Improved Control of CaCO3 Precipitation by Direct Carbon Dioxide

Jun 23, 2006 - has the advantage of effectively introducing only the ions needed for ... and the influence of spectator ions on crystallization can be...
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Improved Control of CaCO3 Precipitation by Direct Carbon Dioxide Diffusion: Application in Mesocrystal Assembly Miles G. Page* and Helmut Co¨lfen Max Planck Institute for Colloid and Interface Science, Potsdam, Germany

CRYSTAL GROWTH & DESIGN 2006 VOL. 6, NO. 8 1915-1920

ReceiVed March 18, 2006; ReVised Manuscript ReceiVed May 11, 2006

ABSTRACT: We introduce a new crystallization method for CaCO3, via CO2 vapor diffusion into a Ca(OH)2 solution. This method has the advantage of effectively introducing only the ions needed for crystallization so that the solution ionic strength can be controlled and the influence of spectator ions on crystallization can be studied. As a further advantage, a pH close to biological conditions is approached at the end of the crystallization reaction, and the kinetics of the reaction can be influenced via the initial pressure in the closed crystallization setup. The method was validated for two examples of CaCO3 crystallization via polymer-mediated mesoscale assembly. Introduction Crystallization control is among the most important techniques of preparation, purification, and application of solid substances. Control of crystal size and shape ensures dissolution rates (for example, of pharmaceuticals), and control of crystal shape and intergrain texture defines flow properties, filling grades, and mechanical properties. A high reproducibility of the chosen procedure ensures the quality of industrial products and is of utmost importance. As nucleation and growth are very sensitive processes, control can be improved by addition of nucleation agents, stabilizers, templates, or ternary components in general. The choice of solvents,1 low molecular weight additives, surfactants, and functional polymers is regularly reported (for recent reviews, see refs 2-5). However, despite decades of intense investigation, crystallization is still not fully understood, as expressed in the lack of complete theoretical models and in the experimental difficulties to fully control crystallization reactions. Therefore, the development of defined crystallization protocols is essential. This is especially true for CaCO3, a well-studied model crystal system due to its industrial and scientific importance.4,6 Research in recent years revealed nonclassical crystallization mechanisms such as oriented attachment, the fusion along crystallographic lines of nanoparticles to form single crystals,7-9 and the formation of so-called mesocrystalssfacetted, colloidal crystals of crystallographically aligned, nonspherical nanoparticles.10 Well-known for example in transition metal oxide7 and amino acid11 crystallization, they have however only recently been identified in calcium carbonate, for both vaterite12,13 and calcite.14,15 Although these are the only so-far reported mesocrystal-driven calcium carbonate structures, it is likely to prove to be a more general phenomenon,10 as identification of mesocrystal structures is not trivial. For example, the case of the rod/dumbbell/sphere morphology observed in calcite crystallization in the presence of the double-hydrophilic block copolymer poly(ethylene oxide)-b-poly(methacrylic acid)16 (PEOb-PMAA) was recently found to be a mesocrystalline assembly,17 but the formation mechanisms are inherently difficult to reveal. Phenomena such as selective crystal face stabilization can account for structural modification by simple ions or polyelectrolytes in the context of ion-by-ion growth. However, in a mesoscale assembly process, one needs to consider other * To whom correspondence should be addressed. Tel.: +49 331 567 9552. Fax: +49 331 567 9502. E-mail: [email protected].

