CO2 Absorption and Sequestration as Various Polymorphs of

Dharmalingam Sivanesan , Young Eun Kim , Min Hye Youn , Ki Tae Park , Hak-Joo Kim , Andrew Nirmala Grace , Soon Kwan Jeong. RSC Advances 2016 6 ...
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CO2 Absorption and Sequestration as Various Polymorphs of CaCO3 Using Sterically Hindered Amine Mari Vinoba,† Margandan Bhagiyalakshmi,‡ Andrews Nirmala Grace,†,§ Dae Hyun Chu,† Sung Chan Nam,† Yeoil Yoon,† Sung Ho Yoon,∥ and Soon Kwan Jeong*,† †

Korea Institute of Energy Research, Daejeon 305-343, Korea Department of Chemistry, Central University of Kerala, Kasaragod 671-314, India § Centre for Nanotechnology Research, VIT University, Vellore632 014, India ∥ Department of Bio & Nano Chemistry, Kookmin University, Seoul 136 702, Korea ‡

S Supporting Information *

ABSTRACT: One aspect of the attempt to restrain global warming is the reduction of the levels of atmospheric CO2 produced by fossil fuel power systems. This study attempted to develop a method that reduces CO2 emissions by investigating the absorption of CO2 into sterically hindered amine 2-amino-2-methyl-1-propanol (AMP), the acceleration of the absorption rate by using the enzyme carbonic anhydrase (CA), and the conversion of the absorption product to stable carbonates. CO2 absorbed by AMP is converted via a zwitterion mechanism to bicarbonate species; the presence of these anions was confirmed with 1H and 13 C NMR spectral analysis. The catalytic efficiency (kcat/Km), CO2 absorption capacities, and enthalpy changes (ΔHabs) of aqueous AMP in the presence or absence of CA were found to be 2.61 × 106 or 1.35 × 102 M−1 s−1, 0.97 or 0.96 mol/mol, and −69 or −67 kJ/mol, respectively. The carbonation of AMP-absorbed CO2 was performed by using various Ca2+ sources, viz., CaCl2 (CAC), Ca(OOCCH3)2 (CAA), and Ca(OOCCH2CH3)2 (CAP), to obtain various polymorphs of CaCO3. The yields of CaCO3 from the Ca2+ sources were found in the order CAP > CAA > CAC as a result of the effects of the corresponding anions. CAC produces pure rhombohedral calcite, and CAA and CAP produce the unusual phase transformation of calcite to spherical vaterite crystals. Thus, AMP in combination with CAA and CAP can be used as a CO2 absorbent and buffering agent for the sequestration of CO2 in porous CaCO3.

1. INTRODUCTION

capture, utilization, and storage (CCUS) approaches, which stress the utilization of CO2. The sequestration of CO2 in mineral carbonates has recently been extensively studied by many researchers; these approaches utilize aqueous saturated CO2 solutions.12−14 Calcium carbonate has commercial application as a filler in paints, paper, plastics, rubber, and cement, and is also a constituent of many biological, geological, and ecological systems. It occurs in six polymorphic phases, viz., calcite, aragonite, vaterite, monohydrate, hexahydrate, and amorphous CaCO3 (ACC).15,16 ACC, a metastable transitional phase formed during cluster

The emission of carbon dioxide from coal combustion and fossil fuel processes is predicted to produce increases in the atmospheric CO2 concentration from 3981,2 to 450 ppm and in the global temperature of more than 2.8 °C by 2100.3 To reduce CO2 emissions, various methods have been identified as carbon dioxide capture and storage (CCS) processes.4−7 CCS approaches have the drawback that CO2 leaked from storage sites could migrate to the land surface and affect soil conditions, groundwater, plant growth, and so on.8,9 Other drawbacks of CCS processes include rapid leakage during earthquakes,10 their energy consumption leading to new carbon emissions, the absence of byproduct output, and poor cost effectiveness.11 An encouraging new technological strategy is provided by carbon © 2013 American Chemical Society

Received: July 16, 2013 Revised: November 29, 2013 Published: December 2, 2013 15655

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Figure 1. Biocatalytic activities of CA in CO2 absorption by AMP vs time: (A) absence of CA (control) and (B) with CA, where [AMP] = 1.5 × 10−2 M, [CO2] = 5 × 10−3, 1.0 × 10−2, and 1.5 × 10−2 M, and [CA] = 1.862 × 10−5 M. Insets show the corresponding Lineweaver−Burk plots.

monohydrate, and calcium nitrate. In this study, three organometallic calcium salts, namely, calcium chloride (CAC), calcium acetate (CAA), and calcium propionate (CAP), were employed as calcium precursors and investigated for the carbonation of hydrated CO2 in AMP in the presence of CA to obtain various polymorphs of CaCO3. The effects of varying the calcium source on the AMP sequestration process were analyzed in detail in terms of the yield, polymorphs, and textural properties of CaCO3.

aggregation, is imperative for the growth process because it is a precursor of other phases and thus plays a vital role in the process of biomineralization. Calcite and aragonite commonly occur in nature, and they are the most thermodynamically stable structures. Vaterite is less abundant in nature because of its rapid transformation to aragonite and calcite, because the least thermodynamically stable.17 The larger specific surface and porous nature of vaterite are found to improve the mechanical properties of the material when used as a filler. Vaterite usually coexists with the other two polymorphs, calcite and aragonite. It is a challenging process for the synthesis of stable vaterite from CO2 emission sources. The vaterite phase depends on the calcium source and mole ratio of the reaction medium. Thus, the production of CaCO3 in a specific polymorph in its pure form can be used in industrial applications. In this study, we have attempted the sequestration of CO2 to the vaterite form through its absorption using a sterically hindered amine sorbent. A crucial aspect of these sequestration processes is the enhancement of CO2 solubility by water-soluble amine solvents compared to water as a result of its low solubility.18 Alkanolamine solvents are very efficient and selective CO2 absorbents for postcombustion processes.19 A variety of amine solvents have been identified for CO2 absorption,20 with CO2 absorption rate constants and loading capacities of MEA > DEA > AMP > MDEA and AMP > MDEA > DEA > MEA, respectively.21,22 These results suggest that sterically hindered amine 2-amino-2-methyl-1-propanol (AMP) should have a high CO2 absorption capacity because of the steric effects of AMP and should enable the direct conversion of absorbed CO2 to bicarbonates via the hydration of CO2. Enzyme carbonic anhydrase (CA) catalyzes the hydration of CO2 to bicarbonate and is reported to have the highest catalytic activity in the biological world.23 The plausible species formed in the hydration process are carbonic acid, bicarbonate, and carbonate.24 CA was shown to enhance the rate of CO2 absorption by alkanolamine.25 CO2 absorbed in AMP can be transformed into carbonate minerals by using bivalent alkali earth metals under atmospheric conditions. The other different calcium precursors available are calcium chloride, calcium acetate, calcium propionate, calcium oxalate, calcium acetylacetonate, calcium 2-ethylhexanoate, calcium D-gluconate

