Crystallization of CaCO3 in the Presence of Ethanolamine Reveals

Oct 3, 2014 - The mechanism of calcite formation in the presence of ethanolamine (ETH) influences the polymorphism, morphology, and size distribution ...
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Crystallization of CaCO3 in the Presence of Ethanolamine Reveals Transient Meso-like Crystals Irene Bonacini,† Silvia Prati,† Rocco Mazzeo,*,† and Giuseppe Falini*,‡ †

Microchemistry and Microscopy Art Diagnostic Laboratory, University of Bologna, Via Guaccimanni 42, 48121 Ravenna, Italy Department of Chemistry “Giacomo Ciamician”, Alma Mater Studiorum - University of Bologna, via Selmi 2, 40126 Bologna, Italy

‡

S Supporting Information *

ABSTRACT: This paper presents a study on the influence of ethanolamine (ETH), an anticorrosion agent, on the polymorphism, morphology, and size distribution of calcium carbonate crystals after room-temperature synthesis. The pH of the solution, the presence of the additive, and the time of crystallization are proved to be important parameters for the morphogenesis of calcium carbonate. Morphologies as diverse as spherulite, rhomobohedral, and scalenohedral micrometric crystals, as well as submicrometric platelike particles, are produced. The mechanism of calcite formation in the presence of ETH involves a mesocrystal transient phase that evolves into classical single crystals. The key parameter of this process is supposed to be the ETH coverage of the primary calcite particles. In particular, it was found that, in particular conditions in the presence of ETH, nanoaggregates of calcium carbonate are formed, which are not stable and, with time, tend to convert into classical rhombohedral crystals of bigger dimensions. This study is relevant for the understanding of basic processes of calcium carbonate precipitation and has a potential applicative impact in the field of restoration when nanomaterials able to penetrate into the porosity of the materials are needed with the aim of reducing the acidity such as on corroded metallic objects or for the deacidification of paper supports. vaterite to calcite. On the other hand, Chen et al.8 observed that pH was the most important factor affecting the polymorphism of CaCO3, and they produced nearly pure vaterite at pH below 8. Besides, mainly vaterite was precipitated in the form of agglomerates of spherical particles by Hostomsky and Jones9 at pH 9.5, and pure vaterite was precipitated by Kralj and Brecevic10 at pH between 9.3 and 9.9. On the other hand, Han et al.11 and Stocks-Fisher et al.12 reported that pH values around 9 induced the heterogeneous deposition of calcite crystals. Many studies show that the addition of several organic additives can influence the morphology, the crystallinity, and the stability of the polymorphic species of calcium carbonate.13−19 In this context, the aim of the present paper is to test the use of ethanolamine (ETH) as an additive to increase pH and to control CaCO3 polymorphism and morphology This compound is used as feedstock in the production of detergents, emulsifiers, polishes, pharmaceuticals, corrosion inhibitors, and chemical intermediates. As corrosion inhibitor, it was tested on different materials, such as reinforced concretes,20 nickel,21 copper, and copper alloys.22 ETH has already been used by Vucak et al.23−25 as an additive in calcium carbonate

1. INTRODUCTION The controlled synthesis of inorganic materials with a specific size and morphology is an important aspect in the development of new materials in many fields. In the past decade, attention has been paid to the synthesis of calcium carbonate (CaCO3) particles due to their significance as a biomineral and their wide applications in industrial fields such as filler for paper, rubber, plastics, paint, etc.1 The morphologies of CaCO3 crystals are generally classified as rhombic calcite, needle-like aragonite, and spherical vaterite; calcite is the most stable phase at room temperature under normal atmospheric conditions, whereas aragonite and vaterite are metastable polymorphs that readily transform into the stable phase calcite. CaCO3 has been intensely studied to understand how crystal polymorphism could be controlled by the preparation conditions, such as reactant concentration, temperature, and addition of organic additives and pH. The precipitation of CaCO3 has been widely investigated,2−6 and even if pH was usually considered to be an important factor affecting this process, just a few reports on the influence of pH on the polymorphism of CaCO3 in conditions of high supersaturation have been so far published. These papers sometimes report ambiguous and contradictory results, probably due to the fact that different crystallization methods were employed. For example, Spanos and Koutsoukos7 concluded that pH had no effect on the transformation of © 2014 American Chemical Society

Received: July 28, 2014 Revised: October 1, 2014 Published: October 3, 2014 5922

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precipitation in order to obtain high-purity calcium carbonate from carbonation of milk lime with carbon dioxide. Vucak et al. employed a mixing method in which ETH maintained the pH of the solution high enough to sustain the precipitation of metal hydroxides. Here, a homogeneous precipitation is carried out with completely diverse starting conditions of precipitation.

