Synthesis and Characterization of Monodispersed Spheres of

Mar 30, 2005 - Parayil Kumaran Ajikumar,Ling Guan Wong,Gayathri Subramanyam,Rajamani Lakshminarayanan, andSuresh Valiyaveettil*. Singapore-MIT Allianc...
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Synthesis and Characterization of Monodispersed Spheres of Amorphous Calcium Carbonate and Calcite Spherules Parayil Kumaran Ajikumar,† Ling Guan Wong,‡ Gayathri Subramanyam,‡ Rajamani Lakshminarayanan,‡ and Suresh Valiyaveettil*,†,‡

CRYSTAL GROWTH & DESIGN 2005 VOL. 5, NO. 3 1129-1134

Singapore-MIT Alliance, Department of Chemistry, National University of Singapore, 3 Science Drive 3, Singapore 117543 Received November 21, 2004;

Revised Manuscript Received February 3, 2005

ABSTRACT: The isotropic property of amorphous calcium carbonate (ACC) is useful for the controlled synthesis of calcium carbonate based biomaterials. A simple and efficient strategy for the synthesis of monodispersed microspheres of ACC is reported using a low-temperature precipitation of calcium carbonate in the presence of magnesium ions. The room-temperature aging of this amorphous phase yielded superstructures of self-assembled calcite crystals. The stability of the precipitated ACC is proportional to an increase in the concentration of the magnesium ions and the precipitation time. The ACC with high magnesium content was stable for a one month period at room temperature under dry conditions. The ACC aged in solution was stable for up to 48 h. It is believed that the low temperature and the presence of magnesium ions facilitated the formation of stable monodispersed spheres of ACC. The morphological studies and characterizations were carried out using SEM, FTIR, XRD, electron diffraction, and Raman spectroscopy. We expect that the low-temperature synthesis of ACC followed by structural manipulation would be an efficient method for the controlled synthesis of interesting calcium-rich biomaterials. 1. Introduction The design and synthesis of calcium-rich materials under mild conditions has potential importance in terms of understanding the fundamental mechanisms in the biosynthesis of mineralized materials as well as for the development of new materials for biomedical applications.1 Many mineralization methodologies have been developed during the last two decades to mimic the natural mineralization process of hard tissue formation.2 In fact, specific aspects of the crystal formation such as crystal orientation, size, shape, and role of macromolecules as nucleators or templates in the biomineralization process have been under intense investigation.3-5 Calcium carbonate (CaCO3), one of the most abundant natural inorganic minerals, exists as three polymorphs (calcite, aragonite, and vaterite), two hydrated forms (monohydrocalcite and calcium carbonate hexahydrate) and an unstable amorphous form.6 Among these forms, the formation of stable amorphous calcium carbonate (ACC) is fascinating; its transformation into crystalline phase is not only thermodynamically favored but also kinetically fast.7 In recent years, research efforts have been focused on the effect of inorganic/ organic additives on crystallization processes to exploit the isotropic properties of ACC for morphological control over the crystalline phase.8 However, the formation and stabilization of ACC is still an active area of research owing to the presence of stable ACC in various biogenic minerals.6,9,10 Biogenic ACC contains substantial levels of magnesium ions in the mineral phase.11 It has been observed that the presence of magnesium is essential for the formation of the transient ACC phase.12 * To whom correspondence should be addressed. Tel: +65 68744327. Fax: +65 67791691. E-mail: [email protected]. † Singapore-MIT Alliance. ‡ Department of Chemistry.