mechanisms that influence growth by self-assembly into higherorder structures. For example, assembly of calcite nanocrystals stabilized by poly(styrene sulfonate)14,15 was suggested to be controlled by an “inner field” in the growing superstructure due to a polymer adsorption-induced dipole in the contributing nanocrystals, but such electric field-mediated assembly, first suggested by Kniep et al.,18 has still not been definitively established experimentally. Complex single-crystal assemblies of calcite and vaterite reported by Gehrke et al.13 were attributed to oriented attachment of vaterite nanoparticles triggered, remarkably, by adsorption of the reaction byproduct ammonia using the vapor diffusion technique19 without other additives. Similar ammonia-driven effects were also reported by Neira et al.,20 who systematically studied variation in crystal morphology with kinetic conditions in this method, and Kojima et al.,21 who studied crystals formed using the Kitano process22 with added ammonia. Apart from effects related to the presence of ammonium, mesoscale assembly processes suggest the need for a crystallization setup that avoids, or better controls, variables such as ionic strength and ionic species. Furthermore, minimization and/ or control and variation of ionic strength are of interest in for example crystallization in surfactant phases and (micro)emulsions, where phase equilibria and physical characteristics are strongly dependent on ionic species and ionic strength. The Kitano process22 is a clean technique in terms of the species present in the system. However, crystallization is favored at the air-water interface due to the concentration gradient perpendicular to the interface and around bubbles of evolving CO2, often with profound effects on the final crystal morphology.23 The double jet technique24,25 is particularly good for timedependent analysis. However, the monotonically increasing ion concentration results in constantly changing conditions, especially the concentration of both crystallizing and spectator (e.g., Na+ and Cl-) ions. Similarly, the double-diffusion technique,26,27 although it provides excellent control of conditions, employs soluble salts of calcium and carbonate that mix under diffusion through a membrane, requiring noncrystallizing species. Furthermore, it requires a specialized experimental setup, which is possibly a reason its use is not more commonly reported in the literature. Another technique, the so-called constant composition method introduced by Nancollas,28,29 aims at a constant composition of the crystallizing ions, but the concentration of spectator ions is constantly increased, as salt is continuously added to maintain constant chemical potential of the precipitating

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species. Finally, a nice technique developed by Matijevic is the thermal30 or enzymatic31 decomposition of urea, although this also involves ammonia and either, respectively, high temperature or added enzymes that were themselves found to modify the growth of the CaCO3.31 Here we present a hybrid Kitano/vapor diffusion crystallization method that opens further possibilities for control and variation of crystallization conditions, whereby CO2 evolves from a supersaturated source solution and diffuses into a crystallizing calcium hydroxide solution. It combines the advantages of simplicity and kinetic reproducibility of the ammonium carbonate vapor diffusion technique, with the minimized presence of extraneous ion species of the Kitano process. Calcium carbonate precipitation from calcium hydroxide solutions or suspensions has also been studied with a reactive crystallization setup by either bubbling CO232 or mixing with CO2 solution33 to study rates of nucleation and growth, and aggregation of calcite forming at the air-water interface.34 Here however we develop a technique with the specific goal of understanding and controlling crystal morphogenesis in the light of recent advances in identification and exploitation of mesoscale assembly systems. We compare and contrast crystals obtained with this method, using two previously studied synthetic polymeric additives that induce mesoscale assembly: A double hydrophilic block copolymer, PEO-b-PMAA16,35 and a simple polyelectrolyte, poly(styrene sulfonate) (PSS),14 with those obtained using, respectively, the vapor diffusion and double jet crystallization techniques. Experimental Section In the vapor diffusion technique, decomposition of solid ammonium carbonate provides the source of carbonate in the system, and the high pH required to induce precipitation is simultaneously achieved by evolution and dissolution of ammonia. It is employed herein as a control technique, with experimental conditions as described by Wang et al.14 In the case of the Kitano process, a calcium-containing, supersaturated CO2 solution is allowed to stand open, resulting in precipitation due to CO2 diffusion in the opposite sense. CO2 evaporation shifts the carbonate/hydrogen carbonate equilibrium to the carbonate side resulting in CaCO3 precipitation.22 The hybrid technique used in this work is essentially a vapor diffusion process, where a separate volume of water, saturated with CO2, is placed with the crystallizing solution in a closed system. The ammonium carbonate diffusion and CO2 diffusion techniques are summarized respectively in Figure 1a,b. The concentration of ions can be kept to a minimum by using exclusively calcium hydroxide solution, in which case only one of [Ca2+] and pH can be selected, the other then being dependent. Alternatively, the pH and calcium concentration in the crystallizing solutions can be decoupled by combining the calcium hydroxide with either (for example) calcium chloride to increase [Ca2+] or sodium hydroxide to increase pH. The required amounts of each for selected [Ca2+] and pH are precalculated such that the minimum amount of Na+ or Cl- is present. The CO2 diffusion is essentially regulated by chemical equilibrium of CO2 concentration in the two solutions, as well as a kinetic component governed by the partial pressure of CO2 under which the solutions are held. Crystallization was performed in 10 mL glass vials, cleaned by sonication in ethanol followed by rinsing in doubly distilled water and air-dried. Solutions were prepared from freshly prepared stock solutions of calcium hydroxide, PEO-b-PMAA, and PSS in boiled, doubly distilled water. Glass slips were placed in the vials, on which the resulting crystals were examined by light microscopy and scanning electron microscopy (SEM). Remaining crystals were collected for wide angle X-ray scattering and transmission electron microscopy (TEM) analysis. Sample volumes used were 6 mL for crystallization with PEOb-PMAA, and 1 mL with PSS. In a typical experiment, CO2 is bubbled through 200.0 g of water until the pH reaches 4.00 (after around 30 min). This source solution

