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Jul 22, 2015 - Pacific Northwest National Laboratory, P.O. Box 999, MS K8-96, Richland, ... ABSTRACT: Magnesite precipitation from aqueous solution,...
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Dynamics of Magnesite Formation at Low Temperature and High pCO2 in Aqueous Solution Odeta Qafoku,*,† David A. Dixon,‡ Kevin M. Rosso,† Herbert T. Schaef,† Mark E. Bowden,§ Bruce W. Arey,§ and Andrew R. Felmy†,∥ †

Pacific Northwest National Laboratory, P.O. Box 999, MS K8-96, Richland, Washington 99352, United States Department of Chemistry, The University of Alabama, Shelby Hall, Tuscaloosa, Alabama 35487, United States § The Environmental Molecular Sciences Laboratory, Pacific Northwest National Laboratory, Richland, Washington 99354, United States ∥ Washington State University, Pullman, Washington 99164, United States ‡

S Supporting Information *

ABSTRACT: Magnesite precipitation from aqueous solution, despite conditions of supersaturation, is kinetically hindered at low temperatures for reasons that remain poorly understood. The present study examines the products of Mg(OH)2 reaction in solutions saturated with supercritical CO2 at high pressures (90 and 110 atm) and low temperatures (35 and 50 °C). Solids characterization combined with in situ solution analysis reveal that the first reaction products are the hydrated carbonates hydromagnesite and nesquehonite, appearing simultaneously with brucite dissolution. Magnesite is not observed until it comprises a minor product at 7 days reaction at 50 °C. Complete transition to magnesite as the sole product at 35 °C (135 days) and at a faster rate at 50 °C (56 days) occurs as the hydrated carbonates slowly dissolve under the slightly acidic conditions generated at high pCO2. Such a reaction progression at high pCO2 suggests that over long term the hydrated Mgcarbonates functioned as intermediates in magnesite formation. These findings highlight the importance of developing a better understanding of the processes expected to occur during CO2 storage. They also support the importance of integrating magnesite as an equilibrium phase in reactive transport calculations of the effects of CO2 sequestration on geological formations at long time scale.

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INTRODUCTION Mechanisms of carbonate mineral formation in geochemical systems have been studied for many years, as a result of their abundance in a variety of natural environments. Recently, this topic has taken on an added importance because formation of these phases, most notably calcium and magnesium carbonates, is one of the primary means of permanently sequestering carbon dioxide in deep geologic formations.1−14 However, carbonate mineral formation is often a slow process at lower temperatures, especially lower than 80 °C, conditions consistent with many potentially important geologic CO2 disposal reservoirs.13,15,16 In particular, it has long been known that the precipitation of magnesite (MgCO3) from aqueous solution is kinetically inhibited, presumably by the strong hydration energy of the Mg2+ ion17−19 or as recently suggested by the intrinsic structural barrier of magnesite.20 However, the fundamental details of factors controlling the aqueous formation of magnesite have proven elusive.17,21 Understanding its mechanism and kinetics of formation could be pivotal for predicting the ultimate fate of geologically stored © 2015 American Chemical Society

CO2 and for developing efficient carbon sequestration technologies. Magnesite nucleation and growth from homogeneous aqueous solutions at low temperature and ambient pressures has proven difficult, so previous work has emphasized understanding its dissolution/precipitation, typically using either natural samples or magnesite synthesized at higher temperatures.17,21−31 Influence of other factors such as pH, ionic strength, and pCO2 on the formation of low temperature magnesite has also been investigated.19 One possible pathway to magnesite formation is via intermediate phases. For example, it is generally understood that higher pH and higher alkalinity favors formation of hydromagnesite (Mg5(CO3)4(OH)2· 4H2O), a basic metastable Mg-carbonate phase. Magnesite could subsequently form upon slow dehydration of hydroReceived: Revised: Accepted: Published: 10736

May 27, 2015 July 20, 2015 July 22, 2015 July 22, 2015 DOI: 10.1021/acs.est.5b02588 Environ. Sci. Technol. 2015, 49, 10736−10744

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Environmental Science & Technology

thermogravimetric analysis with mass spectrometer detector (TGA-MS). To address the objectives, Mg(OH)2 aqueous solutions were saturated with CO2 at elevated partial pressures (90 and 110 atm) and at low temperature (35 and 50 °C) in a closed system reacted for extended periods of time (>4 months). In conjunction, a series of in situ experiments were performed in parallel where solution chemistry at high pCO2 was linked with intermediate carbonate solids formation.

