Effect of Solution Silicate on the Precipitation of Barium Sulfate

Article pubs.acs.org/crystal

Effect of Solution Silicate on the Precipitation of Barium Sulfate Franca Jones,* Tomoko Radomirovic, and Mark I. Ogden Department of Chemistry, Curtin University, GPO Box U1987, Perth, Western Australia 6845 S Supporting Information *

ABSTRACT: The presence of silicate during barium sulfate crystallization has different impacts depending on the pH of the solution. At pH 7 the dominance of the protonated form (H4SiO4) and possible polymerization of the silicate impacts mainly on the aggregation state and on twinning of the barium sulfate formed. At higher pH values (∼10), the silicate ion present is able to influence both morphology and partially substitute for sulfate in the lattice. Interesting fibrous particles are formed under these conditions, but this is not due to mesocrystal formation as the particles are observed to be single crystalline in nature. These fibrous sections are found to be dominant on the surface and are highly porous. These particles are different, however, to the biomorphs formed when crystallization of barium carbonate occurs in the presence of silicate. This is because the speciation of sulfate does not change over a large pH range. The impact of silicate on barium sulfate particles is similar to the impact on calcium carbonate and strontium sulfate crystallization.

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INTRODUCTION Barium sulfate, an inorganic, highly insoluble sulfate salt, is of relevance to the field of biomineralization, both as a biomineral found in some organisms1,2 but also because of the “Barite paradox” whereby Barite appears to form in undersaturated seawater.3,4 This paradox has been linked to both the occurrence of strontium and silicate in seawater,4−6 and the presence of Barite in sediments is often taken as an indicator of significant biological productivity.7 In addition, silica and silicates are ubiquitous in nature and are one of the most common minerals found on Earth, making up >90% of the Earth’s crust. The chemistry of silicates can be complex and can involve inorganic polymerization of the silicate species, depending on the solution composition.8 Perhaps for this reason, few studies on the impact of silicate species can be found on either calcium carbonate or barium sulfate. One of the few studies available is that of the Pina group on calcium carbonate9 in which they found that polysilicic acids have a different effect to the monosilicic acids. Another study is that of Lakshminarayanan which showed that the silicate ion was incorporated into the calcite structure.10 In these studies, the impact of silicate on crystallization has been found to be very sensitive to pH and the mono-/polymeric form. A similar result is reported for strontium sulfate.11 Calcium carbonate and barium sulfate are both minerals of interest from a scaling perspective.12−15 In addition, calcium carbonate is also a significant biomineral,16 and a vast amount of literature on the impact of impurities on calcium carbonate can be found.17−20 This is understandable, but calcium carbonate has some experimental difficulties related to its crystallization that barium sulfate does not. These include the issues of polymorph formation and solution speciation. Three common crystalline polymorphs are known for calcium © 2012 American Chemical Society

carbonate (vaterite, aragonite, calcite) at room temperature and pressure, while only Barite is known for barium sulfate. Carbonate speciation in water is variable depending on pH, while sulfate is essentially 100% as the sulfate ion provided the pH > 4. These considerations make barium sulfate crystallization a suitable candidate as a model system, and we have used it as such to gain some fundamental insights into crystallization generally.21−23 Inorganic solids having curved nongeometrical shapes in nature were previously thought to have had contact with biological or organic material. Work by the group of GarciaRuiz24 showed that some structures having curved surfaces are in fact formed from purely inorganic sources and that, therefore, it is erroneous to make such assumptions. The work of Garcia-Ruiz presented the formation of barium carbonate/silicate structures that they named “biomorphs”. From this work it was shown that silicate at high pH is required to form such structures. The work discussed herein investigates the influence of silicate ions on the formation of barium sulfate at both biological (pH 7) and high pH (≥10). The impact of the silicate ion on the morphology, crystallinity, and structure of barium sulfate in the presence of silicates is discussed with respect to crystallization phenomena and compared to the calcium carbonate,9,10 strontium sulfate,11 and barium carbonate24 systems. Received: February 23, 2012 Revised: April 11, 2012 Published: April 16, 2012 3057

dx.doi.org/10.1021/cg300263f | Cryst. Growth Des. 2012, 12, 3057−3065

Crystal Growth & Design

Article

Figure 1. Effect of silicate on barium sulfate morphology at pH 7 (a) 0 mM SiO2, (b) 0.05 mM, (c) 0.24 mM, and (d) 0.94 mM SiO2 (inset shows close up of side of aggregate).