Page and Co¨lfen

Figure 1. Schematic of vapor diffusion processes in calcium carbonate precipitation. (a) CO2 diffusion by decomposition of ammonium carbonate. (b) CO2 diffusion by desolvation from saturated aqueous solution. is then connected to a 3 L desiccator holding the crystallizing solutions. The entire system is flushed with N2 gas, and the pressure inside the system is then reduced slightly inducing supersaturation in the CO2 source solution. These two steps are not essential to induce crystallization, which anyway will occur under the atmospheric partial pressure of CO2. The effect is simply a Henry’s law equilibration of the partial pressure, pCO2, which in the closed system is reduced by reaction of CO2 with the crystallizing solution, counteracted by desolvation from the then supersaturated reservoir. The reduced pressure acts to increase slightly the rate of CO2 transfer from the source to the crystallizing solutions. Reducing the pressure in this manner introduces a kinetic variable that for the resulting crystals presented here was relatively poorly controlled; however, this is easily rectified by adding a pressure gauge to the setup. In the results section below, such a survey of the progress of the crystallizing reaction was performed under conditions as near as practical to that of the experimental crystallization reaction. The setup was reversed such that a source CO2 solution (500 g) was placed in the desiccator, and this then was connected to a two-necked flask containing calcium hydroxide solution (40 mL, 5.0 mM), to which was fitted a pH electrode through which the pH over the duration of the reaction was monitored versus time using a Metrohm 702 SM “Titrino” auto-titration machine. The system was again flushed with N2 gas, and then the pressure was reduced by a fixed amount, controlled with a vacuum pump and attached Vacuubrand CVC-2 vacuum controller that was also used subsequently to monitor the total pressure in the system. The larger crystallizing solution is necessary to properly immerse the pH electrode. While this could affect the kinetics somewhat, the larger CO2 source ensures that there is more CO2 available than is necessary for the reaction, and it can be reasonably expected that trends observed in the test case would also be followed in different system dimensions even if the precise time scale varies. The following chemical products were obtained from Sigma-Aldrich and employed without further purification: calcium hydroxide, 99% (up to 3% carbonate), PSS (Mw ∼ 70000 g mol-1), calcium chloride dihydrate (>99%). PEO-b-PMAA (PEO ) 3000 g mol-1, PMAA ) 700 g mol-1) was kindly donated by Th. Goldschmidt AG, Essen (Germany), and was exhaustively dialyzed against water using a 1000 g mol-1 molar weight cut off dialysis membrane.

Results In the test cases of Ca(OH)2 solution (40 mL; 5.0 mM) with no additive, the system was first evacuated to the desired initial

Control of CaCO3 Precipitation by CO2 Diffusion

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Figure 3. Light microscope images of CaCO3 rhombohedra obtained from calcium hydroxide solution (5 mM) in the absence of other additives. (a) P0 ) 0.4 atm, (b) P0 ) 0.8 atm. Scalebar: 50 µm. Note the rounded edges on one side as a result of aging at the neutral, final pH. Inset: Crystals grown by fast precipitation (∼2 h) from calcium hydroxide under pCO2 ) P ) 1 atm, showing smaller crystals without signs of aging (same scale).