The formation of ACC in an in vitro environment was first demonstrated by the rapid mixing of supersaturated solutions of CaCl2 and Na2CO3 at room temperature.13 Recently, a number of efforts have been made to understand the formation of the transient amorphous phase during the crystallization of CaCO3 with or without additives.14 However, all of these methods employed the rapid mixing of calcium chloride and sodium bicarbonate under turbulent conditions, and the precipitation occurs within a time scale of seconds, which lacks the detailed kinetic and stability investigation of the amorphous phase. Recently, Faatz et al. demonstrated a new route to the synthesis of ACC by the slow hydrolysis of dialkyl carbonate in calcium chloride solution without any additives.15 In another report, Loste et al. described the formation of calcite with constrained morphologies through the deposition of transient ACC from supersaturated solution at low temperature.16 Moreover, nano- or micron-sized spherical porous CaCO3 particles are potential candidates for the systems of biocompatible materials for drug delivery or scaffolds.17 Here we employ a slow and controlled precipitation of CaCO3 to form stable ACC with optimum properties. This paper describes the synthesis of monodispersed microspheres of stable ACC in the presence of magnesium at a low temperature and discusses the environment suitable for transformation of ACC into spherical self-assembled calcite crystals. 2. Experimental Section CaCO3 Precipitation Experiments. Analytical grade CaCl2, (NH4)2CO3, and MgCl2 were purchased from Merck and used as such without any purification. Deionized ultrapure water was used throughout the experiment. CaCO3 spheres were grown inside a closed desiccator via slow diffusion of CO2 released by the decomposition of ammonium carbonate crystals placed at the bottom of the desiccator.18 Typically, 1 mL of 7.5

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mM CaCl2 solution was introduced into the Nunc dish (6 × 4 wells) containing the cover slips, and the whole set up was covered with aluminum foil with a few pinholes at the top. Aliquots of 0.75 M MgCl2 solution were added to achieve a final Ca/Mg ratio of 1:3, 1:6, and 1:9 in the crystallization medium. The glass cover slips were carefully lifted from the precipitation media after 15, 24, or 48 h, rinsed gently with water, air-dried at room temperature, and then used for further characterization. For aging, samples in solution and dry conditions were kept at room temperature and monitored the phase transformation. Morphological Studies and Characterization. Scanning electron microscope (SEM) images were taken with a Philips XL-30 and JEOL JSM 6700 scanning electron microscope. The samples were carefully mounted on copper stubs with doublesided carbon tape and sputter coated with gold before examination. EDX analyses were performed at various locations on the substrate to confirm the presence of magnesium. Elemental analyses of the particles were carried out using inductively coupled plasma optical emission spectroscopy (ICP-OES) model Thermo Jarrel Ash Duo Axis Plasma. About 1 mg of the sample was digested using concentrated nitric acid for 30 min before making the stock solution in water. The calcium and magnesium stock solutions were diluted to give the required range of standard solutions and used for calibration of the machine before the actual sample was tested. The precipitated CaCO3 was characterized using the Bio-Rad FTS 165 FT-IR spectrophotometer. Small amounts of finely powdered materials were mixed with KBr powder to make a pellet for FTIR investigations. X-ray diffraction studies on the precipitated CaCO3 were done using a D5005 Siemens X-ray diffractometer with Cu-ΚR radiation at 40 kV and 40 mA. The particles grown on the glass slides were mounted on a plastic holder and diffraction studies were carried out. The phase identification was done by comparing the X-ray diffraction patterns of the crystals with the standard data available from Joint Committee on Powder Diffraction Standards. Selected area electron diffraction patterns were recorded on a JEOL JEM CX II machine at an applied voltage of 100 kV. The samples grown on a glass plate were collected on Formvar-coated copper grids and used for electron diffraction studies. For Raman spectroscopy studies, the air-dried samples on the glass plate were observed with a microscope (50×). The spectra were recorded for the samples for 10 s in the range of 100-1200 cm-1 using Jobin Yvon Horiba Raman imaging microscope.

3. Results and Discussion Preparation and Characterization of Spherical Amorphous Calcium Carbonate. The CaCO3 precipitation experiment was performed at 4 °C via slow diffusion of CO2 obtained from the decomposition of ammonium carbonate to 7.5 mM CaCl2 solution.18 Aliquots of a 0.75 M MgCl2 solution were added to achieve the final solution concentration ratio of Ca/Mg ) 1:3, 1:6, and 1:9 in the precipitation media. These ratios were chosen after conducting preliminary experiments with different Ca/Mg ratios, and it found that at a selected Ca/Mg ratio of 1:6, ca. 20 mol % magnesium ions were incorporated in the precipitated CaCO3. The precipitates obtained after 15, 24, and 48 h were removed and gently rinsed with water; air-dried at room temperature for 20 min, and characterized using scanning electron microscopy (SEM), energy-dispersive Xray scattering (EDXS), X-ray diffraction studies (XRD), Fourier transform infrared spectroscopy (FTIR), and Raman spectroscopy. For aging studies, the precipitate collected at various time intervals were kept at room temperature (25 °C), and the solid-solid transformation or morphological changes, if any, was monitored by IR, XRD, and SEM studies. The exact amount of magne-