magnesite at low CO2 pressure, a possible pathway consistent with magnesite formation in sedimentary environments.32 In salt-rich aqueous solutions, the slow transformation of hydromagnesite to magnesite has been found to occur after hydrothermal treatment, i.e., at elevated temperatures, 110− 200 °C, and at moderate pressures.33,34 This transformation could have proceeded either through dehydration and solidstate recrystallization in brines with low Mg concentration or through dissolution of hydromagnesite coupled with magnesite precipitation in brines characterized by high Mg content.34 Conversion of hydromagnesite to stable magnesite was also found to occur at lower temperatures (55 °C) if hydromagnesite was kept in a pCO2 environment of 1 atm.35 A reduced induction time for magnesite formation was also reported in systems with elevated pCO2.19 Overall, it has been suggested that the intermediate phase hydromagnesite, contingent on pCO2, Mg concentration, and ionic strength, could serve as a possible precursor to magnesite formation.19,23 Besides hydromagnesite, brucite (Mg(OH)2) has also been proposed as a precursor to formation of magnesite.23 However, direct nucleation and growth of magnesite from solution, without formation of intermediate metastable phases, has only been reported in studies conducted at high partial CO2 pressure (99−148 atm) and high temperature (80−120 °C).23,36 With respect to current interest in geological CO 2 sequestration, magnesite formation in aqueous solution at low temperatures and high partial CO2 pressure is an increasingly relevant topic. The first evidence for formation of magnesite at high pCO2 (90 atm) and low temperature (35 °C) was reported recently during reaction of synthetic and natural forsterite (Mg2SiO4) in water-saturated supercritical CO2 (scCO2).11,15 It has also been recently demonstrated that magnesite nucleation in “wet” saturated scCO2 was dependent on the thickness of the adsorbed water; increasing water concentration above a certain threshold (4−8 nm layer thickness) led to a continuation in magnesite growth.37 Formation of magnesite in a thin water film introduced the possibility that MgCO3 could be precipitated at low temperature in aqueous solution as well, if key conditions in those films such as local pH, pCO2, saturation state, and electrolyte composition could be determined. An unfavorable aspect in magnesite formation by carbonation of forsterite is the slow kinetics, dominated by the dissolution rate of the silicate mineral, which could result in a lower degree of magnesite saturation, deemed to be critical for MgCO3 nucleation and growth.2 To overcome this solubility constraint, many researchers have explored brucite, Mg(OH)2, as a Mg source to investigate Mg carbonates growth.38−41 Brucite is a common component of ultramafic mine tailings with a significant potential for CO2 sequestration due to high reactivity when exposed to CO2 and water.42 Although present only as a minor component, it was estimated that the amount of CO2 captured via brucite carbonation could significantly offset the annual greenhouse gases emissions in active mining operations.42 Here, we studied magnesite formation at high pCO2 to determine the mechanistic processes leading to its formation at conditions relevant to carbon sequestration. The objectives were to investigate the temporal evolution of magnesite and characterize the mineralogy of Mg-carbonate products at different stages of reaction using a variety of techniques, including scanning electron microscopy and energy dispersive spectroscopy (SEM/EDS), X-ray diffraction (XRD), and