Figure 2. (a) Variation of pH with time (in days, bottom axis) of calcium hydroxide (5 mM, pH0 ) 12.0), in the presence of a CO2 solution, saturated at atmospheric pressure, under initial pressures Patm (in atmospheres, atm) of 0.2 (grey squares), 0.4 (black triangles), and 0.9 (grey circles), and with time in hours, top axis, at Patm ) pCO2 ) 1 atm (black circles). (b) Variation of pH (open circles) and conductivity (closed circles) vs time of a Ca(OH)2 suspension, during bubbling with CO2. The pH follows an acid-base titration-type behavior (a and b), whereas the conductivity (a) passes through a minimum at pH ∼9. This minimum in dissolved ion concentration is observed due to precipitation of calcium carbonate, followed by redissolution as hydrogencarbonate with further addition of CO2.

pressure (P0) of 0.9, 0.4, and 0.2 atm. For P0 ) 0.9 atm, the pH fell from 12.0 to 7.5 over 2.6 days, following an acid-base titration-shaped curve of pH vs time, shown in Figure 2a. Visual observation of surface crystals34 was possible after several hours. The pressure in the system rose from 0.9 atm to about one atmosphere within the first 6 h due to evolution of CO2 (and presumably some water vapor). In the case of P0 ) 0.2 atm and P0 ) 0.4 atm, similar observations result. However, for a well-sealed system, the internal pressure remained below atmospheric pressure throughout the experiment time, followed by similar titration curves. There is an increased “lag time” (Figure 2a) before the pH endpoint for these two cases; however, at the lowest initial pressure of 0.2 atm, the pH time curve near the “end point” is also significantly steeper. This effect is amplified in one final case, in which the system was initially flushed with CO2, resulting in pCO2 of 1 atm at time zero. In this case, the titration process runs to completion in a few hours (Figure 2, top axis). The final pH is reasonably stable to variation in conditions, varying between 6.5 and 7.5, depending upon the final pCO2, but apparently buffered to an extent by the redissolution of calcium carbonate. Thus, it can be seen that the rate of dissolution of the CO2/ reaction with calcium hydroxide can be controlled by varying the total initial pressure in the system: the total final pressure decreasing, with the associated increased lag-time, and the final

pCO2 (approximately Pfinal - P0) increasing, with an increased steepness of the titration curve and decrease in the final pH. This is incidentally a nice demonstration of the kinetics of CO2 dissolution (dependent upon both P and pCO2), and the total solubility (dependent only upon pCO2). Note that the variation in pH is not necessarily coupled to nucleation or crystal growth periods, rather only to the CO2 loading. It may be of later interest to examine, however, if the position in the acid-base equilibrium influences the progress from nucleation to growthdominated phases. Figure 2b shows the variation of pH with time of a suspension of calcium hydroxide (above the saturation concentration, resulting in a solution that remains saturated for the duration of the experiment), bubbled directly with CO2. From this plot one can see that the final pH of the crystallizing solution is rather important, in particular whether the final pH is above or below the minimum in Ca2+ solubility (pH ∼ 8.5 in the case of the initially saturated solution), as measured by the conductivity of the system, which controls whether significant redissolution will occur. The effect of particle aging is revealed by light microscopy (Figure 3) where calcite rhombohedra are rounded on one side where redissolution aging has occurred for P0 ) 0.4 and 0.8 atm. This is not observed in the crystals grown under 1 atm CO2 (inset), where presumably the time was not sufficient for partial redissolution of well-established micron-sized crystals despite a comparable final pH. Crystallization with PEO-b-PMAA. Crystallization using the double jet technique under conditions described by Co¨lfen et al.,16 in the presence of PEO-b-PMAA, follows a roddumbbell-sphere morphological transition with time. The mechanism for this transition was proposed to be the result of orientation of nanocrystals along dipole field lines. It was recently observed17 by microfocused X-ray diffraction that the early rodlike particles are oriented mesocrystal assemblies, which first fuse to single-crystalline rods, then grow into polycrystalline dumbbells and eventually spherical superstructures. Figure 4 shows the result of crystallization by the CO2 diffusion technique, using 0.1 g L-1 of the same block copolymer, PEO-b-PMAA, at two different concentrations of Ca2+, aged for 4 weeks in situ. The results can also be compared to the morphology sequence obtained by simple mixing of CaCl2

