Tuning the Incorporation of Magnesium into Calcite during Its

Article Cite This: Cryst. Growth Des. XXXX, XXX, XXX−XXX

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Tuning the Incorporation of Magnesium into Calcite during Its Crystallization from Additive-Free Aqueous Solution Jacinta M. Xto,†,‡ Huachuan Du,§ Camelia N. Borca,† Esther Amstad,§ Jeroen A. van Bokhoven,†,‡ and Thomas Huthwelker*,† †

Paul Scherrer Institut, 5232 Villigen, Switzerland Institute for Chemical and Bioengineering, ETH Zürich, 8093 Zurich, Switzerland § Soft Materials Laboratory, Institute of Materials, É cole Polytechnique Fédérale de Lausanne (EPFL), 1015 Lausanne, Switzerland Downloaded via UNIV OF SOUTHERN INDIANA on July 22, 2019 at 20:07:25 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

‡

S Supporting Information *

ABSTRACT: Under ambient conditions, marine organisms are able to synthesize a variety of functional materials, ranging from eye lenses to protective shells through the meticulous control over magnesium incorporation into calcite during its crystallization. The mechanistic understanding of how they achieve such exquisite control, at a constant magnesium-to-calcium ratio and at ambient conditions, is important in the development of bioinspired functional materials. However, the replication of these processes in the laboratory is still challenging. Herein, we present a systematic study on how to tune magnesium incorporation into calcite and polymorph selection in the Ca−Mg−CO3 system through the precise control of the inorganic solutions chemistry at ambient conditions of temperature and pressure, and at a magnesium-to-calcium ratio of 5:1, which is analogous to the ratio found in most seas. By varying the pH, cation-to-anion ratio, and solution concentration, the controlled synthesis of magnesium calcites with 10−45% magnesium was achieved at room temperature. The mechanism of formation is consistent with that observed during biomineralization, during which an intermediate magnesium-rich amorphous calcium carbonate (Mg-ACC) phase forms first and later transforms into high magnesium calcite. Once crystallization occurs, the magnesium calcites that form are stable in solution and exhibit slow growth through Ostwald ripening. Our findings suggest that the precise control of saturation levels is key in driving nucleation and crystallization.

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INTRODUCTION

Magnesite and calcite are isomorphs; hence, magnesium ions can substitute calcium in the calcite crystal lattice but not in the aragonite crystal lattice.10,11 Therefore, substitutional defects arising from inclusion of magnesium ions during crystallization from magnesium-rich solutions are inevitable during calcite crystallization but have no effect on aragonite. Recent ab initio calculations showed that the consequence of magnesium inclusion is an increase in the surface energy of the critical nucleus with increasing amounts of magnesium ions, resulting in the preferential precipitation of aragonite above a certain magnesium threshold in solution.12 According to classical nucleation theory, the Gibbs free energy ΔG of a growing crystal nucleus of radius r is given by eq 1,13 which shows that increased surface energy, γ, hinders nucleation

In nature, limited elements are utilized to achieve the formation of unique and complex functional materials.1,2 Marine organisms in particular have developed strategies for synthesizing functional materials such as skeletal supports,3 lenses,4 sensory organs, protective materials,5 and teeth6 by utilizing the freely available magnesium, calcium, and carbonate ions. Some of the remarkable properties displayed by these biominerals have been associated with their ability to control the incorporation of varying amounts of magnesium into calcite during biomineralization.6−9 The unique properties of the sea urchin tooth, for example, have been correlated to the diversity of its structural morphology based on the varying amounts of magnesium that are incorporated in the different structural elements ranging from 5% to 45% magnesium carbonate in the calcite crystals.6 Attempts to synthetically replicate these biomineralization processes with the aim of developing superior bioinspired functional materials remains challenging.2 A key challenge is the lack of insight into the mechanisms of such precise control over magnesium inclusion under ambient conditions and with constant magnesium-tocalcium ratios of 5:1, analogous to the ratio found in oceans. © XXXX American Chemical Society

ΔG = 4πr 2γ +

4 3 πr ΔGv 3

(1)