2. MATERIALS AND METHODS 2.1. Materials. 2-Amino-2-methyl-1-propanol (AMP), thymol blue sodium salt, calcium acetate (CAA), calcium propionate (CAP), calcium chloride (CAC), and bovine carbonic anhydrase (CA) were obtained from Sigma-Aldrich. A 30% CO2 (balance N2) gas cylinder was purchased from Special Gas, Korea, and used to prepare saturated CO2 solutions and provide the absorption capacity for differential reaction calorimetry (DRC) analysis. All solutions were freshly prepared with Milli-Q water and filtered through a 0.22 μm membrane. 2.2. Stopped-Flow Spectrophotometric Study. The absorption rate of CO2 by AMP was estimated in the presence or absence of carbonic anhydrase using an Applied Photophysics SX-20 spectrophotometer. Typically, the decay of the thymol blue indicator intensity at 596 nm was monitored during the reaction between a mixture of 15 mM AMP and 12.5 μM thymol blue indicator and with various concentrations (5, 10, and 15 mM) of saturated CO2 in water solution at 25 °C in the presence or absence of CA .26 The stopped-flow measurements were repeated at least five times, and the average results obtained were fitted to the curve using Pro-Data Viewer software for the analysis of the results. Kinetic parameters kcat and Km values were calculated using Michaels−Menten and Lineweaver−Burk equations.13 2.3. DRC Measurements of the CO2 Absorption Loading. The CO2 absorption capacities and the heat of absorption of AMP were evaluated by using a DRC system.27 Briefly, the DRC apparatus consisted of two mechanically agitated reactors with a temperaturecontrolled water bath that was circulated through the double-layer reactors (250 mL) to maintain the reaction temperature. Each reactor contained a Joule effect calibration probe and thermocouple (±0.01 °C) for measuring the reaction temperature (ΔT). Each reactor was filled with 150 g of AMP solution (10 wt %), stirred at 250 rpm, and maintained at 25 °C. The initial system was measured via the Joule effect calibrated in triplicate, and 30% CO2 gas was supplied in a purged form at 150 cc. The CO2 absorption capacity, heat of reaction, and reaction completion were evaluated by analyzing the reactor vent 15656

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gas with an online gas chromatograph equipped with a TCD detector. The reaction completion was observed on the basis of the outlet gas concentration reaching 30% CO2, at which point the flow of purging CO2 gas was stopped and the system was analyzed again via triplicate calibration to calculate the heat of absorption of CO2 by AMP. The above reaction was performed in the presence of CA (5 μg/mL). 2.4. Sequestration of Absorbed CO2 to CaCO3. The saturated absorbed CO2 solution (0.96 mol/mol) obtained after streaming 30% CO2 gas through the 100 mL aqueous AMP solution (10 wt %) at 25 °C for 5 h in the DRC system was used in sequestration processes with various Ca2+ sources (CAC, CAA, and CAP). Typically, 10 mL of CO2-absorbed AMP solution was added to 10 mL of CAC solution (0.3707 g of CAC in 10 mL water (1:1 mol/mol CO2/Ca2+), and sequestration was carried out on a temperature-controlled shaker at 100 rpm (Lab Companion, SIF 6000). The carbonization processes were carried out at the CO2/Ca2+ ratios of 1:0.5, 1:1, and 1:2 over 60 min at 25 °C. The precipitated CaCO3 was filtered and dried. The sequestration yields were quantified by using ion chromatography (IC),28 and the effects of using the different Ca2+ sources were determined through the characterization of the dried CaCO3.

groups that are attached to tertiary carbon atoms. Substituent groups can affect the donation or withdrawal of electrons via steric hindrance effects. The carbamates produced by AMP are unstable, so a hydrolytic reaction occurs that converts the carbamate to bicarbonate.29 The free amines are then available to react with new CO2 molecules. CA enzyme accelerates the hydration of CO2 to form bicarbonate species because Zn− HOH in CA is ionized to a Zn−OH− species that abstracts C from CO2 to produce HCO3−. The final products obtained through reactions of CO2 with AMP (R = C(CH3)2CH2OH) and CA are the same because of the conversion of CO2 to HCO3−. Thus, CA is an effective promoter of CO2 absorption in AMP solutions, which accelerates the absorption rate and does not reduce the CO2 loading capacity. The formation of bicarbonate is verified by the 1H and 13C NMR spectra in Figure S1. The 1H NMR spectra for pristine and CO2-absorbed AMP show that absorption results in chemical shifts of the triplet signal at δ = 3.34 to 3.52 ppm (−CH2OH) and the signal due to methyl protons at δ = 1.06 to 1.27 ppm. The 13C NMR spectra for pure and CO2-absorbed AMP show that absorption results in chemical shifts of the signals due to the carbon atoms in the −NC(CH3)2, −CNH2, and −CH2OH groups from 25.6, 50.3, and 71.8 ppm to 22.4, 55.4, and 67.3 ppm, respectively. A new peak is present at δ = 161.5 ppm in the spectrum of the CO2-absorbed AMP system, which demonstrates that HCO3− species form during the absorption process. Note that in the 13 C NMR spectra we cannot distinguish between HCO3− and CO32− because of the rapid proton exchange between AMPH+ and AMP, so there is only a single peak for these species at δ = 161.5 ppm.30 We conclude that CO2 absorption by the AMP + H2O + CO2 system does not produce carbamates, which would produce a peak at δ = 164 ppm.31 Therefore, the single peak at δ = 161.5 indicates that CO2 absorption by the AMP system proceeds along the following pathway:

