Research Article pubs.acs.org/journal/ascecg
Tuning Crystal Polymorphisms and Structural Investigation of Precipitated Calcium Carbonates for CO2 Mineralization Ribooga Chang,† Dasol Choi,† Min Hee Kim,† and Youngjune Park*,†,‡ †
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School of Earth Sciences and Environmental Engineering, Gwangju Institute of Science and Technology (GIST), 123 Cheomdangwagi-ro, Buk-gu, Gwangju 61005, Republic of Korea ‡ School of Integrated Technology (Graduate Program of Energy Technology), Gwangju Institute of Science and Technology (GIST), 123 Cheomdangwagi-ro, Buk-gu, Gwangju 61005, Republic of Korea S Supporting Information *
ABSTRACT: Mineral carbonation, which involves spontaneous reactions of CO2 with alkaline earth metals such as calcium or magnesium, is considered one of the most attractive options to sequester CO2 because CO2 can be permanently stored in an inert solid forming stable inorganic carbonate. Moreover, the precipitated CaCO3 has various potential applications in industrial areas, including adhesives, sealants, food, pharmaceuticals, paints, coating, paper, cement, construction materials, etc. In particular, it is expected that the total cost of the carbon capture and storage process could be partly offset by producing value-added CaCO3 materials. In order to add value to the precipitated CaCO3 produced in the ex-situ mineral carbonation process, CaCO3’s polymorphs, as well as other properties such as particle size, shape, density, color, and brightness, must be finely tuned. Among CaCO3’s polymorphs, calcite, aragonite, and vaterite, calcite is considered to be the most thermodynamically favorable structure at ambient temperatures. However, kinematic constraints in the crystallization induced by synthetic factors are known to significantly affect the formation of polymorphs as well. Here, we revisited the effects of the synthetic factors such as pH, temperature, feeding order, and concentration and molar ratio of Ca2+ and CO32− on the formation of CaCO3’s polymorphs to provide fundamental insight into how to control the polymorphism of CaCO3 with the ultimate goal of creating value-added mineral carbonation products. ATR FT-IR spectroscopy and a powder X-ray diffraction analysis were performed on the precipitated CaCO3 using model chemicals, K2CO3 and CaCl2, and CaCO3’s thermal stability was also investigated. KEYWORDS: Carbon Storage, Mineral carbonation, Calcite, Aragonite, Vaterite
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INTRODUCTION Although there is still debate, it is generally accepted that global climate change is caused by excessive global warming gases including CO2. Currently, the CO2 levels in the atmosphere have surpassed 400 ppm,1 and there are warnings that without proper action the global mean surface temperature could rise by 4.8 °C above preindustrial levels by the end of the 21st century.2 This unprecedented rise in temperature is considered to be mainly attributed to the increase in use of fossil fuels such as coal, oil, and natural gas.3,4 However, it is projected that fossil fuels will continue to be one of the most predominant energy sources and the demand for fossil fuels will be further increased by developing countries.5 In order to continue to use fossil fuels in terms of sustainability, there are two preferable strategies: (i) more energy-efficient technologies can reduce energy demand, resulting in less carbon emissions, and (ii) emitted CO2 can be captured, transported, and stored via carbon capture and storage (CCS) technologies so that less CO2 is released into the atmosphere. More importantly, CCS is recognized as the most immediate solution to mitigate CO2 emissions and global climate change.6 © 2016 American Chemical Society
Once CO2 is captured from large point sources such as a coal-fired power plants or other industrial facilities, it must be stored or sequestered permanently, or otherwise for a long time.7 In general, geological CO2 sequestration, which involves CO2 injection into geological underground formations such as deep saline aquifers, unminable coal beds, and oil and gas reservoirs is recognized as one of the most promising options due to the large potential storage volume. However, there are several concerns regarding geological CO2 storage. For example, finding a suitable storage site could be challenging in some regions or countries. In addition, lack of knowledge on geological characterization may increase the possibility of CO2 leakage during and/or after injection, and the costs for characterization and monitoring hence could be increased in some instances. The cost of remediation of leaky storage also should not be ignored. Ocean storage is another technique for CO2 storage. Because the oceans are the largest active carbon Received: October 6, 2016 Revised: November 11, 2016 Published: December 16, 2016 1659
DOI: 10.1021/acssuschemeng.6b02411 ACS Sustainable Chem. Eng. 2017, 5, 1659−1667
Research Article
ACS Sustainable Chemistry & Engineering
important parameters to control the polymorphs of CaCO3.21−23 Although the precipitation of CaCO3 involves a simple reaction between Ca2+ and CO32−, the experimental procedure to form a specific polymorph of CaCO3 varies significantly by the synthetic conditions, and therefore the controllable formation of CaCO3 is still a practical challenge. Here, we revisited the effects of the synthetic factors including concentration and molar ratio of Ca2+ and CO32−, feeding order, temperature, and pH on the formation of CaCO3’s polymorphs to elucidate their exact features and to provide fundamental knowledge and insight into how to control the polymorphism of CaCO3, with the goal of eventually creating value-added mineral carbonation products.