Equation 2 further shows that temperature (T) and saturation (S), which is the log ion activity product of the solution Received: February 7, 2019 Revised: May 22, 2019 Published: June 25, 2019 A

DOI: 10.1021/acs.cgd.9b00179 Cryst. Growth Des. XXXX, XXX, XXX−XXX

Crystal Growth & Design

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divided by log solubility constant Ksp, strongly influence the bulk free energy per unit volume (ΔGv), which drives nucleation k T ln(S) ΔGv = − B vm

Table 1. Summary of Experimental Conditions Time evolution Na2CO3 [M]

9.0 pH effect

(2) Na2CO3 [M]

where kB is the Boltzmann constant and vm is the molar volume of the monomer. Hence, to drive the system toward calcite nucleation in the presence of magnesium ions, it is necessary to increase either the temperature (T) or the saturation level (S) to overcome the increasing surface energy resulting from inclusion of magnesium ions At low temperature, precipitation of high magnesium calcite is additionally postulated to be kinetically limited due to the very strong hydration of Mg2+ compared to Ca2+ and CO32−,10,14 resulting in preferential precipitation of magnesium-free aragonite. Therefore, to form magnesium calcites under ambient conditions, most synthetic approaches rely on organic additives, such as dioxane/THF to reduce the strong hydration behavior of Mg2+, or employ a very high magnesium to calcium ratio.15−23 Inasmuch as these experiments have yielded high magnesium calcites, it remains challenging to control the magnesium content in the calcite crystals or to achieve control over the final polymorph. Furthermore, most of these experiments are conducted under conditions seldom seen in nature, and hence provide limited insights into biomineralization processes. Previous studies have shown that several marine organisms locally control their aqueous solution chemistry to be able to calcify. Foraminifera, for example, are able to produce calcite with very high magnesium content by elevating the intracellular pH at the calcification site leading to elevated levels of magnesium calcite saturation.8 On the other hand, corals increase the concentration of dissolved inorganic carbon at the calcifying site resulting in saturation levels favorable for aragonite precipitation.24 By mimicking these processes, we present a conceptual approach on how to directly tune polymorph selection and magnesium incorporation into calcite in the Ca−Mg−CO3 system directly from aqueous solutions and at ambient conditions. Inspired by the formation of biogenic magnesium calcites, the effect of three key variables, pH, carbonate ions (cation-to-anion ratio), and solution concentration, were systematically studied at a constant magnesium-to-calcium ratio of 5:1. Our work provides a blueprint for tuning the incorporation of magnesium into calcite under conditions analogous to nature.

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pH

0.25

1. 2. 3. 4. 5. 6.

0.25 0.25 0.25 0.25 0.25 0.25

pH

1. 2. 3. 4. 5. 6. 7.

0.65 0.50 0.375 0.25 0.125 0.05

MgCl2 [M]

0.10

0.50

CaCl2 [M]

10.0 0.10 9.5 0.10 9.0 0.10 8.5 0.10 8.0 0.10 7.6 0.10 Concentration effect

(NH4)2CO3 [M] 1. 2. 3. 4. 5. 6.

CaCl2 [M]

pH

CaCl2 [M]

8.9 0.26 8.9 0.20 8.9 0.15 8.9 0.10 8.9 0.05 8.9 0.02 Cation-to-anion ratio

MgCl2 [M] 0.50 0.50 0.50 0.50 0.50 0.50 MgCl2 [M] 1.30 1.00 0.75 0.50 0.25 0.10

(NH4)2CO3 [M]

pH

CaCl2 [M]

MgCl2 [M]