3. RESULTS AND DISCUSSION 3.1. Estimation of Enzymatic Activity Using StoppedFlow Spectrometry. The catalytic efficiency (kcat/Km) of CA was quantified using stopped-flow spectrometry by monitoring the decay intensity of thymol blue at 596 nm during the reaction between the mixture of AMP + thymol blue indicator and different concentration of aqueous CO2 in water at 25 °C. Here, the decay of thymol blue is directly proportional to the formation of bicarbonate. Figure 1 shows the concentration profile of the absorption of CO2 by AMP versus time, where the conversion rate rapidly increases in the first 0.15 s (Figure 1B) and gradually increases after 0.5 s, which indicates the steady absorption rate of CO2 by AMP in the presence of CA. The control system in the absence of CA shows a slow increase in the absorption rate even after 3 s (Figure 1A). The initial velocity (V0) was obtained using exponential curve fitting, and the catalytic efficiency was calculated by applying the Michaels−Menten (eq 1) and Lineweaver−Burk equations (eq 2). The catalytic efficiency (kcat/Km) of CO2 absorption by AMP is 1.35 × 102 and 2.61 × 106 M−1 s−1 with respect to the absence and presence of CA (per mole), respectively, as obtained from the insets in Figure 1A,B of double-reciprocal Lineweaver−Burk plots. Thus, it is observed that CA accelerates the rate of CO2 absorption in the AMP−water− CO2 system. kcat[CO2 ][E0] k m + [CO2 ]

(1)

k 1 1 1 = m + V0 Vmax [CO2 ] Vmax

(2)

V0 =

RNH 2 + CO2 → RNHCOO− + H+

(3)

RNHCOO− + H 2O → RNH 2 + HCO3− (hydrolytic reaction)

(4)

The absorption capacities and the heat of absorption of CO2 by aqueous AMP were quantified by performing DRC with and without CA. The DRC system comprises sample and reference reactors that can be maintained under adiabatic conditions. A refrigerated heating circulator (Julabo, HE V-2) was used to circulate water in the outer layer of the reactor and maintain the temperature at 25 °C. Each reactor contains the same amount of 10% AMP solution, and the CO2 loading capacity was quantified from vent gas using gas chromatography with a Porapak-Q column. Figure S2 shows the progress of CO2 absorption by AMP, which was monitored until reaction completion to reach the initial CO2 gas concentration in the outlet gas of 30 wt % (balance N2). The CO2 absorption capacities of AMP were found to be 0.97 and 0.96 mol/mol in the presence and absence of CA, respectively. In general, CA should accelerate the rate of absorption of CO2 in amine systems,25 whereas with the AMP medium no such significant increase in CO2 absorption loading is observed and almost the same absorption capacities in both the presence and absence of CA at 25 °C is exhibited (Figure S2). The reasons for the propensity of the AMP + CA system might be due to the production of bicarbonate anions. Therefore, the CO2 loading depends only on the properties of AMP. Furthermore, the inset

where V0 is the absorption rate of CO2 by AMP, Vmax is the maximum rate, kcat is the catalytic rate constant, [E0] is the enzyme concentration, [CO2] is the substrate concentration, Km is the substrate concentration when the rate is equal to Vmax/2, and kcat/Km is the catalytic efficiency. 3.2. Mechanism of CO2 Loading as an Anionic Species in AMP. The absorption of CO2 into an aqueous AMP solution was carried out by using CO2 gas to purge the AMP solution at 25 °C for 5 h. CO2 absorbed by AMP undergoes an acid−base interaction to form zwitterions. A plausible mechanism for this process is as follows. The interaction of the sterically hindered amine AMP with CO2 produces carbamates at the amino 15657

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CAP solution, acetic acid or propionic acid is formed, which prevents the further precipitation of CaCO3.40 AMP converts these anions into ammonium carboxylates, which prevents the formation of free acetic acid or propionic acid and thus favors the formation of CaCO3. The mineralization reactions were also investigated for CO2/ Ca2+ ratios of 1:0.5, 1:1, and 1:2. Figure 2 show that the yields

in Figure S2 shows the temperature differences between the reactors, which were determined using Joule effect calibration and calculated as the heat of CO2 absorption. AMP during the absorption process forms unstable CO2− followed by hydrolysis to give HCO3−. The latter process releases a large quantity of heat in the absence of CA. When CA is introduced, the CO2 absorption rate is enhanced and only a small amount of absorption heat is released, perhaps because of the intermolecular proton transfer between AMP and CA, as is observed in salting in/out effects in aqueous amine systems.32 Indeed, the denaturation of CA is negligible in dilute amine− water systems (10%) and the pKa is 9.67,33 even though CA exhibits high catalytic activity at pH 10.534 and maintains the tertiary structure of the enzyme between 25 and 40 °C.35 Therefore, the quantity of heat released during the absorption of CO2 by AMP can be determined as a function of time by measuring the temperature difference between the sample and reference reactors. The quantity of heat (Q) and heat of absorption (ΔHabs) in the presence and absence of CA were found to be −11.09, −11.01 kJ, and −69, −67 kJ/mol, respectively, as obtained from the inset in Figure S2. On the basis of the above results, we conclude that CA accelerates the absorption rate but not the loading capacity, and we carried out the following carbonation reaction in the CO2-saturated AMP solution. 3.3. Carbonization of Anionic Species as CaCO3 by Various Ca2+ Sources. Harvesting stable carbonates from flue gases could be a promising method in the reduction of carbon emission in postcombustion processes. The solubility of CO2 in water is 0.149 wt % at 25 °C,18 and CO2(aq) equilibrates with water to form carbonic acid. The CO2 absorption capacity of an aqueous AMP (10 wt %) solution is 4.17 wt % (according to the DRC analysis). The transformation of CO2 to CO32− proceeds through a zwitterion mechanism, and here AMP produces RNH3+ and HCO3− species and subsequently abstracts H+ by OH− at higher pH to produce CO32−. Ca2+ was obtained by reacting calcium salts with RNH3+. These carbonates are converted to CaCO3 by using calcium sources CAC, CAA, and CAP. The carbonation process proposed here consists of the following steps:36 RNH 2 + CO2 + H 2O → RNH3+ + HCO3−

(5)

RNH 2 + H 2O → RNH+3 + OH−

(6)

Figure 2. Profiles of the sequestration yield of AMP-hydrated CO2 obtained with the various calcium sources.