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MATERIALS AND METHODS Ex Situ Experiments. Batch precipitation experiments were conducted in stainless steel vessels (model 4790, Parr Instrument Company) with a net volume of approximately 33 cm3. Each reaction vessel was loaded with 0.58 or 0.87 g of synthetic Mg(OH)2, respectively, for the 50 and 35 °C experiments. The loading ensured high supersaturation with respect to magnesite upon complete dissolution of brucite in 10 mL of added DI water. Brucite was purchased from SigmaAldrich with a purity of 95%. Examination of the synthetic Mg(OH)2 by SEM/EDS analysis identified Ca as a minor component (∼1−2%). Reaction vessels were filled with CO2, ∼7.3 g for 50 °C and ∼16.2 g or ∼22 g for 35 °C, that upon equilibration resulted in 90 or 110 atm partial CO2 pressure, as indicated by the pressure gauge (the computed CO2 mass is shown in the Supporting Information). The stainless steel vessels filled with brucite/CO2 mixtures were transferred to static incubator ovens set at 35 and 50 °C. Following equilibration to temperature, 1−2 h, the reactor vessels were continuously monitored to ensure no loss of pressure during the course of the experiment. The reaction was terminated at different time periods (Table S1), by quickly depressurizing the vessel and rapidly filtering the suspension through a gentle vacuum filtration unit. Brucite batch studies at 50 °C were conducted in deionized water at a pCO2 of 90 atm from 1 to 56 days (Table S1). Two values of CO2 partial pressure were selected for studies at 35 °C: 90 and 110 atm with reaction time from 63 to 135 days. Solid Characterization Methods. Reaction products from 35 and 50 °C studies were dried immediately at the reaction temperatures and were analyzed by XRD to identify the solids’ mineralogy, by SEM for morphology and particle size characterization, and by TGA-MS for determining the type of carbonation solids and the extent of carbonation reactions. Powdered samples were mounted on zero background quartz slides for bulk XRD analysis that was carried out on a Philips X’Pert MPD system with vertical goniometer. Data were collected for 45 min over a 2θ range of 0 to 75° with a scan step of 0.04°, acceleration voltage of 50 kV, and beam current of 40 mA. The X-ray source was a long-fine-focus ceramic X-ray tube with a Cu anode. Semiquantitative compositions were determined by whole-pattern (Rietveld) fitting of the XRD patterns using TOPAS v4.2 (Bruker AXS). The SEM investigation was performed using a Field Emission Focused Ion Beam-Scanning Electron Microscope, FEI Helios 600, equipped with an energy-dispersive X-ray detector for qualitative analysis at acceleration voltage of 20 keV. The TGA-MS analysis was performed in a TG-209F1 (Netzsch Instruments) by heating the samples to 900 °C with a heating ramp of 2 °C/min under N2 flow (20 mL/min). Mass spectrometric analysis was carried out simultaneously on an Aeolos QMS-403C. Samples were analyzed while monitoring ion currents corresponding to H2O (m/z 18) and CO2 (m/z 44). 10737

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Environmental Science & Technology In Situ Experiments. In situ batch studies at 90 atm pCO2 and 50 °C, with equal amounts of brucite and fluid as in ex situ studies, were conducted using similar stainless steel reactors fitted with ports for inserting a sampling line and a pH electrode. The sample line consisted of a high-pressure stainless steel loop with a volume of 200 μL connected to the reactor. A ceramic porous filter (1-μm pore size) was mounted in front of the sample line to prevent large particles from passing through. At times, the sampling line was open for pressurized solution to fill the closed sample loop. The liquid was released to a 1-mL syringe, was quickly passed through a 0.22-μm Millipore filter, and the filtrate was collected in a tube holding a known volume of 2% HNO3. Elemental analysis on the acidified solution was performed with an ICP-OES (PerkinElmer Optima 2100). Between sampling events, the line was flushed with acid and deionized water. Vessel pressure decreased ∼5% after 10 sampling events. The in situ pH, in gently stirred and not stirred suspensions, was measured with a high-pressure pH electrode (Supporting Information) using a calibration curve that was generated at ambient conditions with standards equilibrated at 50 °C. There was no difference in pH measurement due to the effect of stirring. Thermodynamic Modeling. Thermodynamic modeling calculations were performed using the chemical equilibrium model GMIN which utilizes the ion-interaction model of Pitzer and co-workers.43−45 The use of the ion-interaction model of Pitzer was necessary to accurately describe the chemical equilibria used in this study (Table S1). The thermodynamic data used in the model are based on the 25 °C database of ioninteraction parameters and standard chemical potentials, which has been shown to accurately predict mineral solubilities in the HCO3−CO3−CO2−H2O system to high concentration.46 Because the temperature ranged between 35 and 50 °C, temperature corrections were made to the model while maintaining consistency with the 25 °C database.46 These temperature corrections are described below. Parameters for Mg2+−, −CO32−, and −HCO3− were corrected for temperature using the first derivative with respect to temperature expressions.44 Parameters for Mg2+−HCO3− ion-interactions were taken from a recent critical evaluation.47 Interactions between the divalent ions Mg2+−CO32− were described by the use of MgCO3(aq) ion pairs. Equilibrium constants for these species and equilibrium constants as a function of temperature for the formation of HCO3− and CO3− were taken from literature data.48 Thermodynamic data for solid phase brucite (Mg(OH)2) were corrected for variations in temperature using the van’t Hoff equation using the data for the enthalpy of reaction.48 Thermodynamic data for magnesium carbonate phases, magnesite (MgCO3 ), 47 nesquehonite (MgCO 3 · 3H2O),47 and hydromagnesite (Mg5(CO3)4(OH)2·4H2O),49 were taken from the literature.