2. EXPERIMENTAL SECTION 2.1. Materials. Analytical grade calcium chloride bihydrate and monoethanolamine were obtained from Sigma-Aldrich; sodium carbonate was purchased from Labochimica. Milli-Q water (resistivity 18.2 MΩ cm at 25 °C; filtered through a 0.22 μm membrane) was used for the preparation of aqueous solutions; all chemicals were used without further purification. 2.2. Synthesis. The precipitation of CaCO3 was carried out in a Teflon vessel at room temperature (about 22 °C). Aqueous solutions of CaCl2 (0.5 M) with or without monoethanolamine (4.73 M) and of Na2CO3 (0.5 M) were prepared and mixed together under vigorous stirring by using a magnetic stirrer. The mixture was left under stirring for different times: 1 min, 1 h, and 24 h, closed with Parafilm to avoid interactions with atmospheric CO2, and then filtered and dried in a 105 °C drying oven. The same set of synthesis was carried out adjusting the pH at about 12 by adding some drops of 4 M NaOH. 2.3. Characterization. The pH values were determined with an inoLAB pH/cond 720 pHmeter; the error range of the analytical method for pH was ±0.02. The resulting CaCO3 precipitates were characterized by scanning electron microscopy (SEM) on an Inspect S50 (FEI company). X-ray powder diffraction patterns were recorded on a PanAnalytical X’Pert Pro equipped with an X’Celetrator detector powder diffractometer using Cu Ka radiation generated at 40 kV and 40 mA. The diffraction patterns were collected within the 2θ range from 10° to 60° with a step size (Δ2θ) of 0.05° and a counting time of 109 s. Infrared spectra were collected by using a Thermo Nicolet Avatar 370 FTIR spectrometer on KBr pellets, working in the range of 4000−400 cm−1, at a spectral resolution of 4 cm−1 and 64 scans for each acquisition. An estimation of additive content on the CaCO3 particles was determined by thermogravimetric analysis (TGA) on an STD Q600 simultaneous thermal analysis instrument (TA Instruments). The analysis was performed under a nitrogen flow from 30 to 600 °C with a heating rate of 10 °C min−1. The quantitative XRD phase analysis was carried out applying the Rietveld method using the software Quanto26. The particle average size distribution was calculated from SEM images employing SigmaScan Pro 5 software, measuring 50 crystals at least. Moreover, the specific surface area and the particle size distribution of three calcite/ETH samples were measured with a MASTERSIZER 2000 (Malvern Instrument).

Figure 1. Changes in pH during three sets of CaCO3 precipitation experiments: (blue) Precipitation in the absence of additives. (green) Precipitation at pH 12. (red) Precipitation in the presence of ETH.