0.5 0.4 0.3 0.2 0.1 0.05 0.02

8.9 8.9 8.9 8.9 8.9 8.9 8.9

0.10 0.10 0.10 0.10 0.10 0.10 0.10

0.50 0.50 0.50 0.50 0.50 0.50 0.50

were left to stand in tightly sealed containers for 8 and 74 days without any further agitation. The precipitates were then washed several times with milli-Q water, filtered by Büchner vacuum filtration using Whatman glass fiber filter membranes and left to dry in a desiccator. Ammonium carbonate, which has a pH of 8.9, was used for the experiments on cation-to-anion ratio and the concentration effect in Table 1. The pH of the MgCl2/CaCl2 (pH ≈ 6.5) was not adjusted in all the experiments in Table 1, because at high concentrations, magnesium and calcium ions tend to precipitate as Mg (OH)2 and Ca(OH)2 above pH 9.0. To follow the crystallization mechanism of magnesium calcites, a stock solution of sodium carbonate (pH 9.0) was added to a stock solution of MgCl2:CaCl2 (pH 6.5) to achieve the concentrations in Table 1. 100 mL aliquots of the mixture were transferred to 5 PET bottles from which precipitates were filtered after 1 h, and 1, 4, 8, 32, and 74 days. To study the effect of pH, 0.5 M sodium carbonate stock solution was prepared (pH 11.7), from which 100 mL aliquots were placed into 250 mL PET bottles to make 7 independent solutions. The pH of each solution was adjusted by bubbling with pure CO2 while carefully monitoring the pH with a freshly calibrated Mettler Toledo pH meter to achieve the desired pH values of 7.5, 8.0, 8.5, 9.0, 9.5, and 10, respectively. The pH-adjusted carbonate solutions were immediately poured into the MgCl2/CaCl2. Analysis Procedure. The mineral composition of the synthesized carbonates was analyzed by means of synchrotron-based powder diffraction at 22 keV in the MS (Material Science) beamline at the Swiss Light Source. The experiments were done in a Debye−Scherrer geometry using a solid-state silicon microstrip detector, called MYTHEN (microstrip system for time-resolved experiments). For the measurements, the powder samples were crushed and loaded into 1 mm capillaries with 0.01 mm wall thickness. Data acquisition involved the measurement of a blank capillary and air measurement before measuring each sample for 16 min. Silicon standard was measured at intervals in between the measurements for calibration. Data correction was done by subtracting the signal from the air and

EXPERIMENTAL SECTION

Materials. High purity >99% ACS grade calcium chloride dihydrate, magnesium chloride hexahydrate, sodium carbonate, and ammonium carbonate purchased from Sigma-Aldrich were used in all experiments without further purification. High purity >99% ACS grade calcite purchased from Sigma-Aldrich and a natural dolomite sample from Lagoa vermelha Brazil courtesy of Crisogono Vasconcelos (ETHZ) were used as references. Nitrogen bubbled, 18.2 MΩ milli-Q water from laboratory-based purification system (Millipore, Miliford, MA, USA) was used for all the experiments to ensure no dissolved CO2 is present. Synthesis Procedure. Table 1 summarizes the experimental conditions for all the experiments. All reactions were performed in new low-density polyethylene bottles at standard pressure of 1 bar and temperature of 297 K. The synthesis for all the experiments was done by adding a carbonate solution into a CaCl2/MgCl2 solution as per Table 1 followed by vigorous shaking. The freshly prepared solutions B



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Figure 1. (a) Ca K-edge XANES spectra of calcite (black), natural dolomite (red). (b) Mg K-edge XANES spectra of magnesite (black), natural dolomite (red) accompanied with precipitates filtered from aqueous solutions of composition Mg2+-to-Ca2+-to-CO32− ratio of 5:1:2.5 (Ca2+ = 0.2 M) filtered after 1 h “fresh precipitates” (maroon), after 8 days in solution (dark green) and after 74 days in solution (light green). (c) Synchrotron XRD patterns of calcite and dolomite accompanied with precipitates from solution containing Mg2+-to-Ca2+-to-CO32− ratio of 5:1:2.5 (Ca2+ = 0.2 M) after 1, 4, 8, 32, and 74 days. the empty capillary. The 2θ values were then converted from 22 keV (0.563565 Å) to Cu α (1.54056 Å) using eq 3 and plotted without any further manipulation. ÅÄÅ Å 360 1 × arcsinÅÅÅÅ1.54056 × 2θCuα = ÅÅÇ π 0.563565 É i π zyÑÑÑÑ zzÑÑ × sinjjj2θmeasured × 360 {ÑÑÖÑ (3) k

nm, and lamp current of 4.0 mA. For SEM analysis, 20 nm gold coating was done on the samples, analysis was done using 6 keV acceleration voltage, and detection was done by means of a secondary electron detector.