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Figure 4. SEM images of calcite particles grown by CO2 diffusion in 0.1 g L-1 PEO-b-PMAA and (a, b) 2 mM; (c, d) 6 mM Ca2+. Panel (d) shows a zoom into a fractured portion of the dumbbells seen in (c), showing the fibrous internal structure.

and Na2CO3 in the presence of the same polymer.16 At 2 mM Ca2+ (Figure 4a), the particles maintain the rodlike morphology observed early in the analogous double jet experiment. Interestingly, however, we note also many centrosymmetric, “multipolar” structures. These structures were not observed in the previous double jet precipitation reactions under the same reactant conditions, and the reason for their formation is not yet clear, although quadrupole-shaped species were also observed as a second population in double jet experiments.17 The close-up image (b) shows a similar, although rather smoother, crystal surface as compared to the rodlike particles obtained using the double jet technique. At the higher concentration, 6 mM Ca2+, the main morphology is the dumbbell (c) although some instances of the more developed “twins”, full spheres, and quadrupoles are also observed. A close-up of a fractured dumbbell (d) shows the same internal structure of these particles. However, in several cases the dumbbell character is more pronounced, with the rodlike particle core still exposed between two spheroidal outgrowths, somewhat reminiscent of dumbbells of fluoroapatite grown in gelatin gels.18 These rods (Figure 4a,b) were found17 to consist of single crystals formed by the oriented assembly of nanocrystals, whereas the dumbbells, Figure 4c,d are a polycrystalline, although still highly spatially ordered, nanocrystal assembly.17 This is demonstrated in Figure 4b,d, with the rodlike particle displaying the smooth, probably (104) face, whereas individual, assembled nanocrystals comprise the ends of the dumbbell. The reason for the consistent, time-dependent transition in this case from rod to dumbbell is still not fully established. The rodlike particles observed at 2 mM Ca2+ are similar in length to those obtained by direct mixing of Na2CO3 and CaCl2,16 but significantly longer, at around 15 µm, than those observed by the double jet technique, in which they are not observed to be longer than ∼10 µm.36 This shows that the instability in the nanocrystal aggregation that drives the roddumbbell transition is driven by calcium concentration (both supersaturation and concentration relative to that of the polymer), rather than a property such as the size that is intrinsic to the particle. Similarly, in the dumbbell-forming vapor diffusion conditions (6 mM Ca2+, Figure 4c,d), the required higher calcium concentration is present from the beginning of the experiment, resulting in the smaller dumbbells of 3-4 µm in

Page and Co¨lfen

Figure 5. Vapor diffusion by decomposition of ammonium carbonate into solutions of calcium chloride19 and PSS. Scale bars: 5 µm. (a) Ca2+: 1.25 mM, PSS: 0.1 g/L; (b) Ca2+: 1.25 mM, PSS: 1 g/L; (c) Ca2+: 5 mM, PSS: 0.1 g/L; (d) Ca2+: 5 mM, PSS: 1 g/L.