Figure 1. XRD patterns representing reacted solids at different equilibration periods: (a) 5, (b) 14, (c) 28, and (d) 56 days. Formation of Mg-carbonates (N, nesquehonite; HM, hydromagnesite; M, magnesite) occurred at various stages during equilibration of B, brucite, with CO2 at 50 °C and 90 atm. At 56 days reaction period magnesite was the dominant solid. Low angle reflection observed at 5 days was not positively identified, however, this minor component possibly was a hydrated carbonate solid formed early in the reaction.

rod-shaped particles reaching a length of 30 μm (Figure 2a, Figure S1 A, B). That formation of Mg-carbonates occurred as rapidly as 1 day after reaction is in agreement with previous studies where formation of metastable carbonates simultaneously upon dissolution of brucite was observed in carbonaterich solutions at slightly acidic pH.38 Formation of both nesquehonite and hydromagnesite, however, is intriguing since it is believed that conditions leading to nesquehonite precipitation are different from those of hydromagnesite.50 For example, it has been shown that increasing alkalinity while maintaining a constant activity of carbonate ion would favor formation and stability of hydromagnesite.49,51 Meanwhile, nesquehonite has been confirmed to initially form during brucite equilibration with atmospheric CO2, at conditions that would thermodynamically favor precipitation of hydromagnesite.52 Formation of nesquehonite at these conditions was attributed to the fact that metastable products with simpler structures tend to form more rapidly than the compounds with complex ones.53 Characterization of the 7 day reaction products with SEM in combination with a XRD semiquantitative analysis indicated that hydromagnesite and nesquehonite remained the major carbonate solids. (Figure 1, Table S2). Magnesite formation at this reaction time was not readily evident from bulk XRD analysis, however, submicron size crystallites with the rhombic habit typical of magnesite were observed by SEM (Figure 2b, Figure S1 C, D). Not surprisingly, thermodynamic equilibrium calculations predicted that if the added mass of brucite was equilibrated with CO2 at 50 °C and 90 atm the system would be supersaturated with respect to magnesite (Table S1). Increasing reaction time to 10 and 14 days (Figure 2b, Figure S1 D) caused a slight increase in the size of rhombohedral crystallites, as well as in their abundance. In addition, brucite was reduced to a minor component by 14 days and

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RESULTS AND DISCUSSION Magnesite Formation in Aqueous Solutions at 50 °C and 90 atm pCO2. Formation of the hydrated carbonates, nesquehonite and hydromagnesite, at 1−5 days of reaction was confirmed through the appearance of diagnostic reflections in the XRD patterns for both carbonate phases. Reflections not assignable to carbonates were those corresponding to a portion of the original unreacted brucite (Figure 1a). At this stage of reaction, carbonate precipitates were characterized by two distinct morphologies: hydromagnesite occurred as particles with thin platelet morphology and nesquehonite occurred as 10738

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Figure 2. SEM micrographs of products formed during brucite/water/CO2 experiments at 50 °C and 90 atm depicting particles with different morphologies: platelets characteristic of hydromagnesite, rod-like particles characteristic of nesquehonite, and rhombohedral crystals characteristic of magnesite. By 56 days magnesite appears as the sole carbonation product.

three-step mass loss from decomposition of H2O (18 m/z), indicating the presence of a hydrated solid in the reaction product (Figure 3a, red line). The most significant mass loss, however, was due to the CO2 (44 m/z) release (∼40%, Figure 3a, blue line), which reached a maximum at 430 °C; a temperature within the decomposition of nesquehonite (390 °C) and hydromagnesite (474 °C).54−56 Generally, the total mass loss assigned to H2O and CO2 was ∼59%, which, combined with the XRD analysis, indicated that the reacted solids were composed of a mixture of hydromagnesite, nesquehonite, and brucite. A typical TGA-MS pattern of carbonation products at 7 to 14 days of reaction shown in Figure 3b resembled the previous pattern with the exception of a small CO2 release appearing as a low intensity shoulder at 490 °C. Although minor, this shoulder could be interpreted as a secondary release of CO2 at higher temperature possibly indicating formation of a new carbonate solid. At 28 days, the TGA-MS data showed a bimodal release of CO2 covering a temperature range between 300 and 550 °C (Figure 3c). The initial release of CO2, peaking at 390 °C, is consistent with decomposition of hydrated carbonates. The second CO2 release, which reached a maximum at 510 °C, is indicative of a carbonate solid with a higher degree of crystallinity similar to magnesite. Lastly, at 56 days of equilibration, Figure 3d, the reaction product started to decompose at ∼450 °C and peaked at 540 °C based on mass 44 (m/z); the assignable mass loss was 50.12%, consistent with magnesite’s theoretical mass loss of 52%. The H2O mass loss was ∼5% and the presence of a negligible CO2 release at 380 °C indicated only minor