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RESULTS Formation of Magnesium Calcite from Additive Free Aqueous Solutions. Magnesium calcites were synthesized directly from aqueous solution of Mg2+-to-Ca2+-to-CO32− ratio of 5:1:2.5 ([Ca2+] = 0.2 M), pH 8.9, free of any additives under ambient temperature and pressure. At the onset of the reaction, a gel formed immediately which persisted for several hours and gradually cleared, accompanied with formation of precipitates. The time dependent reaction evolution was monitored for 74 days; the precipitates were filtered after 1 h, and 1, 4, 8, 32, and 74 days. Different chemical phases of calcium carbonate can be identified by both calcium L25 and K edge XANES.26 The Ca K-edge XANES spectrum of the fresh precipitates (filtered after 1 h) shown in Figure 1a displays a single pre-edge peak centered at ∼4039 eV and a single intense white line peak at ∼4049 eV, which is consistent with that of an amorphous calcium carbonate phase.26 Similarly, the Mg Kedge XANES spectrum of the fresh precipitates (filtered after 1 h) in Figure 1b displays only a broad white line feature centered at 1310 eV, which is characteristic of an amorphous magnesium phase,9 ascertaining that amorphous magnesium calcium carbonate is the initial solid phase that forms during formation of high magnesium calcites directly from aqueous solutions. On a longer time scale, the Ca K-edge XANES spectra of the precipitates aged in solution after 8 days, as shown in Figure 1a, display a doubly split weak pre-edge peak at ∼4039 eV, a shoulder peak at ∼4045 eV, and a white line at ∼4047.5 eV that is split into two peaks with a further post edge peak centered at ∼4060 eV. All these features are comparable to those of calcite and dolomite reference spectra, confirming that the precipitates become more crystalline after being aged for a number of days. Similarly, the Mg K-edge XANES spectra in Figure 1b of the precipitates after 8 days display crystallinity

To determine the (104) peak position, single peak fit based on Gaussian function was done using Origin software. Equation 4 shows the Gaussian formula applied 2

y = y0 +

A e−4ln(2)(x − xc) w

π 4ln(2)

/ w2

(4)

where A is the area, yo is the y-axis value of the base of the curve, w is the full width at half maximum, and Xc is the X-axis value of the peak center. The d104 was determined using eq 5 based on the Bragg equation d=

λ 2sinθ

(5)

where λ is the wavelength of the rays, d is the spacing between the atoms, and θ is the angle between the incident rays and the surface of the crystals. For X-ray absorption near edge structure (XANES) measurements, powder samples of the precipitates were deposited on a carbon tape which was attached to a copper plate and mounted in a vacuum chamber for analysis. Both total electron yield and fluorescence measurements were done at the Ca and Mg K edge energies in the PHOENIX beamline of the Swiss Light Source. The composition of the precipitates was further analyzed by atomic absorption spectroscopy (AAS). Magnesium and calcium AAS standard solutions were purchased from Sigma-Aldrich. Through serial dilutions of the standard solutions, 5 calibration standard solutions for each metal were prepared. For the analysis, the precipitates were first dissolved using 0.2 M sulfamic acid and then diluted with milli-Q water. Calcium ion concentration was done using N2O/acetylene flame with a 422.7 nm calcium lamp, slit width of 0.5 nm, and lamp current of 10.0 mA. For magnesium, an air/acetylene flame was used with a 285.2 nm magnesium lamp, slit width of 0.5 C



