.
b
Surface area obtained after heating the sample at 180°C for 1 h.
Characterization Methods. The samples were characterized using powder X-ray
diffraction with a Siemens D-5005 diffractometer (Cu K␣ radiation, a secondary monochromator). Transmission electron microscopic (TEM) photographs were obtained with
a Philips 201 electron microscope. Scanning electron microscopic (SEM) photographs
were obtained using a TESLA BS-350 electron microscope coupled with a Philips EDAX
analyzer. Selected samples were annealed in a laboratory furnace in air for 1 h at a
constant temperature of 200, 300, 400, or 500°C, and their specific surface area was
determined using the BET method with a Coulter SA 3100 device. Chemical analyses
were performed using chelatometry (total Fe), gravimetry (sulfates), and distillation
followed by titration (NH4⫹).
TABLE 2
Samples Obtained by Method B
Sample
B11
B3
B8
B6
B2
B7c
B5
B1
B9
B10
B4
a
Urea/Fe
Fe, M
pHa
t, °C
Intermediatesb
Final productb
1.7
3.38
3.38
3.38
6.75
6.75
6.75
13.5
13.5
27
27
0.183
0.092
0.183
0.183
0.183
0.183
0.183
0.183
0.366
0.183
0.183
1.9–2.6
1.8–5.2
2.1–4.5
1.9–5.0
1.8–5.0
2.6–4.4
1.9–5.5
2.0–5.9
2.7–5.2
3.3–3.3
2.9–6.1
95
80
80
95
80
80
95
80
80
60
80
Sh, Jt
Sh, Jt
Sh
Sh
Sh ⬎ Jt
Gt, Sh
Sh ⬎ Gt
Sh (high sulfate)
Sh
Sh (high sulfate)
Sh
Jt
Jt (ammonium–hydronium)
Sh ⬎ Gt
Jt, Sh
Sh ⬎ Jt (ammonium)
Gt
Sh ⬎ Gt
Sh (low sulfate) ⬎⬎ Gt
Sh
Sh (high sulfate)
Sh (low sulfate)
pH when a significant amount of solid was formed; final pH.
Phase composition: Sh: schwertmannite; Jt: jarosite; Gt: goethite; Hm: hematite.
c
Prepared from ferrous sulfate (in place of ferric sulfate).
b
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FIG. 1
Time dependence of pH during hydrolysis of 0.25 M Fe2(SO4)3 (hydrolytic method A).
Points labeling for intermediate sampling: (1) start of urea addition (pH ⫽ 1.2), (2) the first
appearance of a precipitate in the solution, (3) end of urea addition, (4) beginning of CO2
evolution, (5) full discoloration of the solution, and (6) end of the reaction.
RESULTS AND DISCUSSION
In agreement with literature results [1,13–20,29], the voluminous precipitates of synthetic
ferrihydrite samples prepared by rapid precipitation from aqueous solutions of ferric salts by
alkali metal hydroxides or ammonia, mostly consist of very small particles (⬃4 – 6 nm). The
filtration or dialysis and washing of the obtained ferric gel, as well as grinding of the dried
product, is rather difficult. In contrast, the properties of samples of the ferric gel prepared by
hydrolysis with urea were surprisingly different in many aspects. The sedimentation of the
precipitates in the aqueous media was very rapid, their filtration and washing was easy, and
after drying, they were obtained as loose powder, which did not tend to sinter during drying
and annealing.
A typical curve of the time dependence of pH at 100°C for the hydrolytic method A is
given in Figure 1. In the first stage between the beginning of the experiment and point 1,
before adding urea, a small decrease of pH was observed, due to the partial hydrolysis of the
ferric salt in the boiling solution. At point 1, adding the urea solution at a constant rate was
initiated. At pH ⫽ 2.8, formation of the first fractions of light-ochre jarosite precipitate was
observed. During further addition of the urea solution, the color of the precipitate turned dark
red and its amount increased rapidly. At point 4, formation of gaseous CO2 from urea started
and caused the solution to effervesce vigorously. After reaching this point, the amount of
precipitate increased rapidly and its color turned dark violet. At point 5, the Fe3⫹ ions
Vol. 34, No. 6
HYDRATED IRON OXIDE
909
FIG. 2
SEM micrographs of the final hydrolytic products: A6 (mainly ferrihydrite); B4 (schwertmannite); B9 (schwertmannite); B3/1, the early stage of B3 (jarosite formed from schwertmannite); B3 (jarosite); and B7 (goethite, prepared from ferrous sulphate).
disappeared entirely from the solution. Important properties of these samples are given in
Table 1. A typical example of the particles obtained (A6) is shown in Figure 2.