nesquehonite particles appeared to undergo some degree of morphologic transformations. Consequently, at 14 days of reaction, locating nesquehonite particles with SEM was challenging although particles with the hydromagnesite morphology seemed widespread. A transition of nesquehonite to hydromagnesite has been reported in separate studies; however, it was not clear in the present study if nesquehonite transformed into hydromagnesite.23,50,52 In general, XRD data between 7 and 14 days of reaction confirmed the presence of hydromagnesite and nesquehonite, while magnesite only weakly fit the diffraction pattern (Figure 1b). The most striking changes in the solid products were observed at 28 and 56 days of reaction. Semiquantitative XRD analysis (Table S2) showed that, by 28 days, hydromagnesite and magnesite were the only carbonates to contribute to the crystallographic pattern (Figure 1c). Reflections typically assigned to nesquehonite and brucite were noticeable absent. Consistent with the XRD results, microphotographs at 28 days revealed well-formed, easily identifiable magnesite particles (Figure 2c, Figure S1 E, F). Lastly, after 56 days of equilibration at 50 °C and 90 atm magnesite was identified by XRD as the sole carbonate solid, with individual crystals that grew in size between 2 and 3 μm (Figure 1d, Figure 2d, Figure S1 G, H). Further insights into formation of intermediate carbonates were obtained through TGA-MS analysis. Decomposition patterns collected on reacted solids provided in Figure 3 represent the most significant stages in the brucite carbonation reactions at high CO2 pressure. In detail, TGA-MS data from solids collected between 1 and 5 days of equilibration revealed a 10739

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Figure 3. TGA decomposition patterns represent solids reacted at 50 °C and 90 atm CO2 for (a) 5, (b) 14, (c) 28, and (d) 56 days. Eighteen m/z H2O (red line) and 44 m/z CO2 (blue line) mass spectra analysis revealed formation of hydrated carbonates early in reaction (a). At the end of reaction, formation of a carbonate solid with higher degree of crystallinity and mainly anhydrous was observed, indicative of magnesite. Data in (b) and (c) were composed by a mixture of solids identified in (a) and (d).

equilibration, the results for both pCO2 values of 90 and 110 atm showed that although the metastable phase, nesquehonite, appeared to be the dominant reaction product, formation of distinct rhombohedral crystallites of the same morphology as magnesite was evident in SEM micrographs (Figure S2). Hydromagnesite appeared to form only as a minor component (Table S2), consistent with the expectation that its precipitation is favored at temperatures above 40 °C.23 In fact, direct precipitation of hydromagnesite from aqueous solutions has been reported to be kinetically inhibited at ambient conditions.23,57 At the longer equilibration period of 135 days, magnesite was identified as the primary solid carbonate only at higher pCO2 (110 atm), while at 90 atm nesquehonite remained as the main reaction product. Notably, the significant effect that partial pressure had on magnesite formation and growth is illustrated in Figure 4. The data clearly indicate that given long equilibration times (∼140 days) and high partial CO2 pressure, formation of magnesite would take place in an aqueous environment even at a temperature as low as 35 °C. Further acceleration of magnesite formation could be achieved by increasing pCO 2 well above supercritical conditions. Linking Magnesite Formation at High pCO2 to Aqueous Chemistry. Undoubtedly the most significant finding in this study was the ultimate formation of anhydrous