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Figure 1. Representative electron micrographs of the CaCO3 particles grown at various Ca/Mg ratios (shown in left-hand corner) in the precipitation medium for 15 h at 4 °C. FTIR (A) and XRD (B) pattern of the CaCO3 particles grown in the presence of Mg2+ ions for 15 h. Table 1. The Observed mol % of Magnesium Ion Incorporated into the Precipitated Calcium Carbonate Collected on the Glass Plate after 15 h of Precipitation at 4 °C Was Determined Using Elemental Analysis Ca/Mg ratio in precipitation media

observed mol % of magnesium incorporation

1:3 1:6 1:9

13.3 ( 0.3 21.1 ( 0.2 32.4 ( 0.4

sium incorporated in the precipitate collected was measured using elemental analysis and summarized in Table 1. The results indicated that the percentage of Mg ions incorporated is independent of precipitation time and depends on the initial concentration of the medium. To understand the formation of stable ACC and phase transformation in solution with time, precipitation experiments were performed at 15, 24, and 48 h at 4 °C. Figure 1 shows the morphology of the CaCO3 particles formed after 15 h. In the absence of magnesium ions in the precipitation media, calcite crystals with rhombohedral morphology were observed with a mean size of 10 ( 5 µm (Figure 1). However, after the introduction of the magnesium ions into the medium (Ca/Mg ) 1:3), spherical ACC with an average size of ca. 1.5 ( 0.3 µm was formed (Supporting Information). EDXS on the surface of the spheres showed the presence of Mg and Ca on the surface (Supporting Information). As the concentration of magnesium ion increases in the crystallization medium, the surface of the spheres becomes smoother. FTIR spectra of the powdered samples

Monodispersed Spheres of Amorphous Calcium Carbonate

Figure 2. Representative electron micrographs of the CaCO3 particles grown at various Ca/Mg ratios in the precipitation medium for 24 h at 4 °C. Ca/Mg ratios are indicated in the figure. Micro-Raman spectra (A), FTIR (B), and XRD (C) of the CaCO3 spheres. Table 2. Mean Particle Size Distribution of Precipitated Calcium Carbonate after 15, 24, and 48 h of Precipitation at 4 °C observed mean size (µm) of ACC from different Ca/Mg ratios Ca/Mg ratio precipitation time (h)

1:3

1:6

1:9

15 24 48

1.5 ( 0.3 1.7 ( 0.3 1.4 ( 0.4

2 ( 0.2 1.7 ( 0.4 1.5 ( 0.2

1.0 ( 0.1 1.5 ( 0.3 1.7 ( 0.2

in the KBr matrix is shown in Figure 1B. For the control sample, the peaks at 713, 874, and 1430 cm-1 are characteristic of the calcite phase.19 Introduction of the magnesium ions in the precipitation media at various Ca/Mg ratios (i.e., 1:3, 1:6, or 1:9) resulted in a significant broadening of the peak at 862 cm-1 (ν2), splitting of the 1450 cm-1 (ν3) peak, and the appearance of a broad peak at 1070 cm-1 (ν1) (Figure 1A), all of which indicate the formation of ACC.8i,20 The presence of ACC was further confirmed by X-ray diffraction analyses (Figure 1B). In the absence of Mg2+ ions, diffraction peaks corresponding to the calcite lattice were observed. The lack of sharp peaks in the X-ray diffraction pattern of precipitates obtained from various Ca2+/Mg2+ concentration ratios confirmed the formation of ACC.8c After 24 h, spherical particles of ca. 1.7 ( 0.3 µm in size with a rough surface texture were observed at a Ca/Mg ratio of 1:3 (Table 2, Figure 2). The presence of the broad band at 1080 cm-1 (carbonate symmetric