of CaCO3 with a 1:0.5 mol ratio of CO2/Ca2+ for calcium sources CAC, CAA, and CAP were 42, 46, and 50%, respectively, corresponding to initial CO2 absorbed by AMP (0.96 mol/mol). The low yield of CAC is due to the acidity of the byproduct chloride anions. When the CO2/Ca2+ ratio is increased to 1:1 and 1:2, the yield of CaCO3 increases; of the three Ca2+ sources, CAP produces the maximum yield (75%) for a CO2/Ca2+ ratio of 1:2, and CAA and CAC produce yields of around 72 and 61%, respectively. Figure 2 also shows the amounts of CO2 absorbed by AMP−CO2 solution (0.96 mol/ mol) with CAC, CAA, and CAP, which were found to be 0.61, 0.69, and 0.72 mol/mol of CaCO3, respectively. The order of utilization of the initial Ca2+ ions for a CO2/Ca2+ ratio of 1:2 was found to be CAP > CAA > CAC. This result demonstrates that organic-based Ca2+ salts are good Ca2+ sources for CO2 sequestration in amine absorption processes because their anion byproducts are weaker acids than chloride ions. The quantity of precipitated CaCO3 was quantified by using IC method.28 The crystal phases of the precipitated CaCO3 samples were identified and quantified by using XRD (Figure 3). The phase compositions were calculated by using the relative intensity ratio (RIR) method described elsewhere.41 Figure 3a−c,d−f shows the XRD patterns of CaCO3 obtained with CAA and CAP, respectively, for CO2/Ca2+ ratios of 1:0.5; 1:1, and 1:2. Figure 3g shows the XRD pattern for CAC for a CO2/Ca2+ ratio of 1:2. It was determined with the RIR method that CAA and CAP promote the formation of vaterite over calcite phases to a larger extent than CAC even for low CO2/Ca2+ ratios. The normal polymorph phase transformation of CaCO3 is vaterite to calcite.42 The reverse phase transformation was observed in the CAA and CAP systems as a result of the electrostatic interactions between surface layers containing two orientations of the CO32− groups, and the calcium-rich surface promotes the crystallization of calcite.42,43 Second, the rhombohedrally shaped calcite could be converted to spherical vaterite upon increasing the number of Ca2+ ions because the surface calcium ions attract a layer of carbonate ions that are oriented so as to

2RNH3+ + CaX 2 → 2RNH 2 + Ca 2 + + 2HX (X = Cl−, CH3COO− , C2H5COO−)

2Ca 2 + + 2HCO3− + 2OH− → 2CaCO3 + 2H 2O

(7) (8)

The byproducts of the sequestration of amine-absorbed CO2 with various Ca2+ sources in CaCO3 are HCl, CH3COOH, and C2H5COOH. The pKa values of HCO3−, C2H5COO−, CH 3 COO −, and Cl − are 6.35, 4.87, 4.76, and −6.3, respectively.37 The acidic character produced by the mineralization reactions is expected to affect the yield of CaCO3 formation because acids convert this carbonate to its soluble salts. The aforementioned conversion of carbonate to soluble salts is prevented by the buffering action of AMP. Here, AMP plays a dual role as a CO2 absorber and a buffering agent for the carbonation reaction in the form of the AMP−hydrochloride (R−NH3+Cl−) salt or ammonium carboxylates ([R−NH3+][−OOCR′]).38,39 When CaCO3 precipitates from a CAA or 15658

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3.62 (3.47) 3.39 (3.89) 3.48 (3.74) a

The values indicate that the different calcium source are 1:0.5 (CO2/Ca2+) and 1:2 (within parentheses).

9.3 (46) 1.6 (81.5)

Ca 2p1/2 − Ca 2p3/2 (eV) O 1s/C 1s

1.11 (2.01) 1.33 (1.95) 1.28 (1.57) 0.37 (0.68) 0.43 (0.62) 0.14 (0.48)

Ca 2p/C 1s O 1s/Ca 2p

2.97 (2.96) 3.09 (3.16) 3.10 (3.27) 165.1 (169.3) 130.4 (107.4) 135.3 (112.0)

ΔHd (kJ/mol) peak point (°C)

759.3 (757.1) 761.1 (747.2) 762.7 (755.6) 43.3 (44.1) 44.0 (43.1) 44.1 (42.8)

CO2 loss (wt %) Dp (nm)

28 (23) 48 (37.5) 31.7 (25.5) 0.4 (0.9) 4.0 (32) 37 (24)

SBET (m /g)

CAC CAA CAP

vaterite (%) calcite (%)

100 (100) 90.7 (54) 98.4 (18.5)

calcium sources

thermal properties by TGA-DSC textural properties by BET

Table 1. Profiles of Precipitated CaCO3 Obtained with Various Calcium Sourcesa

facilitate enhanced vaterite formation. In the carbonate plane, the surface CO32− groups rotate to lie flat on the surface and form vaterite particles.44 The precipitated CaCO3 crystal phases were matched with the corresponding JCPDS data files and were found to be calcite (47-1743) and vaterite (33-0268).45 The XRD peaks of the calcite phase obtained (1:0.5) from CAA and CAP have 2θ values of 22.94, 29.28, 35.86, 39.28, and 43.02°, and their respective crystal faces are (1 0 4), (0 1 2), (1 1 0), (1 1 3), and (2 0 2). The intensity of the main calcite peak at 29.28° (1 0 4) decreases with an increase in the Ca2+ ion concentration (1:2), and new peaks emerge at 20.9, 24.8, 27.8, 32.7, 43.8, and 50.07°, which match vaterite JCPDS data file 330268 and correspond to crystal faces (0 0 4), (1 1 0), (1 1 2), (1 1 4), (3 0 0), and (1 1 8), respectively. Table 1 shows the transformation of the calcite phase to the vaterite phase by the CAA and CAP system when the CO2/Ca2+ ratio is increased from 1:0.5 to 1:2; the proportion of the vaterite phase increases from 46 to 81.5%. The CAC system produces pure calcite for all mole ratios of CO2/Ca2+ because this rapid phase transformation produces stable rhombohedral calcite (Figure 3g). N2 adsorption−desorption isotherms were also obtained for the CaCO3 samples precipitated from AMP + CO2 with CAC, CAA, and CAP, and the results are summarized in Table 1. The surface areas of these materials were quantified with the BET method,46,47 and the pore sizes were evaluated with the BJH method (isotherms with hysteresis). The CaCO3 precipitated at a CO2/Ca2+ ratio of 1:2 has a higher specific area and a lower pore diameter than that at a ratio of 1:0.5 (Table 1); this difference in surface area might be due to the significant changes in the textural properties of CaCO3 during crystal formation. It was also found that the CaCO3 samples obtained from CAA and CAP are porous (p/p0 = 0.45−0.98). Figure S3 shows the well-defined type IV hysteresis loops in the partial pressure (p/p0) for the CAP and CAA systems for a CO2/Ca2+ ratio of 1:2, which indicates the formation of porous CaCO3. In contrast, the CAC system produces pure calcite phases at all concentrations (1:0.5 and 1:2), and the surface area and pore volume are nearly constant. Thus, Ca2+ sources CAA and CAP produce porous CaCO3 in AMP + CO2 sequestration processes. Figures S4 and 4 show FE-SEM images of the surface morphologies of CaCO3 precipitated from AMP-absorbed CO2 with various Ca2+ sources and CO2/Ca2+ mole ratios of 1:0.5, 1:1, and 1:2. CaCO3 precipitated with CAA at a CO2/Ca2+ ratio