size, rather than the ∼10 µm dumbbells that evolve out of equivalently sized rods in the double jet experiment. Crystallization with PSS. Shown in Figure 5 are SEM images obtained by vapor diffusion with ammonium carbonate, reproducing as closely as possible the conditions employed by Wang et al.14 We observe nevertheless some variation in the observed morphologies in two of the four cases (Figure 5b,c), which we attribute to other variables in the setup that are not controlled, such as the mass of ammonium carbonate used, the volume of the vapor-diffusion medium, etc.20 As pointed out by Gehrke et al.13 and Neira et al.,20 the sensitivity to such conditions results from the provision of several different crystallization pathways arising from the simultaneous introduction of the carbonate (via CO2) and the changing pH (via NH3), as well as the presence of NH4+ cations. Interestingly, the rosetta-shaped particle shown in Figure 5b are characteristic of vaterite particles and resemble closely those reported by these authors in the absence of any crystallization additive. This is therefore probably a modification resulting from the ammonium ion, which can adsorb to vaterite surfaces thus stabilizing this polymorph,13 rather than due to the polymer additive. The mechanism behind this inconsistency is uncertain, but it illustrates in part the usefulness of being able to avoid extraneous species in the system. Figure 6 shows the result obtained using the direct CO2 diffusion technique, using the same concentrations, collected and examined by SEM after 7 days of crystallization. Samples were examined again after 1 month of crystallization and showed in each case essentially the same morphology, in good general agreement with the results of Wang et al., although the morphologies seem to be generally more developed according to the crystallization model discussed therein. Significantly, the expression of the (001) crystal face is still observed due to charge stabilization by adsorption of the negatively charged polyelectrolyte. However, the crystals generally show a rougher surface than those obtained by the vapor diffusion technique (Figure 5). This can be explained by the fact that the end pH of our crystallization technique was between 8 and 9 (with some variation between samples), which favors increased CaCO3 dissolution-recrystallization (see Figure 2b), as compared to the alkaline end pH of about 9.5-10 in the vapor diffusion technique, which strongly favors CaCO3 crystallization.

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Figure 6. SEM images of crystals produced by diffusion of CO2 into Ca(OH)2/PSS solutions at pH 12.0. Scale bars: 2 µm. (a) Ca2+: 1.25 mM, PSS: 0.1 g/L; (b) Ca2+: 1.25 mM, PSS: 1 g/L; (c) Ca2+: 5 mM, PSS: 0.1 g/L; (d) Ca2+: 5 mM, PSS: 1 g/L.

The crystals were further analyzed by TEM (Figure 7), showing the particulate nature of the substructure. Under electron diffraction (inset), the calcite polymorph was identified as expected,14,15 and crystallographic orientation is observed at low supersaturation ([Ca2+] ) 1.25 mM, at both concentrations of PSS, shown in Figure 7a. The contrast between this and the nanocrystal assembly observed at higher supersaturation is observed in Figure 7b, where the assembled particles display a multicrystalline pattern under electron diffraction, corresponding to several differently oriented nanocrystals. Thus, the same [Ca2+]-dependent transition from stable to unstable nanocrystal assembly, resulting in respectively single- and ordered polycrystal superstructures, is observed whether PSS or PEO-bPMAA (with the commensurate variation in the particle morphology) is the structure-directing polymer. Some suggestion of an overall directing crystallographic unit is still present in Figure 6d, where the structure of the “depleted” surface displays triangular faceting, and to a lesser extent (c), where the same can be seen in the central “hole”. Polarized light microscopy (not shown) allows identification of the depleted surface as the (001) face of the c-axis.14 Therefore, the facet corners appear to retain the geometry of the suppressed (104) face, suggesting a competition between crystallographic orientation and external forces resulting in the ordered but not perfectly oriented nanocrystal assembly. Nevertheless, this picture remains unclear, and it is these open questions10 in mesocrystal assembly, and by extension overall crystal growth control, that we wish to be able to address in the future and that are the primary motivation for development of the presented crystallization technique. Discussion A number of variations from traditional vapor diffusion motivate the development of this technique. No ammonium or other extraneous ions are required, eliminating this variable and allowing ionic strength to be minimized or controlled. The sense of the pH change is reversed compared to the traditional vapor diffusion technique, resulting in an approach to, rather than departure from, biological pH conditions with time, with the additional possibility of controlling the final pH.

Figure 7. TEM image and selected area electron diffraction (inset) of calcium carbonate particles grown under CO2 diffusion with added PSS. (a) [Ca2+] ) 1.25 mM, [PSS] ) 0.1 g L-1. Electron diffraction indicates a single- or highly oriented calcite crystal, with diffraction peaks observed for the (110), (211), (122), and (300) faces. A similar pattern is observed for [Ca2+] ) 1.25 mM, [PSS] ) 1 g L-1. (b) [Ca2+] ) 5 mM, [PSS] ) 1 g L-1. Note the larger number of reflections corresponding to an ordered assembly of calcite subunits, with diffraction peaks observed for the (012), (104), (110), (113), (202), (024), (116) faces.