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Figure 3. Representative electron micrographs of the CaCO3 particles grown at various Ca/Mg ratios in the precipitation medium for 48 h at 4 °C. FTIR spectra (A) of the CaCO3 particles from 48 h precipitation. Note the absence of a peak at 713 cm-1 at higher Ca/Mg ratios (1:6 and 1:9), indicating the existence of an amorphous phase. XRD patterns (B) with sharp diffraction peaks for a pure calcite phase at 1:0 and a small peak corresponds to magnesium calcite at 1:3.

stretching) in the Raman spectra, indicated the formation of ACC in the precipitates collected after 24 h (Figure 2A).21 In addition, a broad featureless hump around 150-300 cm-1 at higher magnesium concentrations (1:6 and 1:9) confirmed the existence of stable ACC.6,10 Further characterization of the ACC phase using FTIR and XRD analyses (Figure 2B,C) also gave consistent results. The absence of a sharp peak at 713 cm-1 in the FTIR spectra and a sharp diffraction pattern in the XRD indicates the presence of an amorphous phase. Analysis of the spheres formed after 48 h using SEM, FTIR, and XRD showed that the low concentration of magnesium ions in the precipitation medium (1: 3) resulted in a slight shift of the peaks compared to the control experiment with no magnesium ions present, indicating the formation of magnesium calcite (Figure 3). Further increasing the magnesium content in the precipitation media (1:6 or 1:9) resulted in a significant broadening of the peak at 862 cm-1 (ν2), splitting of the 1450 cm-1 (ν3) peak, and the appearance of a broad peak at 1070 cm-1 (ν1) (Figure 3A), all of which are characteristics of the ACC phase.20 The absence of sharp X-ray diffraction patterns at a high concentration of Mg2+ ions (1:6 or 1:9) further supports the formation of ACC (Figure 3B). The ACCs precipitated after 48 h appeared to be hollow (Figure 3 inset). A homogeneous and narrow particle size distribution was observed in the presence of high magnesium (1:9) in the precipitation media irrespective of the precipitation time (15, 24, and 48 h) (Table 2). Investigation of the Long-Term Stability and Transformation of ACC to Crystalline Phase. It is

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Figure 4. Micrographs of CaCO3 spheres obtained after 15 h precipitation with 48 h aging at room temperature. Ca/Mg ratios are indicated in the figure. FTIR (A) and XRD patterns (B) of the corresponding CaCO3 particles.

believed that first step in the precipitation of CaCO3 is the formation of an amorphous phase, which at room temperature converts to metastable vaterite followed by a stable calcite phase.14,22 The precipitates obtained from the 15, 24, and 48 h precipitation existed as ACC. The stability of this ACC and its transformation to crystalline phase was further investigated by aging the samples in a dry state and in solution at room temperature. The stability and morphological transformation were monitored at different time periods of 1 day, 2 days, 1 week, and 1 month, respectively. ACC from different precipitation times irrespective of the magnesium content is stable up to 2 days. Figure 4 represents the morphology and stability of the 2 days dry state aged samples of ACC obtained from a 15 h precipitation and shows that ACC is stable without any change in shape or size. A similar observation was seen from the samples obtained from the 24 and 48 h precipitates. However, during one week of aging at room temperature under dry conditions, the low magnesium content (1:3) amorphous phase was partially transformed into calcite, and the high magnesium content ACC is stable with no significant changes in morphology, size, or shape (Figure 5). The broadening of the FTIR peak at 862 cm-1 (ν2), splitting of 1450 cm-1 (ν3), and absence of any sharp peak in the XRD pattern confirmed the ACC phase. Similar results were observed over a period of one month (Figures 6 and 7). The SEM shows that the surface of the low magnesium content (1:3 and 1:6) samples are roughened, whereas the high magnesium content samples have a smooth surface. The XRD and FTIR spectral analysis confirmed the transformation of the low magnesium content ACC (1:3 and 1:6) to magnesium calcite, whereas the high magnesium content ACC samples are stable. The further analysis of the samples after six months showed the complete