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distance of 10 mm (optimal for this instrument) and a voltage of 15 kV. Powder X-ray Diffraction (XRD). In order to obtain sufficient solids for XRD, the crystallization experiment was scaled up from 20 mL to 4 L. The proportions were equivalent to the small-scale experiments, the only differences being that the temperature was not controlled (thus being conducted at room temperature, ∼22 °C), and the solids were magnetically stirred for 3 h after commencement of the reaction by addition of sulfate and then left for 3 days to stand. After this time, the supernatant was decanted and the solids were obtained by filtration through a 0.2 μm membrane. The solids were washed three times with ultrapure water. One set of solids was then dried in a desiccator while a second set of solids was stored wet in a sealed vial for XRD. Wet samples were also investigated by XRD to ascertain whether the presence of silicate induced the formation of a hemihydrate as was observed for strontium.11 The XRD patterns were collected on a Bruker D8 Advance instrument using Cu Kα radiation. A low background holder was used and spun at 30 rpm. The 2 theta range was 15−50° with a step size of 0.001° and a divergence slit of 0.3°. Transmission Electron Microscopy (TEM). From the large batch produced for XRD, a portion of the 1.0 mM silicate (pH 10) dry sample was also set in resin for ultramicrotomy. The thin sections produced from this were viewed with a Jeol 2011 TEM at 200 kV. Thermal Analysis. Approximately 15 mg of sample was heated in a platinum pan using a TA Instruments SDT 2960 simultaneous DSCTGA instrument. The temperature ranged from ambient to 800 at 5 °C per minute in air at a flow rate of 40 mL/min. The temperature of the instrument was calibrated against the melting points of indium, zinc, tin, silver, and gold. In addition, the balance was calibrated over the temperature range with standard alumina weights as provided by the vendor. Fourier Transform Infrared (FTIR). The solids for FTIR were dried in an oven at 100 °C for 24 h to remove surface water. They were then formed into KBr discs before obtaining spectra on a Bruker IFS 66 instrument with 4 cm−1 resolution and 4 scans per spectrum and a 4000−400 cm−1 range.

MATERIALS AND METHODS

All materials were analytical grade reagents used as received. Ultrapure water (resistivity >18 MΩ cm) was used in the preparation of all solutions. Crystallization Morphology. The morphology of the barium sulfate formed in the presence of silicate ions was determined in static, batch crystallization experiments in glass vials. The cleaned glass vials were filled with ultrapure water to the desired volume, and barium chloride (50 μL, 0.1 M) stock solution was then added to the vial. Sodium metasilicate pentahydrate was prepared fresh by dissolving into water at a concentration of 4.7 mM. The pH was then measured using a pH meter (Orion, from Walker Scientific) and adjusted to the desired pH using HCl. It was subsequently diluted to the desired level by adjusting the volume of addition to the vial. A cleaned, round, glass coverslip was placed at the bottom of the vial and the vial placed in a thermostatted water bath at 25 °C to equilibrate for 1/2 h before the addition of a stoichiometric amount of sodium sulfate (50 μL, 0.1 M) to commence the crystallization reaction. The total volume of the crystallization experiment was kept constant at 20.1 mL. After 3 days the glass coverslips were removed and the excess solution was soaked up by tissue paper. The coverslip was then prepared for imaging. The supersaturation (defined as √(IAP/Ksp)) for these experiments was found to be 18.8 at pH 7 and 18.4 at pH 10 (using PHREEQC).25 Unfortunately, attempts to find barium silicate solubility (or complexation) data were unsuccessful, and so this value does not take into consideration these interactions. Scanning Electron Microscopy (SEM). The coverslips obtained from crystallization experiments were placed on carbon coated SEM stubs and carbon paint applied to the circumference to help avoid charging effects. The stubs were placed in a desiccator to let the stubs dry and then were sputter coated with gold prior to viewing on a XL 30 Philips or an Evo Zeiss SEM instrument. In addition, the ultramicrotomed slices of barium sulfate formed in the presence of high concentrations of silicate were prepared for SEM by gold sputtering and analyzed by energy dispersive X-ray spectroscopy (EDX). This was used in a qualitative manner only; thus flat, polished surfaces were not necessary. The spectra were collected at a working 3058