sink on Earth, CO2 injection into the deep ocean offers huge potential to store large amounts of anthropogenic CO2.8 CO2 could be permanently stored in the deep ocean floor by injection at depths greater than 3000 km, where the density of liquefied CO2 becomes higher than that of the surrounding water.9,10 Because of uncertainty in the ecocide and ocean acidification issues, however, ocean storage is not currently considered a feasible option for CO2 storage,7 and therefore more studies are needed. On the other hand, mineral carbonation technology, which involves exothermic reactions of CO2 with alkaline earth metals such as calcium or magnesium, is recently receiving renewed attention although it has been investigated for a few decades. Once converted into mineral carbonate, CO2 can be permanently stored in an inert solid forming thermodynamically more stable CaCO3 or MgCO3. The calcium or magnesium sources can be obtained from naturally occurring silicate minerals (i.e., wollastonite, olivine, serpentine), or industrial wastes (i.e., fly ash, steel slag, cement wastes).11−14 There are several key advantages of mineral carbonation for CO2 storage. For example, first, a leak-free fixation without mid- or long-term monitoring is possible. Second, all exothermic reactions in the carbonation process can eliminate the significant energy requirement for CO2 storage. Finally, industrial wastes and environmentally hazardous materials such as steel slags or asbestos can be recycled and converted into value-added products, and thus there is potential to compensate for the costs occurring in the whole CCS process. In fact, CaCO3 has many industrial applications, including, for example, adhesives, sealants, food and pharmaceuticals, paints, coatings, paper, cements, construction materials, etc.,15 and these applications are dependent on CaCO3’s physicochemical characteristics such as particle size distribution, structure, density, color, brightness, and other properties. In order to add value to the precipitated CaCO3, particularly in connection to ex-situ mineral carbonation schemes for CO2 storage and utilization, the formation of the polymorphs of CaCO3 with specific shape and size must be controlled.16 It has been reported that synthetic factors including pH, temperature, concentration and ratio of carbonate and calcium ions, additives, stirring, and reaction time may significantly affect the formation of the polymorphs. Ševčiḱ et al. investigated the effects of stirring velocities and temperatures ranging from 200 to 600 rpm and 30 to 90 °C, respectively, on the formation of vaterite precipitating simple chemicals of CaCl2·2H2O and K2CO3, and they reported that pure vaterite having ≥99 wt % purity was obtained at estimated optimal reaction conditions of 600 rpm and 60 °C.17 The effect of temperature on the formation of each polymorph was investigated, and it was found that vaterite was formed dominantly at a relatively low temperature condition, although calcite is considered thermodynamically the most stable at ambient temperature, and aragonite could be obtained at higher temperature.18,19 The effects of reaction time and transformation kinetics among the amorphous calcium carbonate (ACC), vaterite, and calcite were investigated via in-situ time-resolved energy dispersive X-ray diffraction by Rodriguez-Blanco et al.,20 and the results revealed that ACC transforms to calcite via a vaterite intermediate through two-step crystallization. Because precipitation occurs from the reaction between carbonate and calcium ions, and further generation of CO32− is achievable at a high pH condition, the solution pH is also considered one of the
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MATERIALS AND METHODS
CaCO3 precipitation was carried out by mixing CaCl2 and K2CO3 solutions with varying molar concentrations ranging from 0.2 to 0.4 M. CaCl2 (p.a.) and K2CO3 (p.a.) were purchased from Junsei Chemical Co., Ltd. (Japan) and Daejung Chemical & Metals Co., Ltd. (Republic of Korea), respectively. In order to explore the effect of Ca2+ and CO32− concentration on the formation of CaCO3 polymorphism, the precipitation experiments were conducted in two different conditions, CO32− rich (1−3 in Table 1) and Ca2+ rich (4−6 in Table 1) cases.
