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Materials Research Bulletin, Vol. 34, No. 6, pp. 905–914, 1999 Copyright © 1999 Elsevier Science Ltd Printed in the USA. All rights reserved 0025-5408/99/$–see front matter PII S0025-5408(99)00078-1 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. 905 906 J. S̆UBRT et al. 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. Vol. 34, No. 6 HYDRATED IRON OXIDE 907 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 908 J. S̆UBRT et al. Vol. 34, No. 6 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 910 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 HYDRATED IRON OXIDE 911 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 912 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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