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Figure 2. Impact of silicate ions at pH 10 on barium sulfate particles (a) 0.24 mM, (b) 1.0 mM, (c) 1.50 mM SiO2, (d) close-up of a particle formed in the presence of 1.0 mM SiO2.

Figure 3. (a) SEM image and (b) EDX map of Si Kα 1 line.

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RESULTS The impact of silicate ions on barium sulfate crystal growth at both neutral and high pH was investigated. At pH 7 we can expect mainly H4SiO4 species to exist, while at pH 10, there are approximately equal amounts of H4SiO4 and H3SiO4− present (ref 24; see Supporting Information, Figure S1). In addition, depending on the concentration, either monomeric or polymeric silicate species are expected.8 At the concentrations used in this work, however, monomers would be expected at high pH while polymeric species could exist at pH 7. Sulfate on the other hand is essentially 100% deprotonated for pH values ≥4. Thus, unlike the carbonate biomorphs formed with silicate,24 the formation of barium sulfate will not involve any subtle pH cycles and the sulfate speciation will remain constant throughout. The barium sulfate morphology in the absence of additives (Figure 1a) is not sensitive to pH at 25 °C.23 When silicate is present at pH 7, the dominant effect is to promote aggregration (see Figure 1b−d) particularly on the (hk0) face.

The work of Pina has shown that, in the case of calcium carbonate, the presence of silicates can be seen in atomic force microscopy (AFM) images as new nuclei.9 Thus, the protonated silicate was acting as a heterogeneous nucleation site. We would assume that the behavior would be similar here for the protonated, and possibly polymerized, silicate species in the presence of barium sulfate. Thus, this heterogeneous nucleation promotion also promotes the formation of crystal aggregates. At pH 10, it was observed that the initial impact of silicate is to limit the c-axis length (and therefore growth in the ⟨001⟩ direction; see Figure 2a). The aggregation previously observed at pH 7 on the (hk0) faces is absent at the higher pH. Similarly, Pina found that the presence of silicate can inhibit (001) growth of strontium sulfate as observed by AFM.11 At higher silicate concentrations, barium sulfate forms a fibrous looking particle (Figure 2b); however, this fibrous appearance may be limited to the surface as the particles formed in the presence of 1.50 mM SiO2 show (Figure 2c). In these small-scale 3059



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Figure 4. Large batch, dry XRD patterns of barium sulfate particles formed in the presence of 1.0 mM SiO2 at pH 7 and 10 (spectra have been offset in the y axis for clarity).

Figure 5. XRD patterns obtained for barium sulfate solids collected at different SiO2 concentrations and pH 10.

mM SiO2, suggesting that changes in the scale of reaction, with associated differences in mixing and so on, did not cause significant changes in the crystal growth processes. The ultramicrotomed sample used for TEM analysis was also prepared for SEM, and EDX mapping was then conducted on the sample. EDX mapping showed Ba as expected (see Supporting Information, Figure S3), and the Si distribution showed no particular preference for the surface suggesting that the Si is probably substituting into the lattice (see Figure 3).

experiments, 5 mM SiO2 produced a solid block that was difficult to filter (and was found to be a barium silicate; see XRD Analysis section). Examining the fibrous particle shown in Figure 2b at higher magnification (Figure 2d) revealed small spherical particles that could not be individually characterized but could conceivably be a barium silicate material. The large (4 L) batch sample at pH 10 was also investigated by SEM and found to have a similar morphology, showing fibrous particles (Supporting Information, Figure S2) at 1.0 3060



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Figure 6. XRD patterns for barium sulfate and barium silicate solids collected at pH 10 and 0.94 and 2.36 mM SiO2.