Table 1. List of CaCO3 Samples Prepared under Different Conditions conc (M) sample notation
CaCl2
K2CO3
reaction conditions
1
0.4
0.2
2
0.4
0.3
3
0.4
0.4
4
0.2
0.4
5
0.3
0.4
6
0.4
0.4
7
0.4
0.2
8
0.4
0.4
9
0.2
0.4
10
0.4
0.4
100 mL of CaCl2 solution (10 mL × 10 times) was added to 300 mL of K2CO3 at 25, 50, and 80 °C 100 mL of CaCl2 solution (10 mL × 10 times) was added to 300 mL of K2CO3 at 25, 50, and 80 °C 100 mL of CaCl2 solution (10 mL × 10 times) was added to 300 mL of K2CO3 at 25, 50, and 80 °C 100 mL of K2CO3 solution (10 mL × 10 times) was added to 300 mL of CaCl2 at 25, 50, and 80 °C 100 mL of K2CO3 solution (10 mL × 10 times) was added to 300 mL of CaCl2 at 25, 50, and 80 °C 100 mL of K2CO3 solution (10 mL × 10 times) was added to 300 mL of CaCl2 at 25, 50, and 80 °C 100 mL of CaCl2 solution was added to 100 mL of K2CO3 at 25 °C 100 mL of CaCl2 solution was added to 100 mL of K2CO3 at 25 °C 100 mL of K2CO3 solution was added to 100 mL of CaCl2 at 25 °C 100 mL of K2CO3 solution was added to 100 mL of CaCl2 at 25 °C
For the CO32− rich case, a total of 100 mL of 0.4 M CaCl2 solution was divided into ten 10 mL solution batches, and each batch was subsequently added dropwise to 300 mL K2CO3 solutions having different molar concentrations from 0.2 to 0.4 M for 2 min. In order to maintain a constant concentration of carbonate ions during the precipitations, K2CO3 was supplemented at the end of injection of each batch. After 1 min of stirring at 350 rpm, another batch was added until a total of 100 mL of 0.4 M CaCl2 solution was added. During the precipitation, temperature was constantly maintained at 25, 50, and 80 °C to examine its effect. Before and after precipitation, pH was recorded by a pH meter (S220, Mettler-Toledo International Inc. (Switzerland)). For the Ca2+ rich cases, a total of 100 mL of 0.4 M K2CO3 solution was added to 300 mL of CaCl2 solution, where the molar concentration of the latter ranged from 0.2 to 0.4 M. The precipitation procedure was the same. 1660
DOI: 10.1021/acssuschemeng.6b02411 ACS Sustainable Chem. Eng. 2017, 5, 1659−1667
Research Article
ACS Sustainable Chemistry & Engineering
Figure 1. Crystal structures of (a) calcite, (b) aragonite, and (c) vaterite. Ca atoms are displayed as large yellow balls, and carbonate groups are illustrated with gray (carbon) and red (oxygen) balls. Vaterite is depicted with a hexagonal P63/mmc structure that accounts for a partial occupancy of 1/3 of the carbonate groups.27 For a comparison of the effect of Ca2+ and CO32− concentration on the formation of CaCO3 polymorphism, CaCl2 or K2CO3 was precipitated without keeping constant concentrations of Ca2+ or CO32−. In the cases of samples 7 and 8 shown in Table 1, 100 mL of 0.4 M CaCl2 solution was added to 0.2 and 0.4 M K2CO3 solutions (100 mL) at once. Stirring at 350 rpm was conducted for 30 min. Similarly, in the case of samples 9 and 10 shown in Table 1, 100 mL of 0.4 M K2CO3 solution was added to 0.2 and 0.4 M CaCl2 solutions (100 mL) at once. The precipitates were filtered using filter paper (0.45 μm, GE Healthcare (Chicago, IL)) and a Büchner funnel, and then washed with deionized water. They were subsequently dried at 60 °C for 24 h in a convection oven (OF-02GW, Jeio Tech. Co., Ltd. (Republic of Korea)). For characterization of the precipitated CaCO3, a Fourier-transform infrared (FT-IR) spectrometer (Nicolet iS5, Thermo Fisher Scientific Inc. (Madison, WI)) equipped with a deuterated triglycine sulfate (DTGS) detector and an attenuated total reflectance (ATR) accessory (single reflection ZnSe crystal) was used. All spectra were recorded with resolution of 4 cm−1 in the 4000−400 cm−1 wavenumber region with 32 scans at room temperature. A variable temperature powder Xray diffractometer (PXRD, D/MAX-2500, Rigaku Co. (Japan)) was used for a quantitative analysis of CaCO3 samples. The PXRD patterns were obtained in the 2θ range 20−60° (0.01° step size and 2°/min scan speed) using Cu Kα radiation (λ = 1.5406 Å) at a generator voltage of 40 kV and a generator current of 300 mA. The obtained diffraction patterns were analyzed through the Rietveld refinement method via the MAUD (material analysis using diffraction) package.24 Morphological structures of the precipitated CaCO3 were examined by using scanning electron microscopy (SEM, S-4700 EMAX System, Hitachi, Ltd. (Japan)) images.