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with features at 1306, 1310, and 1320 eV, characteristic of magnesium calcites.27 Congruent with the XANES spectra, in the synchrotron XRD patterns shown in Figure 1c, precipitates that were aged for at least 4 days, show resolved 2θ reflections at 23.9° (012), at 30.7° (104), at 33° (006), and at 37.1° (110), respectively, which are characteristic of calcites, whereas the sample aged only for 1 day shows a broad pattern with no peaks indicating lack of long-range order after 1 day. This suggests that amorphous magnesium calcium carbonate forms first and gradually transforms into magnesium calcite in aqueous solution. The amount of crystalline magnesium in the precipitates was determined according to eq 6:28−31 XMg = −3.6396d(104) + 11.0405

(6)

where d(104) is the d spacing in Å calculated from the (104) peak of the precipitates and XMg is the amount of magnesium incorporated in calcite. The fraction of crystalline magnesium incorporated into calcite was found to be 40 ± 2 mol % for all crystalline precipitates between 4 and 74 days confirming the formation of very high magnesium calcites. There were no significant changes in the XRD patterns of the precipitates after 4 days, indicating that once crystallization occurred, the precipitated very high magnesium calcites remained stable and did not transform into aragonite or any other crystalline phases. Moreover, in contrast to the sharp and narrow peaks observed in calcite and dolomite reference patterns, the XRD patterns of the precipitates in Figure 1c have broadened peaks indicative of nanosized crystals. Consistent with this, the SEM images of the crystalline precipitates in Figure 2a show that the precipitates are spherical, consisting of aggregates of smaller building blocks. HRTEM (high resolution transmission electron microscopy) and SAED (selected area electron diffraction) analysis of the precipitates in Figure 2b shows that these spherical aggregates are made of nanocrystal building blocks. Further, in spite of the precipitates displaying infinitesimal changes in the XRD patterns after 4 days, morphological variations are observed in the SEM and TEM images. In particular, the precipitates filtered after 8 and 74 days have pores within the spherical clusters, and the HRTEM images in Figure 2b show an increase in the crystal sizes from about 5 nm (4 days precipitates) to about 10 nm (74 days precipitates) indicative of crystal growth. Effect of Cation-to-Anion Ratio. The main anions involved in nucleation and growth in the Ca−Mg−CO3 system are the carbonate ions. To establish a correlation between cation-to-anion ratio and nucleation in the Ca−Mg− CO3 system, we performed a series of precipitation experiments with a varying concentration of carbonate ions, at a constant carbonate buffer pH of 8.9 and Mg2+-to-Ca2+ ratio of 5:1 ([Ca2+] = 0.1 M). When the concentration of carbonate ions increased from 0.02 M to 0.2 M, a shift in the calcite reflections to higher 2θ values is observed in the XRD patterns in Figure 3a, corresponding to increased magnesium incorporation in the calcite crystal lattice from 11% to 33% (Figure 3b). Additionally, in Figure 3a, we observe that the carbonate ions play a critical role in determining which crystalline polymorph is preferentially precipitated from aqueous solutions. When the [CO32− ] = 0.02 M, such that Mg2+-to-Ca2+-to-CO32− ratio is 5:1:0.2, the XRD pattern is characterized by intense aragonite reflections at 26°, 27°, 33°, 36°, and 38° and a very small amount of magnesium calcite

Figure 2. (a) SEM images of precipitates synthesized from Mg2+-toCa2+-to-CO32− ratio of 5:1:2.5 (Ca2+ = 0.2 M) and filtered from solution after 4, 8, 32, and 74 days, respectively. (b) High resolution TEM images with SAED pattern inserts of precipitates synthesized from Mg2+-to-Ca2+-to-CO32− ratio of 5:1:2.5 (Ca2+ = 0.2 M) and filtered from solution after 4 and 74 days. (The arrows and shaded regions highlight a single magnesium calcite crystal.)