Figure 5. Representative scanning electron micrographs of the CaCO3 particles obtained after 15 h (A-C) and 48 h (DF) precipitation time followed by one week aging at room temperature. The Ca/Mg ratios are indicated in the figure. FTIR spectra (G) of the CaCO3 particles grown in the presence of Mg2+ ions after 15 h precipitation and X-ray diffraction patterns (H) of the CaCO3 spheres collected after 48 h precipitation. In both cases, one week aging at room temperature was carried out before characterization.

transformation into the calcite phase (Supporting Information). During the aging in solution, the samples obtained with different magnesium content were stable up to 48 h; further aging studies showed that within a time period of one week, the precipitates obtained were partially transformed into the crystalline phase (Figure 8; Figure S6, Supporting Information). In summary, ACC formed in the presence of magnesium ions is stable for at least 48 h in solution and one month in the solid state at room temperature. Similar observation was reported by Loste et al. by the rapid mixing of sodium bicarbonate and calcium chloride solutions.8c Here also the presence of high amounts of magnesium ions yielded an ACC precursor phase which then transformed to crystalline form within 14 h through a series of metastable phases. Although the method of precipitation employed in our studies is

Monodispersed Spheres of Amorphous Calcium Carbonate

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Figure 8. XRD of the CaCO3 particles (from 15, 24, and 48 h precipitation) aged for 2 days in solution at room temperature.

Figure 6. Representative electron micrographs of the CaCO3 particles collected after 15 h (A-C) and 48 h (D-F) precipitation with one month aging at room temperature. Ca/Mg ratios used in the precipitation media are indicated in the figure.

content. It is expected that magnesium forms strong bonds with water molecules, and these hydrated spheres increase the degree of supersaturation in the medium, which results in the formation of amorphous phases. However, the low temperature and high magnesium content yielded more stable and highly monodispersed spherical particles with a hollow structure. At low temperatures and in the presence of a high magnesium ion content, formation of colloidal amorphous phase rather than the nucleation of crystalline phases was initiated. Such a colloidal phase is expected to minimize its surface contact with the surroundings through the formation of spherical structures. In addition to this, by aging the samples under a dry environment, we achieved the solid-to-solid transformation of ACC into assembled spherical superstructures of magnesium calcite without losing its shape or size from the amorphous precursor. It is interesting to note that we achieved control over the stability, size (micron to submicron), shape, and morphology of the magnesiumincorporated ACC by changing the Ca/Mg ratios and lowering the temprature. 4. Conclusion

Figure 7. FTIR and XRD of the CaCO3 particles obtained after 15 h (A and B) and 48 h (C and D) precipitation with one month aging at room temperature.

different, the results corroborate with the findings that ACC is stabilized in the presence of high magnesium

The results presented here demonstrate a simple route to synthesize highly monodispersed stable spherules of magnesium-incorporated ACC and the preparation of uniform spherical self-assembled magnesium calcites. In fact, the stability of ACC can be achieved by performing the precipitation at a low temperature and in the presence of a high amount of magnesium. The control over the morphology of the materials was achieved through the addition of Mg2+ ions and lowtemperature precipitation. The ACC obtained showed a hollow spherical structure for particles. Micron- or submicron-sized spheres of CaCO3 materials is appealing owing to many potential applications in bitechnology, biomedical sciences, chemical storage, and catalysis.17,23,24 Thus, the design and development of novel strategies for the preparation of such hollow spheres of CaCO3 is an active area of material research. This is the first paper describing the preparation of stable ACC using a low-temperature precipitation without any organic additives. Further studies toward the design and development of more simple methods for the production of submicron- or nanosized hollow spherical ACC are under progress. Acknowledgment. The authors thank the Singapore-MIT Alliance for financial support. We also

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acknowledge the technical support from the Department of Chemistry and Department of Materials Sciences, National University of Singapore. Supporting Information Available: EDXS data of calcium carbonate crystals grown with different Ca/Mg ratios (Figure S1); FTIR and XRD (Figure S2), representative electron micrographs (Figures S3 and S4), and XRD (Figures S5 and S6) of CaCO3 grown in the presence of Mg2+ ions. This material is available free of charge via the Internet at http:// pubs.acs.org.