could have different industrial applications with specific shape and size. For example, spherical microporous vaterite particles are used in regenerative medicine and drug delivery due to their biodegradability.30 Vaterite can be also applied as a coating pigment to ink jet paper due to its higher hydrophilicity relative to calcite.31 A needle-like aragonite can be applied to lightweight plastics, and the medical, pharmaceutical, and cosmetics industries by altering optical properties or by achieving a high aspect ratio.32−37 In order to explore the effect of concentration and molar ratio of Ca2+ and CO32−, feeding order, temperature, and pH on the formation of CaCO3 polymorphs, precipitation reactions were performed using CaCl2 and K2CO3 in an aqueous solution at constant temperature conditions of 25, 50, and 80 °C. It is known that pH and concentration of Ca2+ and CO32− ions influence the formation of CaCO3’s polymorphs.38,39 In order to prevent significant changes in the pH and the concentration of Ca2+ and CO32− ions during the reactions, the precipitations proceeded divided into several steps, filling up CaCl2 or K2CO3 within each step to maintain the initial concentration. The results are summarized in Table 2 (precipitated at 25 °C) and Table 3 (precipitated at 50 and 80 °C). During the precipitations, pH was also monitored, and no significant pH changes were observed. On the other hand, as a comparison group, CaCO3 precipitations via CaCl2 and K2CO3 solutions were also performed without maintaining constant concentrations of Ca2+ or CO32− ions (Table 1).
RESULTS AND DISCUSSION As shown in Figure 1, CaCO3 is known to form three polymorphs, calcite (hexagonal R3̅2/c), vaterite (hexagonal P63/mmc), and aragonite (orthorhombic Pmcn),25,26 and each polymorph has various applications owing to distinct physicochemical characteristics.27 Among the polymorphs of CaCO3, calcite is thermodynamically the most stable at ambient conditions, and thus, it is recognized as an industrially important inorganic mineral for various applications depending on its morphological structures. For example, ultrafine rhombohedral calcite is applied in the plastic and sealant industry, whereas scalenohedral calcite is used in paint and paper applications.28 On the other hand, thermodynamically less stable aragonite and metastable vaterite, which can be stabilized kinetically or biochemically at ambient synthetic conditions,29 exhibit different morphological structures rendering distinct properties compared to calcite, and therefore, they
Table 2. Analysis Results of CaCO3 Samples Precipitated at 25 °C
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25 °C composition (%)
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sample notation
calcite
1 2 3 4 5 6 7 8 9 10
3 4 2 19 17 15 29 36 26 35
aragonite
vaterite
initial pH
final pH
yield (%)
97 96 98 81 83 85 71 64 74 65
10.9 11.0 11.4 7.8 8.1 7.5 10.9 11.4 7.8 7.5
11.6 11.6 11.6 7.5 7.5 7.5 7.4 9.5 11.2 9.6
115.8 101.2 95.8 95.0 98.5 96.3 96.8 97.4 97.3 95.4
DOI: 10.1021/acssuschemeng.6b02411 ACS Sustainable Chem. Eng. 2017, 5, 1659−1667
Research Article
ACS Sustainable Chemistry & Engineering
(712 and 700 cm−1), assigned as the peaks from aragonite, were observed (Figure 2c−f).41 The results indicate that, at room temperature conditions, precipitated CaCO3 existed in the form of vaterite rather than calcite. On the other hand, aragonite favorably formed at the elevated temperature of 50 °C, and it became the most dominant polymorph at 80 °C. For a detailed quantitative analysis, diffraction patterns were obtained using PXRD measured at room temperature, as shown in Figure 3. The diffraction patterns were analyzed by Rietveld refinement.42 The refinement fits for sample 8 precipitated at 25, 50, and 80 °C are shown in Figure 4, and the polymorph compositions (%) of the CaCO3 samples are tabulated in Tables 2 and 3. As shown in Table 2, samples 1−6 precipitated at 25 °C included vaterite as the main component as well as a relatively small amount of calcite. The change of pH and the change of the concentration of K2CO3 or CaCl2 were minimized in the process of precipitating samples 1−6. For samples 1−3, a total of 100 mL of 0.4 M CaCl2 solution was added to 300 mL of K2CO3 solution in 10 steps. K2CO3 in an amount equivalent to the amount of precipitation in each step was added to maintain the concentration and keep the pH constant. For samples 4−6, a total of 100 mL of 0.4 M K2CO3 solution was added to 300 mL of CaCl2 solution in the same manner as described above. An interesting finding was that the calcite content was higher in samples 4−6 than in samples 1−3. In samples 1−3, the pH was about 11, and the CO32− ions were richer than the Ca2+ ions.
Table 3. Analysis Results of CaCO3 Samples Precipitated at 50 and 80 °C 50 °C composition (%) sample notation
calcite
aragonite
1 2 3 4 5 6
6 8 6 52 45 46
94 92 93 46 54 54
vaterite