(11%) characterized by weak magnesium calcite reflections. Increasing the carbonate ion concentration results in a higher intensity of magnesium calcite reflections at 23−24° (012), 29−31° (104), 31−33.5° (006), and 36−37.5° (110), indicating that precipitation of magnesium-rich calcites is increasingly favored relative to aragonite. Interestingly, coprecipitation of monohydrocalcite is additionally observed in the 8 days precipitates (dotted line) from solutions of Mg2+to-Ca2+-to-CO32− ratio of 5:1:2 and 5:1:3 but not in the 74 days precipitates of the 5:1:2 solution, indicating that monohydrocalcite is a metastable precipitate in these solutions. Most interestingly, when the carbonate concentration is such that the Mg2+-to-Ca2+-to-CO32− ratio is greater than or equal D



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32°, 34°, and 35.7° indicating coprecipitation of nesquehonite (MgCO3·3H2O). This is also consistent with the high content of magnesium revealed by atomic absorption spectroscopy (AAS) measurements compared to the amount calculated from the XRD measurements in Figure 3b. Influence of pH on the Formation of Magnesium Calcites. Since pH influences carbonate speciation in aqueous solutions, it is expected that changes in pH influence the Ca− Mg−CO3 chemistry. To elucidate the role of pH in the nucleation and growth in the Ca−Mg−CO3 system, a series of experiments at environmentally relevant pH values between 7.5 and 10 were carried out at a constant Mg2+-to-Ca2+-to-CO32− ratio of 5:1:2.5 ([Ca2+] = 0.1 M). Synchrotron XRD data in Figure 4a show that below pH 9.0, the diffraction patterns are very similar, with both patterns having aragonite reflections of similar intensities at 26°, 27°, 33°, 36°, and 38°, and the magnesium calcite peaks do not show significant shifts in the 2θ positions. However, above pH 9.0, there is a shift in the magnesium calcite reflections to higher 2θ values, concomitant with a marked increase in the amount of incorporated magnesium from 20 mol % at pH 7.5 to 42 mol % at pH 10 (Figure 4b). Even though the amount of magnesium increases significantly at pH 9.5 and above, the formation of monohydrocalcite becomes more dominant under these high pH conditions as revealed by the increase in intensity of the monohydrocalcite reflections at 20.5°, 29°, 32°, and 38° in Figure 4a and the significantly lower percentage of magnesium detected in AAS compared to the amount calculated from the XRD patterns in Figure 4b. Effect of Concentration. By varying the overall concentration of the Ca2+, Mg2+, and CO32− ions (at a constant ratio of 5:1:2.5 and a carbonate buffer pH of ∼8.9) in the mother liquor, we show the strong influence of concentration on crystallization in the Ca−Mg−CO3 system. XRD patterns in Figure 5a show that at low concentrations, in addition to magnesium calcite reflections, aragonite reflections are also evident at 26°, 27°, 33°, 36°, and 38°. However, as the initial concentration of the solution is increased, the intensity of the aragonite reflections decrease and eventually disappear. A shift in the magnesium calcite reflections to higher 2θ angles also takes place with increasing concentration, indicative of greater magnesium incorporation in the calcite precipitates. Analysis of the fraction of magnesium carbonate in the precipitated magnesium calcites as a function of concentration in Figure 5b shows that as the concentration increases, the amount of incorporated magnesium increases from 25 mol % at the lowest concentration to 45 mol % (disordered calcium rich dolomite) at the highest concentration used in the experiments. Consistent with observations in Figure 3a, monohydrocalcite reflections at 20.5°, 29°, 32°, and 38° are observed in the 8-day precipitates from the solutions consisting of [Ca2+] = 0.05, 0.10, and 0.15 M (dotted line) but not in the precipitates filtered after 74 days (solid line), indicating that monohydrocalcite is a metastable phase in these solutions.

Figure 3. Effect of cation: anion ratio (a) synchrotron XRD patterns of precipitated solids after 8 days (dotted line) and 74 days (solid line) from Mg2+-to-Ca2+ ratio of 5:1 ([Ca2+] = 0.1M), pH 8.9, and varying carbonate concentration of 0.02 M (5:1:0.2), 0.05 M (5:1:0.5), 0.1 M (5:1:1), 0.2 M (5:1:2), 0.3 M (5:1:3), 0.4 M (5:1:4), and 0.5 M (5:1:5). (b) Correlation between carbonate concentration and amount of magnesium carbonate in the precipitated CaxMg1−xCO3 crystals.