The SEM observations showed that, after 1 min of crystallization, the majority of CaCO3 particles appeared as aggregates of nanospherulitic crystallites in spherical or cauliflower-like shapes typical of the presence of vaterite2 (Figure 2a). Some rhombohedral crystals were also observed, which indicated the concurrent formation of a few nuclei of calcite. The average crystal size distribution varied from 3.5 to 6.3 μm (Table 1). FTIR spectroscopy and XRD analyses showed that vaterite and calcite precipitated (Figure 1SIa,d, Supporting Information). In the FTIR spectra, calcite was identified by the three main diagnostic bands, ν3 (1420 cm−1, asymmetric stretching of the carbonates), ν2 (875 cm−1, out-plane bending of carbonates), and ν4 (712 cm−1, in-plane bending of the carbonates), and two other minor bands at 1798 and 2513 cm−1.27 Distinction between calcite and vaterite in FTIR is based on the ν3 band at 1464 cm−1, which is wider for vaterite, and on the ν4 band, which for calcite is at 712 cm−1 and for vaterite at 744 cm−1.28 In the diffraction pattern, vaterite shows three intense peaks at 2θ = 24.8, 27.0, and 32.7° corresponding to the (110), (111), and (112) planes, respectively, and less intense peaks centered at 2θ = 43.98, 49.19, and 50.12° corresponding to (200), (202), and (114) planes (JCPDS: 741867).29 For calcite, the most intense peak is at 2θ = 29.4°, assigned to the (104) plane, plus other peaks centered at 23, 35.9, 39.4, 43.15, 47.5, and 48.5° that corresponded to (012), (110), (113̅), (202), (018), and (116̅) planes, respectively (JCPDS: 47-1473).30 The quantitative XRD phase analyses revealed that the precipitate contained 92.5 ± 0.5% of vaterite and 7.5 ± 0.5% of calcite (Table 1, Figure 4SIa1, Supporting Information). After 1 h of crystallization, in the precipitate well-defined rhombohedral particles in a mixture with spherical particles were observed (Figure 2b); the average crystal size ranged from 4 to 9 μm (Table 1). FTIR spectroscopy identified calcite and vaterite (Figure 1SIb), and accordingly, XRD analyses showed the diffraction peaks of vaterite and calcite (Figure 1SIe). The quantitative XRD phase analysis allowed us to calculate 80.4 ± 0.5% of vaterite and 19.6 ± 0.5% of calcite (Table 1, Figure 4SIa2). Reaching 24 h of crystallization, only well-defined rhombohedral particles, sometimes aggregated, were observed in the SEM images (Figure 1c), with a particle size distribution from 5.7 to 12.6 μm (Table 1). In the FTIR spectrum, only the three main bands of calcite, ν3, ν2, and ν4, were observed, and

3. RESULTS Three sets of crystallization experiments were carried out with the aim of understanding the role of ETH on the homogeneous precipitation of CaCO3: in the absence of additives (section 3.1), adding NaOH to keep the pH equal to 12 during the whole crystallization process (section 3.2), and adding 4.73 M ETH, which gives a starting pH of 12 (section 3.3). 3.1. Precipitation of CaCO3 in the Absence of Additives. These precipitation trials were carried out in order to study the phase and size distribution of CaCO3 particles, shape, and morphology and to follow the pH variation during the precipitation process from 1 min up to 24 h. The pH change during the precipitation process was measured (Figure 1): after 1 min of reaction, a sharp reduction of pH from 9.6 ± 0.1 to 8.1 ± 0.1 was recorded. Then, the pH remained constant after 1 h (pH = 8.1 ± 0.1) and 24 h of crystallization (pH = 8.0 ± 0.1) 5923

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Figure 2. SEM pictures of CaCO3 particles precipitated with the three different sets of crystallization: in the absence of additives at 1 min (a), 1 h (b), or 24 h (c); adding NaOH to keep the pH equal to 12 at 1 min (d), 1 h (e), or 24 h (f); and adding 4.73 M ETH, which gives a starting pH of 12 at 1 min (g), 1 h (h), or 24 h (i).

Table 1. CaCO3 Particles Precipitated via the Three Different Crystallization Methods synthesis CaCO3

CaCO3 pH 12

CaCO3/ETH

a

time 1 min 1h 24 h 1 min 1h 24 h 1 min 1h 24 h

synthesized powder amount (g)

crystals size distribution (μm)

± ± ± ± ± ± ± ± ±

3.5−6.3 4−9 5.7−12.6 1.8−5.3 2.7−5.6 4.5−6.9 0.65−2.6a 0.21−0.64a 3.6−7.1a

6.9 6.8 6.8 7.3 7.3 7.5 7.0 7.5 7.3

0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3

XRD quantitative analysisb 92.5% V 80.4% V 100% C 68.8% V 41.6% V C C C C

+ 7.5% C + 19.6% C + 31.2% C + 58.4% C

TGA weight loss 1.7 ± 0.1% 1.3 ± 0.1% 0.9 ± 0.1% 1.5 ± 0.1% 1 ± 0.1% 1.4 ± 0.1% 2.2 ± 0.1% 1.4 ± 0.1% 1.8 ± 0.1%

Value of crystallite distribution. bV and C indicate vaterite and calcite, respectively.