Figure 7. Thermal analysis curves of weight loss versus temperature for different silicate levels present during barium sulfate crystallization.

It is clear that the pH 7 sample is the most similar to the control XRD (Supporting Information, Figure S4) with some differences in peak intensities. In particular, the 102 and the 211 peaks are significantly more intense than found in the control. This is reasonable since the Barite formed at pH 7 in the presence of silicate appears to be a more aggregated version of the normal morphology. Thus, the XRD reflects that more of the Miller planes are accessible to the beam owing to its spherical agglomeration but otherwise similar morphology to the control.

The roughness of the surface prevented quantitative assessment of the silicate present. The smaller subparticles evident in the SEM would suggest that the solids are delicate and may have broken (this was also found for Barite formed in the presence of mellitic acid23). XRD Analysis. XRD of the dried barium sulfate samples formed in the presence of silicate show that there is no substantial line broadening observed when compared to the control (Supporting Information Figure S3) regardless of the pH of the sample. 3061



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Figure 8. FTIR spectra of Barite with different concentrations of silicate present − straight gray lines indicate the appearance of barium silicate bands.

weight loss, barium silicates tend to have melting points above 1000 °C; thus, the weight loss cannot be due to the loss/ decomposition of SiO2 or related species and must be due to some other phenomenon, including the possibility of water. If we assume for the moment that this is true and that the majority of solids are barium sulfate, then at 2.36 mM SiO2, with a weight loss of 9.49 wt % this would equate to ∼0.5 mol of H2O per 0.4 mol of barium sulfate (more than 1:1!). We could instead consider the weight loss below 600 °C as being more likely to be due to structural water loss in which case this would give ∼0.5 mol of H2O per mol of barium sulfate. However, no hemihydrate is observed in the XRD patterns, and thus water in a fixed structural position can be ruled out. No doubt barium silicate is also precipitating with barium sulfate as seen in the XRD patterns in Figure 5; however, if the silicate is amorphous it is not straightforward to determine the relative amounts of barium silicate and barium sulfate. The presence of such an amorphous silicate, that has water associated with it, could explain such a significant weight loss on heating. The FTIR spectra of the two barium sulfate samples with silicate present were compared to pure Barite (Figure 8). This showed that as the amount of silicate increased, bands commensurate with barium silicate became more and more prominent.27 The barium sulfate bands were still discernible but the silicate bands are significant: in the presence of 2.36 mM SiO2, the silicate bands are stronger than the sulfate bands. This is even more surprising when one considers that in the XRD patterns the majority of the silicate peaks only begin to appear above baseline for SiO2 concentrations ≥1.89 mM. The one peak that is observable even at 0.94 mM SiO2 is that at ∼43° 2 theta. This could be because the silicate phase has substantial amorphous character. Perhaps of more interest is that in the presence of silicate, even at very low concentrations, significant O−H stretching is observed in the 3000−4000 cm−1 region. If we take the ratio of the OH stretch peak height to the strongest sulfate peak