One limitation is the initial, pH0, of the system: With pH0 ) 11, using the conditions employed herein, no crystallization took place due to direct conversion from hydroxide to hydrogen carbonate ions via the carbonate intermediate. The rather high employed pH0 of 12.0 could, for example, reduce the size of precursor particles or nanocrystals due to more favored nucleation, compared with those found by ammonium carbonate decomposition; however, the final morphologies show such effects to be minimal. It has also been shown37 that nucleation at high pH can favor aragonite over calcite, although since this was only at pH > 13, it is not considered here. Meanwhile, the results indicate that a pH0 of 12, as an upper limit, does not adversely affect the role of the controlling additives for the above-presented cases. The final pH of the system, however, can be “frozen” at an arbitrary pH before reaching the plateau (Figure 2) by closing

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the connection to the CO2 reservoir and flushing the system briefly with inert gas, allowing particles to age at the chosen pH without further nucleation. A lower final pH allows further investigation of structures resulting from partial redissolution, whereas moderate (8-9) or higher (>9) values favor complete precipitation. Furthermore, the process can be repeated by alternating the atmosphere above the solution between inert and CO2-rich. This allows dissolution-precipitation aging by pH variation, using only the CO2 solubility equilibrium. Alternatively, the “auto-titration” monitoring system employed in the test-cases (with no additive) could be straightforwardly used instead for pH control, with calcium hydroxide as the titrant, allowing full control of the pH (necessarily coupled in this case with the Ca2+ concentration) throughout the reaction, without forfeiting other advantages of the technique. Conclusion The crystallization of carbonate by direct diffusion of CO2 into calcium hydroxide solution allows expanded control over contributing factors in crystal growth including the nature and concentration of ions and the resulting ionic strength. The validity of the technique introduced here is demonstrated by the effective reproduction of the newly observed mechanism of structure determination by self-assembly of mesocrystals, as shown by studying two model systems exhibiting this emerging phenomenon. It opens further possibilities for investigation in this direction through in situ aging at biologically comparable pH, and the possibility of crystallizing at very low, or systematically varied, ionic strength to investigate field-driven and other assembly mechanisms. Kinetics of the system may also be controlled by varying internal pressure, as can the level of supersaturation, etc. The aging of crystals at mild pH favors more thermodynamically stable final structures, resulting in intensification of observed morphological effects of the growthdirecting additives when compared with other techniques, and the final crystals can partially be dissolved by selecting a low final pH to reveal internal structural features. Acknowledgment. We thank the CNRS, CEA, and DFG for financial support of this project within the German-French Network. Th. Goldschmidt AG, Essen, is thanked for generously supplying the PEO-b-PMAA. References (1) Lahav, M.; Leiserowitz, L. Chem. Eng. Sci. 2001, 56, 2245-2253. (2) Estroff, L. A.; Hamilton, A. D. Chem. Mater. 2001, 13, 3227-3235. (3) Co¨lfen, H.; Mann, S. Angew. Chem., Int. Ed. 2003, 42, 2350-2365.

Page and Co¨lfen (4) (5) (6) (7) (8) (9) (10) (11) (12) (13) (14) (15) (16) (17) (18)

(19) (20) (21) (22) (23) (24) (25) (26) (27) (28) (29) (30) (31) (32) (33) (34) (35) (36) (37)