the diffraction pattern showed only the peaks due to calcite (Figure 1SIc,f). The quantitative XRD analysis confirmed that only calcite was present in the precipitate (Table 1). The precipitate obtained after 24 h was also analyzed employing the FTIR grinding curves method.31 This methodology, applicable only in the presence of a pure phase, allows having an estimation of the atomic order. The achieved information related to the crystallinity of a mineral can be used to differentiate the biogenic minerals from the geogenic ones. Among biogenic samples, this parameter provides a qualitative estimation of the interaction between the organic matrix and the mineral phase, thus measuring the interaction between the two components. In the present case, calcite was obtained in

the absence of additives, and the grinding curve overlapped with the one obtained for geogenic calcite (Figure 3). The weight loss in the range of 150−410 °C, probably ascribable to release of water, is about 1.5%. Moreover, an endothermic peak at about 423 °C, associated with the thermal transition from vaterite to calcite, was revealed by the heating flow profile both for 1 min and 1 h crystallization products. 3.2. Precipitation of CaCO3 at pH 12. This set of precipitation experiments was carried out adjusting the pH at about 12, by continuously adding a few drops of a 4 M NaOH solution. The crystallization was then performed up to 24 h with an almost constant pH (Figure 1). After 1 min, 1 h, and 24 h of crystallization, the morphologies of calcium carbonate 5924

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M ETH, which is the highest concentration below the solubility limit. In the presence of 4.73 M ETH, the pH during the precipitation time (Figure 1) was almost constant and showed a small decrease from the beginning of the reaction (12.0 ± 0.1) to end of the process (11.8 ± 0.1). After 1 min of crystallization, the particles were mainly constituted by a quite regular assembly of rhombohedral crystallites of micrometric and submicrometric sizes (Figure 2g). Some rhombohedral crystals with well-defined habits were also present. The specific surface area was 0.648 m2/g, the average size of crystallites ranged from 650 nm to 2.6 μm (Table 1), and the average size of particles was about 10.36 μm (Figure 5SI, Table 1SI, Supporting Information). From the FTIR spectra and the XRD patterns, calcite was identified as the only phase present (Figure 3SIa,d, Supporting Information). SEM observations showed (Figure 2h) that, after 1 h, the only morphology present was the regular assembly of rhombohedral crystallites observed after 1 min. However, these crystallites appeared as nanoaggregates, having an average size from 210 to 640 nm. The average size of aggregates was about 3.3 μm, and the specific surface area, 2.03 m2/g (Figure 5SI, Table 1SI), was greater than that observed for particles obtained after1 min. XRD and FTIR analyses revealed only the presence of calcite (Figure 3SIb,e). Rhombohedral particles, with well-defined growth planes, sometimes aggregated, were observed by SEM after 24 h (Figure 2i). The nanoaggregate structures, observed after crystallization for 1 h, disappeared, and the size distribution was in the range of 3.6−7.1 μm for the crystallites and showed an average value of 16.83 μm for the aggregates (Table 1, Figure 5SI, Table 1SI). The specific surface area was of 0.401 m2/g. The FTIR spectrum and the XRD pattern allowed detecting only calcite (Figure 3SIc,f). The precipitates obtained at different times were also analyzed employing the FTIR grinding curve method, as illustrated in Figure 3. The curves related to calcite obtained after 1 min and 1 h of crystallization in the presence of ETH have the same slope, which is higher than the one observed for geogenic calcite and for the other synthesized calcite powders. It can be noted that the position of the curve achieved just considering the points obtained on the 1 h ETH calcite is shifted in the left lower part of the graphic with respect to the geogenic curve. Thermogravimetic analysis showed in the range around 150−410 °C a weight loss of 2.2% for the product after 1 min. This value cannot be justified only by water lost, according to the previous experiments, and can be partially attributed to absorbed and/or adsorbed ETH in the calcite crystals. The weight lost tends to decrease after 1 h and at 24 h of crystallization at values that are not significantly different from the ones observed for the controls. However, the morphology of the particles and the grinding curves suggested that a residual content of ETH absorbed and/or adsorbed in the calcite crystals under the TGA limit of detection (less than 0.1%) has to be present at least on the particles obtained after 1 h of crystallization. What is relevant for the discussion on the mechanism of formation is that, for sure, a decrease of the content of ETH with the crystallization time occurs.