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DISCUSSION We have systematically studied the crystallization products in the Ca−Mg−CO3 system under quiescent conditions at a constant initial magnesium-to-calcium ratio of 5:1 for time scales of up to 74 days. While it is not possible to formulate a quantitative phase diagram of this complex system, Figure 6 illustrates the general trends observed experimentally. Our findings summarized in Figure 6 show that in aqueous solution

to 5:1:3 ([CO32−] = 0.3M), additional reflections characteristic of nesquehonite are observed at 23°, 24.8°, 27.4°, 29.4°, 30°, E



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Figure 4. Influence of pH: (a) Synchrotron XRD patterns of precipitated solids after 8 days (dotted line) and 74 days (solid line) from Mg2+-to-Ca2+-to-CO32− ratio of 5:1:2.5 (Ca2+=0.1 M and a pH of 7.5, 8.0, 8.5, 9.0, 9.5, and pH 10. (b) Correlation between pH and the amount of magnesium in the precipitated CaxMg1−xCO3 crystals.

Figure 5. Concentration effect: (a) synchrotron XRD patterns of precipitated solids after 8 days (dotted line) and 74 days (solid line) from solutions containing Mg2+-to-Ca2+-to-CO32− ratio of 5:1:2.5 at Ca2+ concentrations of 0.02, 0.05, 0.1, 0.15, 0.2, and 0.26 M. (b) Correlation between concentration of all ions and the amount of magnesium in the precipitated CaxMg1−xCO3 crystals (Note: all ions scale up with the Ca2+ to maintain a constant ratio of 5:1:2.5).

with constant initial magnesium-to-calcium ratio of 5:1, polymorph control can be achieved through changing the pH, concentration, and cation-to-anion ratio. Consistent with previous experimental19,32−34 and theoretical studies,12 ara-

gonite precipitation is favored (Figures 3a and 4a) in dilute solutions and at low carbonate concentrations corresponding F



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summarized in Figure 6 show that increasing the saturation levels by increasing the pH or carbonate/overall solution concentration favors the formation of thermodynamically unstable very high magnesium calcites.41,42 A time-dependent reaction evolution is also observed during the formation of this very high magnesium calcite, pointing to a multistep formation pathway (Figure 7). A clearly visible dense liquid (gel) formed

Figure 6. Summary of results showing a qualitative overview of the reaction conditions and the crystalline polymorphs formed after 74 days. (*overview picture not to scale.)

to carbonate-to-calcium ratios (2:1, we found that formation of very high magnesium calcites with up to 45 mol % magnesium occurs (Figure 5b). Similar compositions have been reported from laboratory synthesis in aqueous solutions with either high magnesium-to-calcium ratio, or in the presence of organic additives or at elevated temperatures and in a variety of sedimentary deposits.18−23,29,30,35−37 At elevated concentrations (with concomitant high carbonate ion concentrations), very high saturation levels are attainable. Thermodynamic calculations shown in Figures S2, S3, and S4 illustrate that in the Ca−Mg−CO3 system, increasing the pH or the overall solution/carbonate concentration increases the supersaturation with respect to almost all calcium, magnesium, and mixed calcium magnesium carbonate crystalline polymorphs. While classical thermodynamics predicts the degree of supersaturation and the relative stability of different polymorphs, kinetics plays a critical role in determining the crystallization pathway, and may well affect the final polymorph and its composition.38,39 Wang and coworkers argue that, in the Ca−Mg−CO3 system at sufficiently high saturation levels, formation of very high Mg-ACC becomes kinetically favorable; the subsequent crystallization of the Mg-ACC results in the formation of very high magnesium calcites.40 Consistent with this, our findings

Figure 7. Proposed mechanism for formation of very high magnesium calcites directly from aqueous solutions at ambient conditions.