Meldrum, F. C. Int. Mater. ReV. 2003, 48, 187-224. Yu, S. H.; Co¨lfen, H. J. Mater. Chem. 2004, 14, 2124-2147. Co¨lfen, H. Curr. Opin. Colloid Interface Sci. 2003, 8, 23-31. Penn, R. L.; Banfield, J. F. Geochim. Cosmochim. Acta 1999, 63, 1549-1557. Penn, R. L. J. Phys. Chem. B 2004, 108, 12707-12712. Polleux, J.; Pinna, N.; Antonietti, M.; Niederberger, M. AdV. Mater. 2004, 16, 436-439. Co¨lfen, H.; Antonietti, M. Angew. Chem., Int. Ed. 2005, 44, 55765591. Wohlrab, S.; Pinna, N.; Antonietti, M.; Co¨lfen, H. Chem. Eur. J. 2005, 11, 2903-2913. Zhan, J.; Lin, H.-P.; Mou, C.-Y. AdV. Mater. 2003, 15, 621-623. Gehrke, N.; Co¨lfen, H.; Pinna, N.; Antonietti, M. Cryst. Growth Des. 2005, 5, 1317-1319. Wang, T.; Antonietti, M.; Co¨lfen, H. J. Am. Chem. Soc. 2005, 127, 3246-3247. Wang, T.; Antonietti, M.; Co¨lfen, H. Chem. Eur. J. 2006, in press, DOI: 10.1002/chem.200501019. Co¨lfen, H.; Qi, L. Chem. Eur. J. 2001, 7, 106-116. Reinecke, A.; Page, M.; Co¨lfen, H.; Smarsly, B.; Do¨bereiner, H.-G., unpublished results. Busch, S.; Dolhaine, H.; DuChesne, A.; Heinz, S.; Hochrein, O.; Laeri, F.; Podebrad, O.; Vietze, U.; Weiland, T.; Kniep, R. Eur. J. Inorg. Chem. 1999, 1643-1653. Addadi, L.; Moradian, J.; Shay, E.; Maroudas, N. G.; Weiner, S. Proc. Natl. Acad. Sci. U.S.A. 1987, 84, 2732-2736. Neira, A.; Fernandez, M. S.; Reteurt, J.; Arias, J. L. Mater. Res. Soc. Symp. Proc. 2004, EXS-1, H6.18.11. Kojima, Y.; Sadotomo, A.; Yasue, T.; Arai, Y. J. Ceram. Soc. Jpn. Intl. Ed. 1992, 100, 1128-1135. Kitano, Y.; Park, K.; Hood, D. W. J. Geophys. Res. 1962, 67, 4873. Rudloff, J.; Co¨lfen, H. Langmuir 2004, 20, 991-996. Zhong, Q. P.; Matijevic, E. J. Mater. Chem. 1996, 6, 443-447. Sedlak, M.; Antonietti, M.; Co¨lfen, H. Macromol. Chem. Phys. 1998, 2, 247. Wada, N.; Yamashita, K.; Umegaki, T. J. Colloid. Interface Sci. 1998, 201, 1-6. Becker, A.; Becker, W.; Marxen, J. C.; Epple, M. Z. Anorg. Allg. Chem. 2003, 629, 2305-2311. Tomson, M. B.; Nancollas, G. H. Science 1978, 200, 1059-1060. Koutsoukos, P.; Amjad, Z.; Tomson, M. B.; Nancollas, G. H. J. Am. Chem. Soc. 1980, 102, 1553-1557. Wang, L.; Sondi, I.; Matijevic, E. J. Colloid Interface Sci. 1999, 218, 545-553. Sondi, I.; Matijevic, E. J. Colloid Interface Sci. 2001, 238, 208214. Lin, R. Y.; Zhang, J. Y.; Zhang, P. X. J. Cryst. Growth 2002, 245, 309-320. Kontrec, J.; Kralj, D.; Brecevic, L.; Falini, G.; Fermani, S.; NoethigLaslo, V.; Mirosavljevic, K. Eur. J. Inorg. Chem. 2004, 4579-4585. Wickman, H. H.; Korley, J. N. Nature 1998, 393, 445-447. Co¨lfen, H.; Antonietti, M. Langmuir 1998, 14, 582. Reinecke, A.; Co¨lfen, H. MPI, Potsdam, Germany, unpublished data. Kitamura, M.; Konno, H.; Yasui, A.; Hirokatsu, M. J. Cryst. Growth 2002, 236, 323-332.

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