Figure 3. Plot of the ν2 versus ν4 peak heights of geogenic calcite, calcite synthesized at 24 h in the absence of additives (CaCO3) and adding NaOH (CaCO3 pH 12) and calcite/ETH at 1 min, 1 h, and 24 h crystallization (CaCO3/ETH), after each spectrum was normalized to the corresponding ν3 peak height. For each type of calcite, data points correspond to successive grindings of the same specimen.

particles (Figure 2d−f) were similar to those observed in section 3.1. Spherical particles composed of aggregated spherulitic nanocrystallites were the major items formed after 1 min of crystallization. Some rhombohedral crystals with well-defined growth planes were also present (Figure 1d). The average size distribution varied from 1.8 to 5.3 μm (Table 1). The FTIR spectrum indicated that the product obtained after 1 min of crystallization was composed of vaterite and calcite (Figure 2SIa, Supporting Information). XRD qualitative and quantitative analysis confirmed the presence of vaterite (68.8 ± 0.5%) and of calcite (31.2 ± 0.5%) (Figure 2SId, Supporting Information, Table 1, Figure 4SIb1). At 1 h of crystallization, rhombohedral particles with welldefined growth planes, in a mixture with a few spherical particles, were observed by SEM (Figure 1e) with an average size between 2.7 and 5.6 μm (Table 1). The FTIR spectroscopy XRD analysis identified the copresence of calcite and vaterite. Quantitative analyses based on XRD data quantified 41.6 ± 0.5% of vaterite and 58.4 ± 0.5% of calcite (Table 1, Figure 2SIb,e, Figure 4SIb2). After 24 h of crystallization, the spherical particles disappeared and only rhombohedral crystals with pronounced growth planes, and an average distribution from 4.5 to 6.9 μm, were observed in the SEM images (Figure 1f, Table 1). The FTIR spectrum and the X-ray diffraction pattern revealed only the presence of calcite (Figure 2SIc,f; Table 1). The latter precipitate showed an FTIR grinding curve (Figure 3) that moved lightly to the left with respect to the one related to geogenic calcite. Thermogravimetic profiles showed in the range around 150− 410 °C a water lost of about 1.5% for all the samples, as previously observed. Moreover, endothermic peaks at about 441 and 410 °C were observed in the heating flow profile for the products after 1 min and 1 h of crystallization, due to the thermal conversion of vaterite to calcite. 3.3. Precipitation of CaCO3 in the Presence of ETH. Crystallization of CaCO3 was performed in the presence of 4.73 5925

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4. DISCUSSION In this research, effects of ETH on the crystallization of calcium carbonate were studied in conditions of high supersaturation (4.9 saturation index (SI) of calcite calculated by Visual MINTEQ). This condition allows the nucleation of all polymorphs of calcium carbonate.32 ETH is a strong base that kept the pH values always around 12 during the whole crystallization process; this allowed the speciation of inorganic carbon in strong favor of carbonates. To discriminate the role of ETH as pH enhancer or crystal growth modifier, an additional two sets of crystallization trials were carried out: one at constant pH 12 by continuously adding NaOH and one in the absence of additive, ETH or NaOH. The use of ETH in the crystallization of calcium carbonate is not new. ETH was used as an additive in CaCO3 precipitation by Vucak et al.23−25 In these experiments, a heterogeneous precipitation system involving three phases (gas, liquid, and solid) was employed. At 30 °C, ETH seemed to have no effects on either the morphology or the polymorphism of CaCO3. In particular, vaterite grew initially and then transformed into calcite. In the present research, the control experiment performed without ETH maintaining pH at 12 during the whole crystallization process showed that the shape and morphology of the CaCO3 precipitates did not change with respect to the experiment in which pH was not controlled (Figure 2). Only the polymorphic distribution showed a small increase of the calcite content at pH 12. This is in agreement with some previous reports showing that pH in conditions of high supersaturations is not a key parameter for the evolution of calcium carbonate precipitation process along the time.11,33 Indeed, in both the controls performed without ETH, the systems precipitated first a mixture of vaterite and calcite that, with time (after 24 h), was converted completely into calcite. Han et al.33 described the transformation of vaterite to calcite as a classical two-step crystallization process: first, the dissolution of vaterite and then the crystallization of calcite, which was the rate-determining step of the overall transformation.34 Here, it was observed that the high pH favored the precipitation of calcite. Thus, according to the two-step crystallization process, in conditions of high pH, the crystallization of calcite is favored. The effect of the pH on the aggregation and shape of the crystalline particles has been poorly investigated in the literature.35 At constant pH 12, spherical particles in a mixture with rhombohedral ones were observed after 1 min, whereas only rhombohedral and scalenohedric crystals appeared after 24 h. By comparison of the control experimentsone with the final pH at about 8 and the other with a constant pH 12we observed that pH did not influence the morphology, contrary to what was observed by other authors reporting an increase of crystals with an irregular morphology as the pH became more basic.35 Moreover, the two control experiments showed similar particle size distributions. In the presence of ETH, only calcite precipitated, independently from the crystallization time. At 1 min and 1 h of crystallization, the calcite particles appeared as regular assemblies of micro- and nanocrystallites, respectively, the latter showing a platelike shape. These assembled particles were not observed in the absence of ETH or at pH 12. After 24 h of crystallization, only calcite crystals similar to those obtained in the control experiments were observed. The content of ETH adsorbed and absorbed in calcite crystals decreased from 1 min