at the onset of the reaction which slowly cleared out as the formed precipitates began to settle, suggesting a possible coalescence and solidification of nanoscale droplets resulting in formation of an initially amorphous solid phase consistent with theoretical predictions of spinodal decomposition.43,44 This has been observed in the calcium carbonate system at sufficiently large saturations for calcium concentrations above 5 mM.45,46 Amorphous magnesium calcium carbonate (Mg-ACC) is the first solid phase observed to form, consistent with observations from biomineralization processes and laboratory synthesis of calcium magnesium carbonates.11,15,16,30,47,48 The Mg-ACC exhibited enhanced stability compared to impurity free ACC,49 which could be attributed to the presence of structural water bound to the magnesium ions in the Mg-ACCs, which slows down crystallization as suggested by Lin and co-workers.48 Previous work shows that the magnesium content in the MgACC strongly depends on the initial solution composition.15,30,49−52 In particular, the presence of organic additives, magnesium:calcium ratio, and carbonate concentration have been shown to strongly influence the magnesium amount in Mg-ACC.15,16,30,47,49,53−55 At a constant magnesium-to-calcium ratio, pH and carbonate concentration have also been shown to influence the amount of magnesium incorporated in the Mg-ACC.30,56 In extension to this work, our findings show a strong correlation between the initial solution concentration, G



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pH and cation-to-anion ratio, and the amount of magnesium incorporated in the final magnesium calcite crystals. This strong influence of solution composition on the final crystalline polymorph can be attributed to the critical role of the amorphous intermediate phase. Blue and co-workers showed a 1:1 dependence between the magnesium content in the MgACC and the magnesium content in the magnesium calcite crystalline phase,51 implying that by tuning the initial solution conditions, we tune the amount of magnesium incorporated in the Mg-ACC and thereby effectively control the amount of magnesium incorporated in the magnesium calcites. In the crystallization phase in Figure 7, crystallization is envisioned to occur through a dissolution and recrystallization mechanism of the Mg-ACC as recently observed from high resolution TEM experiments.15 When left in the mother liquor for up to 74 days, the crystalline magnesium calcites neither transformed to any other crystalline polymorph nor significantly changed in the fraction of magnesium in the crystal, further supporting the critical role of the initial solution composition, in defining the magnesium content in the crystalline phase. Finally, based on morphological changes in the precipitates in Figure 2a and the crystal growth observed in Figure 2b, one can envision that the last step in Figure 7 is dominated by very slow growth through Ostwald ripening to form well-defined very high magnesium calcite crystals. The high kinetic stability of the magnesium calcites, against further conversion, is likely due to a high activation energy barrier required to transform to the most stable phase (dolomite).28 The solution composition and degree of saturation not only influences the amount of magnesium incorporation into calcite. Under some solution conditions, monohydrocalcite is observed to occur as a metastable intermediate phase (Figures 3a and 5a) which disappears after prolonged aging in the mother liquor. This is consistent with previous literature which has reported the formation of monohydrocalcite as an intermediate phase in aragonite formation.55,57,58 However, at very high pH conditions (Figure 4a), preferential precipitation of monohydrocalcite occurs and it remains stable in solution even after prolonged aging in the mother liquor (74 days). This preferential precipitation of monohydrocalcite at higher pH in the presence of magnesium ions has also been observed in previous laboratory studies and in nature.59,60 The formation of monohydrocalcite may be attributed to the pH driven increase in carbonate-to-bicarbonate ratio, which increases the carbonate-to-calcium ratio (Figure S5). It is however not yet clear why the monohydrocalcite displays enhanced stability under high pH conditions. Nesquehonite coprecipitation is also observed at very high carbonate concentrations (carbonate:calcium ratios ⩾3:1) possibly due to the excess carbonate ions with respect to calcium ions. Magnesium calcites are known to display varying crystal morphology depending on the amount of magnesium carbonate in the crystals. Figure 8 shows that high magnesium calcites formed directly from aqueous solution look mesostructured and the sizes of the individual composites decrease with increasing magnesium content in the crystals. With 10 mol % magnesium incorporated in the calcite crystal lattice, the precipitates display a plate-like crystalline morphology stacked together into spheres. With 20 mol % magnesium incorporated, the individual crystal units are much smaller (

Tuning the Incorporation of Magnesium into Calcite during Its

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