The morphology, appearance, and color of the products obtained using method B were
very similar to those of method A. It is remarkable that the overall spherical shape of
aggregates was preserved even when the phase composition of the solid phase varied
substantially (Fig. 2). The aggregates were made of a variety of primary particles, such as
submicrometer needles or irregularly shaped particles, and from much bigger, rounded plates.
Chemical analysis and powder X-ray diffraction (Fig. 3) showed that at pH ⬃ 2 and for
a low urea/Fe ratio, jarosites are the preferred products. This is in line with the thermody⫹
⫹
namics of the Fe-SO2⫺
4 –H –NH4 system. In contrast to schwertmannite, jarosites are
crystalline, as shown by XRD. According to chemical analysis, the jarosites were usually
deficient in NH4⫹. The XRD patterns (Fig. 2) match best that of ammonium jarosite (JCPDS
26-1014), but they are also similar to that of hydronium jarosite (JCPDS 31-0650). In
addition, the diffraction lines (021) and (113) are separated, yielding a characteristic wellresolved doublet at 2␪ between 28° and 29°. Therefore, the chemical composition and the
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J. S̆UBRT et al.
Vol. 34, No. 6
FIG. 3
XRD patterns. From top to bottom: A6 (ferrihydrite with traces of goethite), B9 (schwertmannite), B8 (schwertmannite with traces of goethite), and B3 (jarosite).
XRD patterns suggest that the formed jarosite denoted as Jt (ammonium– hydronium) in
Table 2 contained about 70 –90% NH4⫹ and 30 –10% H3O⫹ of monovalent cations, i.e., the
jarosite formula can be written as (NH4)1⫺x(H3O⫹)xFe3(SO4)2(OH)6 with 0.7 ⬍ x ⬍ 0.9.
At pH ⬃ 3 and for a large urea/Fe ratio, schwertmannite is preferred. The above-mentioned
pH values typical for the formation of jarosite and schwertmannite are in line with those
given in the literature [5,7]. However, the actual phase composition of the solids depends not
only on thermodynamics, but also on kinetic factors. Schwertmannite is thermodynamically
unstable with respect to jarosite (at lower pH) or to goethite (at higher pH) [6], but its
formation from soluble ferric oligomers requires less structural rearrangements [7]. Therefore, the schwertmannite-promoting effect of urea could possibly be related to its stabilizing
effect on ferric sols [28].
The schwertmannite formula has been reported as Fe16O16(OH)16⫺2z(SO4)z with Fe/S
between 4.6 and 8 according to earlier reports [4,6], or Fe4O4(OH)2An⫺2/n where A is an
appropriate anion [7]. In line with its variable composition, the observed Fe/S ratio of our
Vol. 34, No. 6
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TABLE 3
Samples Obtained by Method C
Sample
B14
B15
B16
B17
HMTA/Fe
pH
t, °C
Intermediatesa
Final producta
4
0.8
0.4
0.13
5–7
2
2
1–2
80
80
80
80
Sh
Sh
Jt ⬎ Gt, Hm
Jt ⬎ Sh
Sh (low sulfate)
Sh (low sulfate) ⬎ Gt
Jt ⬎ Gt, Hm
Jt (ammonium–hydronium)
a
Sh; schwertamannite; Jt: jarosite; Gt: goethite; Hm: hematite.
samples of schwertmannite varied from 5 to 20 (see Table 2 and the denotation “high sulfate”
Sh and “low sulfate” Sh, respectively). The low-sulfate schwertmannite was usually obtained
at higher temperatures and at final pH ⬃ 5. According to Barham [7], Fe/S ratios as high as
25 to 50 could be reached by anionic exchange at pH ⬃ 7. As follows from the composition
of sample A6, after heating for 5 h at pH ⬃ 6 and at a temperature of 100°C, the majority
of the sulfates was indeed removed from the solid, but the schwertmannite structure collapsed
(compare corresponding XRD patterns in Fig. 3).