to 1 h crystallization time. Considering a maximum intake of ETH on the 1 min obtained particles of about 0.6% in weight, we can calculate a molar ratio of ETH/CaCO3 of 9.4%, which agrees with the significant interaction during the crystallization process. The specific surface area drastically increased from 1 min to 1 h, and subsequently drastically decreased from 1 to 24 h. An opposite trend was observed in the average size of the particles. All of this information suggested that the initial particles formed after 1 min, were made of assembled crystallites, and underwent a disassembly process after 1 h, which produced smaller particles with higher specific surface areas. The latter were instable and, with time, transformed into classical rhombohedral crystals. The particles obtained after 1 min of crystallization could be classified as mesocrystallike,27−29 since they were birefringent (Figure 6SI, Supporting Information), even if not perfectly, and showed a superficial patterned roughness representative of primary building blocks. A schematic representation of the proposed mechanism is reported in Figure 4.

Figure 4. Schematic representation of mechanisms of growth for single crystals of calcite via a classical pathway or via a mesocrystal assembly process.

Results obtained using the “grinding curve” method supported the proposed mechanism31,36−39 (Figure 3). The distance of a synthetic calcite grinding curve from the grinding curve of a geogenic calcite was a measure of the reduction of the crystalline perfection of the crystals, which was generally associated with the presence of entrapped additives. In this case, the 1 min and 1 h calcite/ETH composites grinding curves were far from the geogenic one and overlapped. Also, the position of the 24 h calcite/ETH composites grinding curve, closer to the geogenic one, agreed with the proposed mechanism, which implied a partial ripening process of the primary building blocks. The grinding curves showed also that the particles precipitated after 1 h were characterized by lower values of the ν2/ν3 and ν4/ν3 ratios, when compared to those for 1 min 5926

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particles, due to the intensification and narrowing of the ν3 band. Moreover, the shift of the ν3 band from 1420 to 1447 cm−1 was evident (Figure 5).

crystals of bigger dimensions, through the three-dimensional attachment mechanism.

5. CONCLUSION This work has shown that ETH favored formation of calcite since the early stages of the precipitation process by stabilization of primary calcite particles (crystallites). Moreover, it has underlined that a highly basic pH has no influence on the shape and morphology of CaCO3 particles, in conditions of high supersaturation. Most importantly, this study has shown that, in the presence of ETH, the formation of “classical” single crystals of calcite occurs through a meso-like crystal transient phase mechanism in which the key parameter is the ETH coverage of the primary calcite particles. This research has also applicative implications, besides the fundamental ones on the crystallization mechanism reported above. Indeed, calcite nanoaggregates with a high superficial area have been obtained, which can be easily disaggregated after treatment with ultrasounds. A dispersion of this material may have a practical application in the restoration field for deacidification processes. Indeed, thanks to the particles’ dimensions, calcite crystals may penetrate into the porosity of the materials, neutralizing the excess of acidity such as on corroded metallic objects or on degraded paper supports.

Figure 5. FTIR spectra of particles synthesized adding 4.73 M ETH at 1 min (a), 1 h (b), and 24 h (c).