polymorphs by XRD

Figure 3. XRD patterns of precipitated CaCO3 obtained at various CO2/Ca2+ ratios: (a−c) CAA, (d−f) CAP, and (g) CAC.

2

relative intensities ratio and BE difference between core levels of Ca 2p by XPS

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Figure 4. FE-SEM images of CaCO3 precipitated with different calcium sources at (A, C, E) 1:2 CO2/Ca2+ and (B, D, F) related HR-TEM images.

of 1:0.5 contains a combination of two different crystal shapes, viz., bricklike calcite following a brick-by-brick formation mechanism48 and spherical vaterite crystals (Figure S4A), whereas at CO2/Ca2+ ratios of 1:1 and 1:2 homogeneous asteroid-shaped particles were observed (Figures S4B and 4A). The CaCO3 obtained with CAP consists of coexisting calcite and vaterite crystals in the form of bricklike, asteroid, and porous spherical vaterite particles for CO2/Ca2+ ratios of 1:0.5, 1:1, and 1:2, respectively (Figures S4C,D and 4C) whereas the CAC system reveals calcite uniquely at all ratios (Figures S4E,F and 4E). In general, although the vaterite polymorph is spherical, it can also be synthesized in other morphologies such as hexagonal plates, disk, and snow and flower types depending on the methodology adopted. Diverse routes have been reported for the synthesis of spherical vaterite by adding additives,49 polymers,50 surfactants,51 and varying reaction parameters.52 This study reveals the influence of the calcium source (CAP and CAA) wherein propionate and acetate anions might influence the crystallization to obtain the porous vaterite form of CaCO3. The porous nature of vaterite was also verified with BET and TEM studies. HR-TEM images of the CAA and CAP system revealed the formation of CaCO3 with a porous nature at a 1:2 mol ratio, whereas the CAC system produced nonporous calcite as given in Figure 4B,D,F. The thermal stability of polymorphs of CaCO3(s) was studied by TGA-DSC analysis (Figure S5). CaCO3 decomposes to CaO and CO2, and an endothermic weight loss of about 43− 44% by mass in the range of 600−780 °C was observed.53,54 This reveals that the release of CO2 on thermal decomposition depends on the physical form and particle size of CaCO3. The enthalpy of decomposition (ΔHd) of precipitated CaCO3

obtained from DSC analysis was found to be 107−169 kJ/ mol (Table 1). The enthalpy of decomposition depends on crystal polymorphs because it is expected that lattice energies of rhombohedral calcite are higher than those of porous vaterite.44 Thus, on increasing the Ca2+ ion concentration from 1.0:5 to 1:2 in CAA and CAP systems, a decrease in the endothermic enthalpy values was observed, which is attributed to the phase transformation of calcite to vaterite. The CAC system showed similar values of ΔHd in all ratios (CO2/Ca2+), producing a pure calcite phase (Table 1). Elemental analysis of the produced CaCO3 from AMP revealed no contamination because the results showed 12 wt % carbon and the absence of nitrogen. The polymorphs of CaCO3 precipitated with CAC, CAA, and CAP were examined with FT-IR (Figure S6). The calcite spectrum contains vibrational bands at 712, 876, and 1421 cm−1, and the vaterite spectrum contains peaks at 746, 878, 1088, 1412, and 1495 cm−1.55 There is an increase in Ca2+ content in the CAA and CAP systems with respect to the precipitation of CaCO3 (1:2), producing more intense symmetric (ν1) and asymmetric (ν3) C−O stretching at 1085.7 cm−1 and at 1415.1 and 1494.2 cm−1, respectively, and decreasing out-of-plane bending (ν4) of C−O at 711.7 cm−1. In addition, CO32− bending (ν2) and O−C−O bending (ν4) produce characteristic bands at 878.1 and 746 cm−1, respectively (Figure S6a−d). The spectrum of CaCO3 formed in the CAC system contains peaks at 711.7, 876.2, and 1421.9 cm−1, which are assigned to the C−O bond out-of-plane (ν4) and in-plane bending (ν2) and stretching vibrations (ν3) in calcite (Figure S6e). Hence, the FT-IR spectra confirm that the CAC system produces pure calcite and that coexisting calcite 15660

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Figure 5. Deconvoluted XPS spectra of precipitated CaCO3 obtained with CAC.

splitting of two Ca 2p (2p1/2−2p3/2) core levels (i.e., 3.62 and 3.47 eV, respectively) (Table 1). However, increasing the number of Ca2+ ions in the CAA and CAP system reveals an increasing energy gap for the Ca 2p core as shown in Table 1. These results suggested that the CAA and CAP system yields a vaterite phase at 1:2. The O 1s core level of CO3 in each of the three systems’ CaCO3 polymorphs was found in the XPS results at 531.2 ± 0.3 eV, as shown in Figures 5 and S8 and S9. The high-energy resolution of CaCO3 from the CAC, CAA, and CAP system exhibited an apparent lower binding energy of CO peaks at 2 eV from the relative intensity of the main O 1s. The relative intensity ratios of C 1s, Ca 2p, and O ls of precipitated CaCO3 are presented in Table 1, which shows that the O/Ca ratio of the CAC system matched the reported data for the calcite phase at both 1:0.5 and 1:2 ratios, respectively.59 In other cases, because of the coexistence of both phases the ratio was observed to be smaller (∼3.2 eV) than that of pure vaterite.59 Hence, AMP and organic calcium are suitable materials for CO2 absorption processes and subsequent sequestration because of their high CO2 absorption capacities, low enthalpies, and cost effectiveness and can be used to convert industrial CO2 emissions into stable CaCO3.