References (1) (a) Weiner, S.; Addadi, L. J. Mater. Chem. 1997, 7, 689702. (b) Baeuerlein, E. Biomineralization: From Biology to Biotechnology and Medical Application; Wiley-VCH Publishers: New York, 2000. (c) Mann, S.; Web. J.; Williams, R. J. P. Biomineralization: Chemical and Biochemical Perspectives; VCH Publishers; Weinheim; 1989. (2) (a) Aizenberg, J.; Hendler, G. J. Mater. Chem. 2004, 14, 2066-2072. (b) Meldrum, F. C. Int. Mater. Rev. 2003, 48, 187-224. (c) Mann, S. Biomineralization: Principles and Concepts in Bioinorganic Materials Chemistry; Oxford University Press: New York, 2001. (3) (a) Mann, S.; Davis, S. A.; Hall, S. R.; Li, M.; Rhodes, K. H.; Shenton, W.; Vaucher, S.; Zhang, B. J. Chem. Soc., Dalton 2000, 21, 3753-3763. (b) Walsh, D.; Mann, S. Nature 1995, 377, 320-323. (b) Walsh, D.; Lebeau, B.; Mann, S. Adv. Mater. 1999, 11, 324-328. (4) (a) Co¨lfen, H. Curr. Opin. Colloid Interface Sci. 2003, 8, 2331. (b) Colfen, H.; Antonietti, M. Langmuir 1998, 14, 582589. (c) Rudloff, J.; Antonietti, M.; Colfen, H.; Pretula, J.; Kaluzynski, K.; Penczek, S. D. Macromol. Chem. Phys. 2002, 203, 627-635. (5) (a) Zhan, J. H.; Lin, H. P.; Mou, C. Y. Adv. Mater. 2003, 15, 621-623. (b) Yang, D.; Qi, L. M.; Ma, J. M. Chem. Commun. 2003, 1180-1181. (c) Meldrum, F. C.; Hyde, S. T. J. Cryst. Growth 2001, 23, 544-558. (d) Rautaray, D.; Kumar, P. S.; Wadgaonkar, P. P.; Sastry, M. Chem. Mater. 2004, 16, 988993. (e) Yu, S.-H.; Colfen, H. J. Mater. Chem. 2004, 14, 2124-2147. (f) Aizenberg, J. Adv. Mater. 2004, 16, 12951302. (6) (a) Addadi, L.; Raz, S.; Weiner, S. Adv. Mater. 2003, 15, 959-970. (b) Braga, D. Angew. Chem., Int. Ed. 2003, 42, 5544-5546. (7) (a) Clarkson, J. R.; Price, T. J.; Adams, C. J. J. Chem. Soc., Faraday Trans. 1992, 88, 243-249. (8) (a) Reddy, M. M.; Nancollas, G. H. J. Cryst. Growth. 1976, 35, 33-38 (b) Donners, J. J. J. M.; Heywood, B. R.; Meijer, E. W.; Nolte, R. J. M.; Roman, C.; Schenning, A. P. H. J.; Sommerdijk, N. A. J. M. Chem. Commun. 2000, 1937-1938. (c) Donners, J. J. J. M.; Heywood, B. R.; Meijer, E. W.; Nolte, Sommerdijk, N. A. J. M. Chem.sEur. J. 2002, 8, 25612567. (d) Loste, E.; Wilson, R. M.; Seshadri, R.; Meldrum, F. C. J. Cryst. Growth 2003, 254, 206-218. (e) Raz, S.; Weiner, S.; Addadi, L. Adv. Mater. 2000, 12, 38-42. (f) Kjellin, P.; Holmberg, K.; Nyden, M. Colloid Surf. A 2001,

(9)

(10)

(11) (12) (13) (14)

(15) (16) (17)

(18) (19) (20) (21) (22) (23)

(24)