When ferric sulfate was substituted by ferrous sulfate, the primary particles also aggregated to the characteristic spherical particles. The first subsample contained schwertmannite,
but it recrystallized to goethite, preserving the original spherical shape (see sample B7 in Fig.
2). When ferric sulfate was substituted by ferric nitrate, goethite and haematite were formed.
This result is in line with previous reports [26,27]. The phases were of poor crystallinity
(XRD) and the primary particles did not aggregate into large spheres.
The precipitation of ferric sulfate with HMTA using method C yielded ammonium–
hydronium jarosite at pH ⬃ 2 and schwertmannite at pH ⬃ 5 (see Table 3), but the primary
particles did not aggregate into large spheres at all. This comparison clearly demonstrates the
indispensable role of urea that promotes the adhesiveness of the primary particles.
We have also observed that when a suitable support is present in the solution during
hydrolysis, all the precipitate adheres on the surface. This phenomenon can be utilized to
cover various surfaces with layers of metal hydroxides, e.g., to obtain colored layers on
appropriate solid particles [30]. This process was used for coating mica, with the aim of
substituting specularite in corrosion-protective pigments. The reaction mechanism of coating
is certainly based on the unique properties of primary particles of schwertmannite.
The thermal dehydration of ferric gels takes place at ⬃250°C and leads to crystalline
␣-Fe2O3 (haematite) [15]. The resulting haematite particles are irregular, with a broad
particle size distribution.
When heated in air, the samples of ferrihydrite and schwertmannite dehydrated in a
relatively broad temperature interval between 50 and 400°C. The presence of a significant
amount of pores in the particles is responsible for their high surface area [31]. The sample
remained amorphous to X-ray until ⬃450°C. The specific surface areas of samples heated for
1 h at given temperatures are plotted in Figure 4.
CONCLUSIONS
We found that slow hydrolysis of aqueous solutions of Fe2(SO4)3 with urea in the temperature range 60 –100°C leads to a characteristic form of iron(III) hydrous oxides and basic
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J. S̆UBRT et al.
Vol. 34, No. 6
FIG. 4
Temperature dependence of specific surface areas of the hydrolytic products after annealing
at given temperatures for 1 h.
sulfates. The solids consisted of uniform spherical aggregates with a diameter in the order of
tens of ␮m (see Fig. 3). Basic iron(III) sulfates, with a structure and chemical composition
corresponding to ammonium– hydronium jarosite, were formed at low pH and low urea/Fe
ratio, whereas generally metastable schwertmannite was obtained at high concentration of
urea and at 3 ⬍ pH ⬍ 6. Figure 5 shows the generalized scheme of the observed transformations of ferric oxides obtained when ferric sulfate is hydrolyzed with urea. Hydrolysis of
aqueous solutions of Fe(NO3)3 and FeCl3 with urea, or Fe2(SO4)3 with hexamethylene
tetraamine, led to solids of significantly different particle morphology under otherwise
similar reaction conditions.
FIG. 5
Generalized scheme of hydrolysis of Fe2(SO4)3 by urea. Ferrihydrite is the solid phase
suitable for the preparation of amorphous ferric oxide with large specific surface area,
schwertmannite is responsible for the unique particle morphology. Legend: Sh: schwertmannite; Jt: jarosite; Fh: ferrihydrite; Gt: goethite; Hm: hematite.
Vol. 34, No. 6
HYDRATED IRON OXIDE
913
The unique adhesiveness of schwertmannite primary particles suggests a possible application of the hydrolysis of ferric sulfate with urea for easy production of films of Fe oxides
on suitable solid surfaces. The production of pigments comparable to specularite (␣-Fe2O3
with platy crystals) [30] may serve as an example.
Thermal dehydration of these spherical aggregates of iron(III) hydrous oxides yielded
amorphous Fe2O3 with high specific surface area (about 200 m2/g) that is stable up to
450 –500°C. This feature allows us to propose this product of homogeneous hydrolysis of
ferric sulfates as a possible sorption material. The high specific surface area and uniform
spherical particles could also make the hydrolytic products suitable for production of
catalysts and ferrites.
ACKNOWLEDGMENT
This work was supported by the Ministry of Industry and Trade of the Czech Republic, grant
no. PZ-CH/06/98.
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