In the case of the Barite formed at higher pH with silicate present, the peak intensities differ significantly from the control. However, the lack of significant line-broadening in the XRD pattern, even when the barium sulfate appears to be very fibrous, suggests that this is not due to mesocrystal formation.26 In the case of strontium sulfate,11 the presence of silicate induced the hemihydrate form to crystallize. Barium sulfate is not known to have a hemihydrate under ambient conditions but wet samples were investigated with XRD in case drying resulted in a phase change. Wet samples showed similar XRD patterns (Figure 6) to the dry samples’ (SEM images can be found in the Supporting Information, Figure S5). The block of solids formed at 5 mM SiO2 had an XRD pattern (labeled Basilicate) that corresponds best to a barium silicate pattern (albeit with significant line broadening, a mixture of PDFD number 004-0422 - Ba 2 SiO 3 O 8 - and 027-0647 − Na2Ba4Si10O25). In the “wet” XRD patterns at 2.36 mM SiO2, a mixture of barium sulfate with barium silicate is obtained. In addition, at 2.36 mM SiO2, in the “wet” samples, the peaks observed for the barium silicate are discernible with much sharper peaks. We can also see, by overlaying the BaSO4 + 0.94 mM SiO2 and the BaSO4 + 2.36 mM SiO2 XRD patterns, that the 2 theta value of the barium sulfate peaks are shifting to lower 2 theta values or larger d spacings (Figure 6) with increasing silicate concentration, as would be expected for silicate substitution.10,11 The calculated changes in the a-, b-, c- lattice parameters were 0.3−0.8%, suggesting only a small degree of incorporation. Thermal Analysis and FTIR. The XRD patterns did not show any impact of water within the structure when silicate is present. Thus, thermal analysis and FTIR were undertaken to investigate this further. The thermal analysis curves (Figure 7) show that in fact as the silicate content increases, so too does the weight lost from ambient (∼23 °C) up to 850−900 °C. Despite this significant 3062



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Figure 9. (a, b) TEM images of barium sulfate particles formed at pH 10 and in the presence of 1.0 mM silicate − colored circles represent areas where selected area diffraction patterns were collected.

(∼1079 cm−1) height we obtain ratios of 0.07, 0.72, and 0.62 for the presence of 0, 0.94, and 2.36 mM SiO2 respectively. While this is not strictly quantitative it shows that much more water is present in the Barite structure when silicate is also present. The broadness of the O−H stretch region suggests that the water is very liquid like. Given the results from the thermal and the infrared analysis, water is probably present as inclusions and/or associated with amorphous silicate. Clearly, water is present in the structure when silicate is also present, and the maximum of the O−H stretch is at a slightly higher wavenumber than is found for pure water (∼3420 cm−1 in the presence of silicate c/f ∼3400 cm−1 for pure water28). This may reflect a slightly stronger hydrogen bond or more mass associated with the hydrogen bond. TEM Analysis. Analysis of the particles formed at high pH and 1.0 mM SiO2 were further investigated by TEM. The results show that these particles are single crystals. The fiberlike structures observed by SEM appear to be brittle and break off easily in the ultramicrotomy process. This is consistent with the low number of particles with obvious fibrous protusions observed by TEM. The fibrous features (shown as colored circles) are selected areas for electron diffraction (SAED). It was found that the particles are crystalline (Figure 10) as spots are observed. In the main body of the barium sulfate particle some polycrystallinity is observed (hint of rings), but otherwise the pattern is consistent with single crystal formation.

These SAED patterns are most similar to the 001 zone simulated spectra meaning the long fibrous particles are elongated in the zone perpendicular to this axis. We assume that the long axis is the ⟨010⟩ when these data are compared to the SEM images at low silicate concentrations where the c-axis is noticeably shortened by the presence of silicate. Higher magnification of the particles from the fibrous regions shows some channels (Figure 11). However, it is clear that the

Figure 11. (Left, right) High magnification TEM images of barium sulfate formed at pH 10 and in the presence of 1.0 mM SiO2 showing porous channel-like structure for the fibrous sections.

presence of silicate is not inducing “biomorphic” (curved) structures during the crystallization of barium sulfate, despite the significant change in morphology. The fibrous nature of the solids observed in this work is consistent with that observed for strontium sulfate although those particles were attributed to the hemihydrate.11 The lack of biomorph formation for barium sulfate when compared to barium carbonate24 is hypothesized to be due to the protonation state of sulfate. As stated previously, unlike the carbonate anion, the sulfate ion speciation does not change above pH 4. Thus, there are no pH fluctuations great enough in this system to alter the dominating sulfate species.24 However, it is interesting to note that silicate is still able to incorporate into the structure, cause morphological changes, and form fibrous-looking particles or cluster-like particles depending on pH. The work by Lakshminarayanan10 suggested that the silicate ion substituted for carbonate in the calcium carbonate lattice and that this was particularly prevalent on positively charged