These unusual narrowing and the shift of the ν3 band were reported to be due to a reduction of the size of the mineral particles40,41 and agreed with the above observations and the proposed mechanism of crystallization. In the proposed mechanism, the interaction force among the primary particles, the crystallites, had a central role. These interactions changed with the ETH coverage of the crystallites. Completely covered crystallites were stable, whereas naked ones tended to aggregate to reduce their surface energy. Thus, it can be supposed that mesocrystal-like particles, the ones observed after 1 min, were formed by crystallites stabilized by ETH coverage, whereas, in “classical” rhombohedral single crystals, the ones obtained after 24 h, the degree of ETH coverage was very low. The partial coverage found a reasonable justification also considering crystallographic issues. Mesocrystals diffract as “classical” single crystals. This requires an almost perfect tridimensional crystallographic register of the primary building units. This process appears more reasonable when the interaction between crystalline faces is direct, similarly to a heterogeneous secondary nucleation event, and not mediated by an additive, which, even in the most likely case, should bring a certain degree of mismatch among primary particles. The above considerations fit with what is reported for a natural model of mesocrystals: the sea urchin spine. In this composite constituted by macromolecule/calcite material, the macromolecules cover the crystalline building block only partially, without forming continuous layers.42 Along the same lines, the reported existence of mineral bridges among the aragonitic tablets making the nacre implies a partial coverage on the primary building blocks.43,44 Finally, the key parameter of this proposed process was the ETH coverage of the primary calcite particles: as Figure 4 shows mesocrystals formed after 1 min which contain a certain amount of ETH underwent to a disassembly process after 1 h. Due to this transformation a reduction of ETH coverage occurred with the formation of smaller particles. These nanoaggregates have a high lattice energy due to the scarcity of ETH coverage, which led to the crystallographic fusion of the nanoaggregated building units to classical rhombohedral

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ASSOCIATED CONTENT

S Supporting Information *

FTIR spectra and XRD patterns of particles synthesized in the three different methods, Rietveld quantification plot figures of particles synthesized at basic pH (in the absence or in the presence of ETH), particle size distribution graphic, PLM figure, and table with particle dimensions and specific area of CaCO3/ETH crystals. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*Phone: +39 051 20 9 9484. E-mail: [email protected] (G.F.). *Phone: +39 0544/937160. E-mail: [email protected] (R.M.). Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Jaime Gómez Morales Laboratorio de Estudios Crystalográficos, IACT (CSIC-UGR), for the surface area measurements. G.F. thanks the Ministero degli Affari Esteri (MAE) for funding the Italy-Israel binational project CaFuMa.

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REFERENCES

(1) Wua, Z.; Davidson, J. H.; Francis, L. F. J. Colloid Interface Sci. 2010, 343, 176−187. (2) Martos, C.; Coto, B.; Pena, J. L.; Rodriguez, R.; Merino-Garcia, D.; Pastor, G. J. Cryst. Growth 2010, 312, 2756−2763. (3) Garcià-Carmona, J.; Morales, J. G.; Clemente, R. R. J. Cryst. Growth 2003, 249, 561−571.

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(41) Yue, L. H.; Shui, M.; Xu, Z. Spectrosc. Lett. 2001, 34 (6), 793− 802. (42) Berman, A.; Addadi, L.; Weiner, S. Nature 1988, 331, 546−548. (43) Cölfen, H.; Mann, S. Angew. Chem., Int. Ed. 2003, 42, 2350− 2365. (44) Meldrum, F. C.; Cölfen, H. Chem. Rev. 2008, 108, 4332−4432.