and vaterite phases are obtained from both the CAA and CAP systems. Figure S7 shows the XPS spectra for CaCO3 precipitated with CAC, CAA, and CAP at CO2/Ca2+ mole ratios of 1:0.5 and 1:2, which reveal the nature of the surface interfacial bonds formed during crystallization. The Ca 3p, Ca 3s, C 1s, Ca 2p, Ca 2s, and O 1s binding energies were observed at 26, 45, 289, 347, 440, and 531 eV, respectively.56−58 Figures 5 and S8 and S9 show high-resolution spectra of C 1s, Ca 2p, and O 1s, which reveal the core-level characteristics of CaCO3. The main C 1s core level of CO3 is at 289.4 eV, and the adventitious carbon relative intensities with respect to the main CO3 peak appear at 8.7 ± 1.5 (Cx) and 6 ± 1.5 eV (Cy) lower binding energies, respectively. The adventitious carbon (Cx) intensity is higher at a CO2/Ca2+ ratio of 1:0.5 than at 1:2 because there are insufficient Ca2+ ions to form CaCO3 so CO3 ions become surrounded by several layers of excess adventitious carbon of AMP that accumulates at the surface. At a CO2/Ca2+ ratio of 1:2, the Cx intensity decreases with an increase in Ca2+ as the number of excess CO3 ions decreases. This clearly supports the adventitious carbon layer formation on CaCO3 and lowers the shifting of the binding energy of Cx from C ls (289.4 eV) for the CAC, CAA, and CAP systems to 3.79, 3.88, and 4.14 eV, respectively (Figures 5 and S8 and S9). This result reveals that at a CO2/Ca2+ ratio of 1:2 in the CAC system the Cl− anions is not involved in the surface layers and showed a lower shifting than did others. The high-energy resolution of the Ca 2p core level of the polymorphic CaCO3 samples indicates that the binding energies of Ca 2p3/2 and Ca 2p1/2 are 346.8 ± 0.3 and 350.3 ± 0.3 eV, respectively, as shown in Figures 5 and S8 and S9. The increase in the proportion of Ca2+ ions in the CAC system from 1:0.5 and 1:2 shows a smaller difference in the spin−orbit

4. CONCLUSIONS The activity of CA in the absorption of CO2 by AMP was exactly quantified using stopped-flow spectrometry. The catalytic efficiency (kcat/Km) of aqueous AMP in the presence or absence of CA was found to be 2.61 × 106 or 1.35 × 102 M−1 s−1, respectively, which clearly shows that CA enhances the CO2 absorption. The carbon dioxide absorption capacities and heat of reaction of AMP were found by measurements with a DRC system to be unaffected by the presence or absence of the 15661

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(8) Zhou, X.; Apple, M. E.; Dobeck, L. M.; Cunningham, A. B.; Spangler, L. H. Observed response of soil O2 concentration to leaked CO2 from an engineered CO2 leakage experiment. Int. J. Greenhouse Gas Control 2013, 16, 116−128. (9) Amonette, J. E.; Barr, J. L.; Dobeck, L. M.; Gullickson, K.; Walsh, S. J. Spatiotemporal changes in CO2 emissions during the second ZERT injection, August−September 2008. Environ Earth Sci. 2010, 60, 263−272. (10) Steven, M. D.; Smith, K. L.; Colls, J. J. Environmental Risks and Impacts of Carbon Dioxide (CO2) Leakage in Terrestrial Ecosystems. In Developments and Innovation in Carbon Dioxide (CO2) Capture and Storage Technology; Maroto-Valer, M. M., ed.; Woodhead Publishing: Cambridge, U.K., 2010; Vol. 2, pp 324−343. (11) Rogelj, J.; McCollum, D. L.; Reisinger, A.; Meinshausen, M.; Riahi, K. Probabilistic cost estimates for climate change mitigation. Nature 2013, 493, 79−83. (12) Mirjafari, P.; Asghari, K.; Mahinpey, N. Investigating the application of enzyme carbonic anhydrase for CO2 sequestration purposes. Ind. Eng. Chem. Res. 2007, 46, 921−926. (13) Vinoba, M.; Lim, K. S.; Lee, S. H.; Jeong, S. K.; Alagar, M. Immobilization of human carbonic anhydrase on gold nanoparticles assembled onto amine/thiol-functionalized mesoporous SBA-15 for biomimetic sequestration of CO2. Langmuir 2011, 27, 6227−6234. (14) Hwang, E. T.; Gang, H.; Chung, J.; Gu, M. B. Carbonic anhydrase assisted calcium carbonate crystalline composites as a biocatalyst. Green Chem. 2012, 14, 2216−2220. (15) Yao, Y.; Dong, W.; Zhu, S.; Yu, X.; Yan, D. Novel morphology of calcium carbonate controlled by poly(L-lysine). Langmuir 2009, 25, 13238−13243. (16) Ihli, J.; Kulak, A. N.; Meldrum, F. C. Freeze-drying yields stable and pure amorphous calcium carbonate (ACC). Chem. Commun. 2013, 49, 3134−3136. (17) Grasby, S. E. Naturally precipitating vaterite (μ-CaCO3) spheres: unusual carbonates formed in an extreme environment. Geochim. Cosmochim. Acta 2003, 67, 1659−1666. (18) Dodds, W.; Stutzman, L.; Sollami, B. Carbon dioxide solubility in water. Ind. Eng. Chem. Chem. Eng. Data Series 1956, 1, 92−95. (19) Rochelle, G. T. Amine scrubbing for CO2 capture. Science 2009, 325, 1652−1654. (20) Puxty, G.; Rowland, R.; Allport, A.; Yang, Q.; Bown, M.; Burns, R.; Maeder, M.; Attalla, M. Carbon dioxide postcombustion capture: a novel screening study of the carbon dioxide absorption performance of 76 amines. Environ. Sci. Technol. 2009, 43, 6427−6433. (21) Vaidya, P. D.; Kenig, E. Y. CO2-alkanolamine reaction kinetics: a review of recent studies. Chem. Eng. Technol. 2007, 30, 1467−1474. (22) Vinoba, M.; Bhagiyalakshmi, M.; Grace, A. N.; Kim, D. H.; Yoon, Y.; Nam, S.; Baek, I.-H.; Jeong, S. K. Carbonic anhydrase promotes the absorption rate of CO2 in post-combustion processes. J. Phys. Chem. B 2013, 117, 5683−5690. (23) Bond, G. M.; Stringer, J.; Brandvold, D. K.; Simsek, F. A.; Medina, M.-G.; Egeland, G. Development of integrated system for biomimetic CO2 sequestration using the enzyme carbonic anhydrase. Energy Fuels 2001, 15, 309−316. (24) Silverman, D. N.; Lindskog, S. The catalytic mechanism of carbonic anhydrase: implications of a rate-limiting protolysis of water. Acc. Chem. Res. 1988, 21, 30−36. (25) Penders-van Elk, N. J.; Derks, P. W.; Fradette, S.; Versteeg, G. F. Kinetics of absorption of carbon dioxide in aqueous MDEA solutions with carbonic anhydrase at 298K. Int. J. Greenhouse Gas Control 2012, 9, 385−392. (26) Xiang, Q.; Fang, M.; Maeder, M.; Yu, H. Effect of sarcosinate on the absorption kinetics of CO2 into aqueous ammonia solution. Ind. Eng. Chem. Res. 2013, 52, 6382−6389. (27) Lim, J.-A.; Kim, D. H.; Yoon, Y.; Jeong, S. K.; Park, K. T.; Nam, S. C. Absorption of CO2 into aqueous potassium salt solutions of lalanine and l-proline. Energy Fuels 2012, 26, 3910−3918. (28) Vinoba, M.; Bhagiyalakshmi, M.; Jeong, S. K.; Yoon, Y. I.; Nam, S. C. Capture and sequestration of CO2 by human carbonic anhydrase