Figure 10. (a, b) Selected area electron diffraction (SAED) patterns of circled regions labeled (a) and (b) respectively shown in Figure 9b and (c) theoretical SAED patterns for the 001 zone of barium sulfate as constructed by CrystalDiffract. 3063



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surfaces. Thus, in this work we hypothesize the silicate is also substituting for sulfate, as supported by the XRD results. The sulfate ion has a S−O bond length of about 149 pm29 while that of Si−O is larger at 163 pm.30 Thus, even in barium sulfate the substitution of silicate will require accommodation of a 9% larger ion.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: (618) 9266 7677. Fax: (618) 9266 4699. Notes

The authors declare no competing financial interest.

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CONCLUSIONS Biomorph formation in the presence of silicates is not observed for barium sulfate. The sulfate is essentially deprotonated from pH 4, and small changes in pH at either 7 or 10 are not sufficient to cause a cycling of the sulfate speciation as occurs for carbonate. The impact of silicate on calcite as investigated by AFM is also consistent with the results discussed here in that the growth island is altered from rhombic to ellipsoid;9 however, in the case of calcite one would predict the possibility of biomorph formation for the same reasons that it is observed for barium carbonate24 that being that the carbonate speciation is sensitive to small changes in pH. At pH 7 the presence of silicate during barium sulfate crystallization results in particles that are highly aggregated and roughened, leading to structures reminiscent (although not exactly the same) of “Barite roses”. In terms of the “Barite paradox”, the results here require further investigation into the impact of dissolved silicate on nucleation rates and, in particular, the role of polymeric versus monomeric silicate on the morphological/thermodynamic and kinetic impacts of silicate on barium sulfate crystallization. With this information, a clearer understanding of why barium sulfate crystallizes in seawater can be obtained. This will be the focus of future work. At pH 10, an unusual fibrous structure appears that is crystalline and highly porous. XRD analysis shows that the silicate partially substitutes for sulfate. TEM images and analysis show that these particles appear to be single crystals rather than mesocrystals. The TEM images would suggest that the fibrous nature of the particles is limited to the surface; however, SEM/ ultramicrotomy shows that the internal structure of the particles is fragile, and thus, silicate is also impacting on the structure as a whole. In addition, TEM clearly shows the fibres are quite porous. The elongation zone of the fibers appears to be in the ⟨010⟩ direction. Finally, thermal analysis and FTIR clearly demonstrate the presence of water in the structure when silicate is present. This water is not observed in the XRD pattern and so cannot be associated with the formation of a barium sulfate hemihydrate. Nor is the separate silicate phase observed in the XRD until much higher silicate concentrations are present compared to the FTIR results. This would suggest an amorphous silicate phase is formed that crystallizes as more silicate is present. However, it appears that when silicate incorporates, it does so with some water. The water does not appear to have a defined structural position as indicated by the broad O−H stretch in the FTIR.

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ACKNOWLEDGMENTS We thank Peter Chapman for performing the thermal analyses, and we wish to acknowledge the Curtin Centre for Materials Research for use of the SEM, TEM, and XRD facilities. We thank the anonymous reviewers for their constructive comments.

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REFERENCES

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ASSOCIATED CONTENT

S Supporting Information *

Speciation graph, SEMs of batches prepared for XRD, XRD pattern of control barium sulfate, and EDX result for barium. This information is available free of charge via the Internet at http://pubs.acs.org/. 3064



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(29) Baird, N. C.; Taylor, K. F. J. Comput. Chem. 1981, 2, 225. (30) Boisen, M. B., Jr; Gibbs, C. V.; Downs, R. T.; D’Arco, P. Am. Mineral. 1990, 75, 748.

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Effect of Solution Silicate on the Precipitation of Barium Sulfate

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