(4) Peric, J.; Vucak, M.; Krstulovic, R.; Brecevic, Lj.; Kralj, D. Thermochim. Acta 1996, 277, 175−186. (5) Wolthers, M.; Nehrke, G.; Gustafsson, J. P.; Van Cappellen, P. Geochim. Cosmochim. Acta 2012, 77, 121−134. (6) Meldrum, F. C. Int. Mater. Rev. 2003, 48, 187−224. (7) Spanos, N.; Koutsoukos, P. G. J. Cryst. Growth 1998, 191, 783− 790. (8) Chen, P. C.; Tai, C. Y.; Lee, K. C. Chem. Eng. Sci. 1997, 52, 4171−4177. (9) Hostomsky, J.; Jones, A. G. J. Phys. D 1991, 24, 165−170. (10) Kralj, D.; Brecevic, L. J. Cryst. Growth 1990, 104, 793−800. (11) Han, Y. S.; Hadiko, G.; Fuji, M.; Takahashi, M. J. Cryst. Growth 2006, 289, 269−274. (12) Stocks-Fisher, S.; Galinat, J. K.; Bang, S. S. Soil Biol. Biochem. 1999, 31, 1563−1571. (13) Yang, Q. F.; Liu, Y. Q.; Gu, A. Z.; Ding, J.; Shen, Z. Q. J. Colloid Interface Sci. 2001, 240, 608−621. (14) Andritsos, N.; Kontopoulou, M.; Karabelas, A. J.; Andritsos, N.; Kontopoulou, M.; Karabelas, A. J.; Koutsoukos, P. G. Can. J. Chem. Eng. 1996, 74, 911−919. (15) Ketrane, R.; Saidani, B.; Gil, O.; Leleyter, L.; Baraud, F. Desalination 2009, 249, 1397−1404. (16) Gopi, S. P.; Subramanian, V. K. Desalination 2012, 297, 38−47. (17) Yan, G.; Wang, L.; Huang, J. Powder Technol. 2009, 192, 58−64. (18) Yu, J.; Lei, M.; Cheng, B.; Zhao, X. J. Solid State Chem. 2004, 177, 681−689. (19) Colfen, H.; Qi, L. Chem.Eur. J. 2001, 7 (1), 106−116. (20) Ngala, V. T.; Page, C. L.; Page, M. M. Mater. Corros. 2004, 55 (7), 511−519. (21) Muralidharan, V. S.; Veerashanmugamani, M.; Paruthimal, G. Bull. Electrochem. 1985, 1, 373−375. (22) Golfomitsou, S.; Merkel, J. F. In Metal 2004. Proceedings of the International Conference on Metals Conservation, National Museum of Australia, Camberra, Australia, Oct 4−8, 2004. (23) Vucak, M.; Peric, J.; Krstulovic, R. Powder Technol. 1997, 91, 69−74. (24) Vucak, M.; Pons, M. N.; Peric, J.; Vivier, H. Powder Technol. 1998, 97, 1−5. (25) Vucak, M.; Peric, J.; Pons, M. N.; Chanel, S. Powder Technol. 1999, 101, 1−6. (26) Altomare, A.; Burla, M. C.; Giacovazzo, C.; Guagliardi, A.; Moliterni, A. G. G.; Polidori, G.; Rizzia, R. J. Appl. Crystallogr. 2001, 34, 392−397. (27) Adler, H. H.; Kerr, P. F. Am. Mineral. 1963, 48, 839−853. (28) Andersen, F. A.; Brecevic, L. Acta Chem. Scand. 1991, 45, 1018− 1024. (29) Meyer, H. J. Z. Kristallogr. 1969, 128, 183−212. (30) Effenberger, H.; Mereiter, K.; Zemann, J. Z. Kristallogr. 1981, 156, 233−243. (31) Suzuki, M.; Dauphin, Y.; Addadi, L.; Wainer, S. CrystEngComm 2011, 13, 6780−6786. (32) Ogino, T.; Suzuki, T.; Sawada, K. Geochim. Cosmochim. Acta 1987, 51, 2757−2767. (33) Han, Y. S.; Hadiko, G.; Fuji, M.; Takahashi, M. J. Mater. Sci. 2006, 41, 4663−4667. (34) Kralj, D.; Brecevic, L.; Kontrec, J. J. Cryst. Growth 1997, 177, 248−257. (35) Ma, Y. F.; Gao, Y. H.; Feng, Q. L. J. Cryst. Growth 2010, 112, 3165−3170. (36) Lane, M. D. J. Geophys. Res.: Planets 1999, 104, 14099−14108. (37) Regev, L.; Poduska, K. M.; Addadi, L.; Weiner, S.; Boaretto, E. J. Archaeol. Sci. 2010, 37, 3022−3029. (38) Gueta, R.; Natan; Addadi, L.; Wainer, S.; Refson, K.; Kronik, L. Angew. Chem., Int. Ed. 2007, 46, 291−294. (39) Poduska, K. M.; Regev, L.; Boaretto, E.; Addadi, L.; Weiner, S.; Kronik, L.; Curtarolo, S. Adv. Mater. 2011, 23, 550−554. (40) Yue, L. H.; Shui, M.; Xu, Z. J. Zhejiang Univ., Sci. 2000, 1, 178− 183. 5928

dx.doi.org/10.1021/cg501133n | Cryst. Growth Des. 2014, 14, 5922−5928