CA enzyme and to have values of approximately 0.96 mol/mol and −68 kJ/mol, respectively. The bicarbonate formation mechanism during CO2 absorption by AMP was verified by performing 13C NMR and 1H NMR spectral analyses. Calcium sources CAC, CAA, and CAP were tested in sequestration processes and were found to provide yields of CaCO3 of 61, 72, and 75%, respectively. AMP has a dual role as a CO2 absorber and a buffering agent in its chloride or carboxylate salt forms that favor the formation of CaCO3. The polymorph phase compositions of the CaCO3 materials from CAC, CAA, and CAP at a CO2/Ca2+ ratio of 1:2 were 100, 54, and 18.4% calcite (balance vaterite). These results suggested that the abnormal polymorph phase transformation was observed in the CAA and CAP system. FE-SEM was used to show that CAA and CAP produce coexisting calcite and vaterite with asteroid shapes or smooth spherical particles whereas CAC produces rhombohedral calcite particles. We conclude that AMP is a potential candidate for CO2 absorption and sequestration processes in combination with CAA and CAP for synthesizing porous CaCO3 in postcombustion technologies.

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ASSOCIATED CONTENT

S Supporting Information *

1

H and 13C spectra of CO 2 absorption by AMP. CO2 absorption capacities and heats of reaction. Nitrogen adsorption/desorption isotherms of the calcium carbonates. FE-SEM images of CaCO3 precipitated with different calcium sources. TGA/DSC curves of different CaCO3 samples obtained with CAC, CAA, and CAP. FT-IR spectra of the CaCO 3 samples. XPS spectra of the CaCO 3 phases. deconvoluted XPS spectra of precipitated CaCO3 obtained with CAA and CAP. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported through Korea CCS R&D Centre funded by the Ministry of Science, ICT & Future Planning of the Korean Government.

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REFERENCES

(1) www.CO2now.org. (2) Monastersky, R. Global carbon dioxide levels near worrisome milestone. Nature 2013, 497, 13−14. (3) Nick, F. M.; Vieli, A.; Andersen, M. L.; Joughin, I.; Payne, A.; Edwards, T. L.; Pattyn, F.; Van De Wal, R. S. W. Future sea-level rise from Greenland’s main outlet glaciers in a warming climate. Nature 2013, 497, 235−238. (4) Wee, J.-H. A review on carbon dioxide capture and storage technology using coal fly ash. Appl. Energy 2013, 106, 143−151. (5) Zapp, P.; Schreiber, A.; Marx, J.; Haines, M.; Hake, J.-F.; Gale, J. Overall environmental impacts of CCS technologiesA life cycle approach. Int. J. Greenhouse Gas Control 2012, 8, 12−21. (6) Bobicki, E. R.; Liu, Q.; Xu, Z.; Zeng, H. Carbon capture and storage using alkaline industrial wastes. Prog. Energy Combust. Sci. 2012, 38, 302−320. (7) Johnsson, F. Perspectives on CO2 capture and storage. Greenhouse Gas Sci. Technol. 2011, 1, 119−133. 15662