194, 49-55. (g) Gower, L. B.; Odom, D. J. J. Cryst. Growth 2000, 210, 719-734. (h) Xu, G. F.; Yao, N.; Aksay, I. A.; Groves, J. T. J. Am. Chem. Soc. 1998, 120, 11977-11985. (i) Aizenberg, J.; Muller, D. A.; Grazul, J. L.; Hamann, D. R. Science 2003, 299, 1205-1208. (a) Raz, S.; Testeniere, O.; Hecker, A.; Weiner, S.; Luquet, G. Biol. Bull. 2002, 203, 269-274. (b) Becker, A.; Bismayer, U.; Epple, M.; Fabritius, H.; Hasse, B.; Shi, J. M. J. Chem Soc., Dalton Trans. 2003, 551-555. (a) Aizenberg, J.; Lambert, G.; Addadi, L.; Weiner, S. Adv. Mater. 1996, 8, 222-226. (b) Tsutsui, N.; Ishii, K.; Takagi, Y.; Watanabe, T.; Nagasawa, H. Zool. Sci. 1999, 16, 619628. Weiner, S.; Levi-Kalisman, Y.; Raz, S.; Addadi, L. Connect. Tissue Res. 2003, 44, 214-218. Raz, S.; Hamilton, P. C.; Wilt, F. H.; Weiner, S.; Addadi, L. Adv. Funct. Mater. 2003, 13, 480-486. Brecevic, L.; Nielsen, A. E. J. Cryst. Growth 1989, 98, 504510. (a) Rieger, J.; Thieme, J.; Schmidt, C. Langmuir 2000, 16, 8300-8305. (b) Bolze, J.; Peng, B.; Dingenouts, N.; Panine, P.; Narayanan, T.; Ballauff, M. Langmuir 2002, 18, 83648369. (c) Pontoni, D.; Bolze, J.; Dingenouts, N.; Narayanan, T.; Ballauff, M. J. Phys. Chem. B 2003, 107, 5123-5127. (d) Colfen, H.; Qi, L. M. ChemsEur. J. 2001, 7, 106-116. Faatz, M.; Grohn, F.; Wegner, G. Adv. Mater. 2004, 16, 9961000. Loste, E.; Meldrum, F. C. Chem. Commun. 2001, 901-902. (a) Sukhorukov, G. B.; Volodkin, D. V.; Gunther, A. M.; Petrov, A. I.; Shenoy, D. B.; Mohwald, H. J. Mater. Chem. 2004, 14, 2073-2081. (b) Volodkin, D. V.; Petrov, A. I.; Prevot, M.; Sukhorukov, G. B. Langmuir 2004, 20, 33983406. Albeck, S.; Aizenberg, J.; Addadi, L.; Weiner, S. J. Am. Chem. Soc. 1993, 115, 11691-11697. Andersen, F. A.; Brecevic, L. Acta Chem. Scand. 1991, 45, 1018-1024. Aizenberg, J.; Lambert, G.; Weiner, S.; Addadi, L. J. Am. Chem. Soc. 2002, 124, 32-39. (a) Weiss, I. M.; Tuross, N.; Addadi, L.; Weiner, S. J. Exp. Zool. 2002, 293, 478-491. (b) Li, M.; Mann, S. Adv. Funct. Mater. 2002, 12, 773-779. Sohnel, O.; Garside, J. Precipitation: Basic Principles and Applications; Butterworth-Heinemann: Oxford UK, 1992; p 145. (a) Zhong, Z.; Yin, Y.; Gates, B.; Xia, Y. Adv. Mater. 2000, 12, 206-209. (b) Hollow and Solid Spheres and Microspheres; Science and Technology Associated with Their Fabrication and Application; Wilcox, D. L., Sr., Berg, M., Bernat, T., Kellerman, D., Cochran, J. K., Jr., Eds.; MRS Proceedings, Pittsburgh, PA, 1995; Vol. 372. (c) Jiang, P.; Bertone, J. F.; Colvin, V. L. Science 2001, 291, 453-457. (d) Wong, M. S.; Cha, J. N.; Choi, K. S.; Deming, T. J.; Stucky, G. D. Nano Lett. 2002, 2, 583-587. (a) Byrn, S.; Pfeiffer, R.; Ganey, M.; Hoiberg, C.; Poochikian, G. Pharm. Res. 1995, 12, 945-954. (b) Hancock, B. C.; Parks, M. Pharm. Res. 2000, 12, 397-404.

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