dx.doi.org/10.1021/la403671y | Langmuir 2013, 29, 15655−15663

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Article

covalently immobilized onto amine-functionalized SBA-15. J. Phys. Chem. C 2011, 115, 20209−20216. (29) Xu, S.; Wang, Y.-W.; Otto, F. D.; Mather, A. E. Kinetics of the reaction of carbon dioxide with 2-amino-2-methyl-1-propanol solutions. Chem. Eng. Sci. 1996, 51, 841−850. (30) Ciftja, A. F.; Hartono, A.; da Silva, E. F.; Svendsen, H. F. Study on carbamate stability in the Amp/CO2/H2O system from 13C-NMR spectroscopy. Energy Procedia 2011, 4, 614−620. (31) Jakobsen, J. P.; Krane, J.; Svendsen, H. F. Liquid-phase composition determination in CO2-H2O-alkanolamine systems: an NMR study. Ind. Eng. Chem. Res. 2005, 44, 9894−9903. (32) Le Tourneux, D.; Iliuta, I.; Iliuta, M. C.; Fradette, S.; Larachi, F. Solubility of carbon dioxide in aqueous solutions of 2-amino-2hydroxymethyl-1, 3-propanediol. Fluid Phase Equilib. 2008, 268, 121− 129. (33) Conway, W.; Wang, X.; Fernandes, D.; Burns, R.; Lawrance, G.; Puxty, G.; Maeder, M. Toward the understanding of chemical absorption processes for post-combustion capture of carbon dioxide: electronic and steric considerations from the kinetics of reactions of CO2(aq) with sterically hindered amines. Environ. Sci. Technol. 2013, 47, 1163−1169. (34) Kernohan, J. The pH-activity curve of bovine carbonic anhydrase and its relationship to the inhibition of the enzyme by anions. Biochim. Biophys. Acta 1965, 96, 304−317. (35) Sarraf, N.; Saboury, A.; Ranjbar, B.; Moosavi-Movahedi, A. Structural and functional changes of bovine carbonic anhydrase as a consequence of temperature. Acta Biochim. Pol. 2004, 665−672. (36) Vučak, M.; Perić, J.; Krstulović, R. Precipitation of calcium carbonate in a calcium nitrate and monoethanolamine solution. Powder Technol. 1997, 91, 69−74. (37) http://www.wiley.com/college/pratt/0471393878/student/ index.html. (38) Pinkert, A.; Marsh, K. N.; Pang, S. Alkanolamine ionic liquids and their inability to dissolve crystalline cellulose. Ind. Eng. Chem. Res. 2010, 49, 11809−11813. (39) Prah, J.; Maček, J.; Dražič, G. Precipitation of calcium carbonate from a calcium acetate and ammonium carbamate batch system. J. Cryst. Growth 2011, 324, 229−234. (40) Eloneva, S.; Teir, S.; Salminen, J.; Fogelholm, C.-J.; Zevenhoven, R. Steel converter slag as a raw material for precipitation of pure calcium carbonate. Ind. Eng. Chem. Res. 2008, 47, 7104−7111. (41) López-Periago, A. M.; Pacciani, R.; Vega, L. F.; Domingo, C. N. Monitoring the effect of mineral precursor, fluid phase CO2−H2O composition, and stirring on CaCO3 crystallization in a supercritical: ultrasound carbonation process. Cryst. Growth Des. 2011, 11, 5324− 5332. (42) Xiao, J.; Wang, Z.; Tang, Y.; Yang, S. Biomimetic mineralization of CaCO3 on a phospholipid monolayer: from an amorphous calcium carbonate precursor to calcite via vaterite. Langmuir 2010, 26, 4977− 4983. (43) Yang, B.; Nan, Z. Abnormal polymorph conversion of calcium carbonate from calcite to vaterite. Mater. Res. Bull. 2012, 47, 521−526. (44) de Leeuw, N. H.; Parker, S. C. Surface structure and morphology of calcium carbonate polymorphs calcite, aragonite, and vaterite: an atomistic approach. J. Phys. Chem. B 1998, 102, 2914− 2922. (45) Wojtas, M.; Wołcyrz, M.; Oz̀yhar, A.; Dobryszycki, P. Phosphorylation of intrinsically disordered Starmaker protein increases its ability to control the formation of calcium carbonate crystals. Cryst. Growth Des. 2012, 12, 158−168. (46) Bang, J.-H.; Jang, Y. N.; Kim, W.; Song, K. S.; Jeon, C. W.; Chae, S. C.; Lee, S.-W.; Park, S.-J.; Lee, M. G. Specific surface area and particle size of calcium carbonate precipitated by carbon dioxide microbubbles. Chem. Eng. J. 2012, 198−199, 254−260. (47) Wei, S.-H.; Mahuli, S.; Agnihotri, R.; Fan, L.-S. High surface area calcium carbonate: pore structural properties and sulfation characteristics. Ind. Eng. Chem. Res. 1997, 36, 2141−2148.

(48) Wang, S. S.; Picker, A.; Cölfen, H.; Xu, A. W. Heterostructured calcium carbonate microspheres with calcite equatorial loops and vaterite spherical cores. Angew. Chem. Int. Ed. 2013, 52, 6317−6321. (49) Vinoba, M.; Bhagiyalakshmi, M.; Jeong, S. K.; Nam, S. C.; Yoon, Y. Carbonic anhydrase immobilized on encapsulated magnetic nanoparticles for CO2 sequestration. Chem.Eur. J. 2012, 18, 12028−12034. (50) Dong, W. Y.; Cheng, H. X.; Yao, Y.; Zhou, Y. F.; Tong, G. S.; Yan, D.; Lai, Y. J.; Li, W. Bioinspired synthesis of calcium carbonate hollow spheres with a nacre-type laminated microstructure. Langmuir 2011, 27, 366−370. (51) Chen, Z.; Nan, Z. Effects of 1,3,5-trimethylbenzene on morphology and polymorph of CaCO3 crystals in the presence of SDS. Mater. Res. Bull. 2013, 48, 4319−4328. (52) Chu, D. H.; Vinoba, M.; Bhagiyalakshmi, M.; Baek, I. H.; Nam, S. C.; Yoon, Y.; Kim, S. H.; Jeong, S. K. CO2 mineralization into different polymorphs of CaCO3 using an aqueous-CO2 system. RSC Adv. 2013, 3, 21722−21729. (53) Judd, M. D.; Pope, M. I. Energy of activation for the decomposition of the alkaline-earth carbonates from thermogravimetric data. J. Therm. Anal. 1972, 4, 31−38. (54) Radha, A. V.; Forbes, T. Z.; Killian, C. E.; Gilbert, P. U. P. A.; Navrotsky, A. Transformation and crystallization energetics of synthetic and biogenic amorphous calcium carbonate. Proc. Natl. Acad. Sci. U.S.A. 2010, 107, 16438−16443. (55) Nebel, H.; Neumann, M.; Mayer, C.; Epple, M. On the structure of amorphous calcium carbonate: a detailed study by solid-state NMR spectroscopy. Inorg. Chem. 2008, 47, 7874−7879. (56) Stipp, S. L.; Hochella, M. F., Jr. Structure and bonding environments at the calcite surface as observed with X-ray photoelectron spectroscopy (XPS) and low energy electron diffraction (LEED). Geochim. Cosmochim. Acta 1991, 55, 1723−1736. (57) Baer, D.; Moulder, J. High resolution XPS spectrum of calcite (CaCO3). Surf. Sci. Spectra 1993, 2, 1. (58) Järvinen, L.; Leiro, J. A.; Bjondahl, F.; Carletti, C.; Eklund, O. XPS and SEM study of calcite bearing rock powders in the case of reactivity measurement with HCl solution. Surf. Interface Anal. 2012, 44, 519−528. (59) Ni, M.; Ratner, B. D. Differentiating calcium carbonate polymorphs by surface analysis techniquesan XPS and TOF-SIMS study. Surf. Interface Anal. 2008, 40, 1356−1361.

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