occurrences of hydrated carbonates at this reaction time. In general, solid characterization techniques showed that early formation of magnesite ∼7 days was followed by a gradual transformation to a sole carbonation solid at long equilibration time, while hydrated magnesium carbonates appeared to have almost entirely disappeared. The importance of hydrated metastable phases on the formation of magnesite in the early stages of reaction could be significant, because it has been proposed that magnesite could be facilitated by the presence of a crystalline matrix.35 In this study, individual submicron magnesite crystallites were located spatially close to the hydrated carbonates (Figure 2b, Figure S1 C, D). However, magnesite particles appeared structurally separated from hydrated solids, implying that magnesite grew by precipitation at the expense of metastable intermediates. Microscopic evidence, particularly, seems to indicate that magnesite nucleation and growth may have coincided with dissolution of metastable hydrated carbonates. Magnesite Formation in Aqueous Solutions at 35 °C. Formation of magnesite in aqueous solutions at lower temperature, 35 °C, was investigated for long equilibration periods, 63 and 135 days, and at pCO2 of 90 and 110 atm. The earliest equilibration time, 63 days, was selected to closely match the time where full conversion of magnesite was observed in the 50 °C experiments. At 35 °C and 63 days of 10740

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Figure 4. XRD patterns and SEM micrographs of solid products from 35 °C reaction of brucite suspension with scCO2 (90 and 110 atm). Nesquehonite appeared as the major reaction product during the course of the study except for the experiment conducted at pCO2 of 110 atm at longer time scale. In contrast, magnesite which appeared as a minor component at earlier time and/or lower pressure grew as a primary carbonate solid at 35 °C, 135 days, and 110 atm.

MgCO3 in aqueous solutions at 35 and 50 °C, and, to the best of our knowledge, this is the first study that reports magnesite synthesis in laboratory at near ambient temperature. The data reveal a strong dependency between MgCO3 formation in aqueous solution and high pCO2 environment. We may interpret the collective observations as follows. Initially, at high pCO2, dissolution of brucite, dissolution and protonation of CO2 in solution, and formation of Mgcarbonates took place simultaneously (Figure 5). Results from in situ experiments showed that brucite dissolution occurred immediately upon pressurization of a Mg(OH)2 suspension with scCO2 (Figure 5c, Table S3). Following pressurization, slightly acidic conditions were established rapidly as demonstrated by the comparison of the bulk solution pH values that were monitored for a period of 14 days with ex situ pH calculations at equilibrium (Table S1). The pH values, however, gradually increased with time, ∼0.9 units (Figure 5b). We should note that the in situ pH measurements at high pressure were problematic, because the pH electrode could only be calibrated once, at the beginning of the experiment. The longer-term pH data, therefore, could not be corrected for probable drifts with time (details in Supporting Information). Ex situ solid characterization data revealed that carbonation also occurred as early as 1 day of equilibration. In fact, based on XRD semiquantitative analysis (Table S2), reaction products consisted of hydrated Mg-carbonates (nesquehonite, hydromagnesite) and a fraction of unreacted original brucite. Despite the fact that initially carbonation reactions seemed to follow a stable trend with time (Figure 5a, see 1 to 5 days reaction data), the correspondingly linear increase in aqueous Mg2+ concentration was indicative of an active dissolution process (Figure 5c). Yet, semiquantitative XRD analysis (Table S2) during this time period revealed that the unreacted proportion of the original brucite appeared to remain largely

unchanged. That a fraction of the Mg(OH)2 stayed unreacted for a period of time was possibly due to brucite surface passivation by the extensive formation of hydrated carbonates.58 However, to account for the significant increase in aqueous Mg2+ concentration it can be inferred that the newly formed hydrated Mg carbonates were likely dissolving. In fact, the early observation that both metastable carbonates, nesquehonite and hydromagnesite, precipitated simultaneously was suggestive that their formation occurred in localized regions characterized by spatial/temporal fluctuation in pH and in Mg concentrations. These localized regions, we hypothesize, were far from equilibrium with respect to bulk aqueous phase. The metastable phases formed therein were thus likely to dissolve over time as the relevance of localized domains gave way to contact with bulk solution, which was evidently undersaturated with respect to intermediate carbonates, nesquehonite and hydromagnesite (Table S1). The formation of magnesite at >7 days of reaction correlated directly with the highest in situ Mg2+ concentration, 0.56 mol/L (Figure 5). Given that magnesite saturation was reached after 1/2 h of reaction (saturation index, SI = 0.28, Mg2+ = 0.04 mol/ L, Table S1), undoubtedly the solution was highly oversaturated with respect to magnesite at 7 days equilibration (SI = 2.36, Table S1). The SIs were calculated assuming equilibrium with the scCO2 phase. The calculation results revealed that the solutions became rapidly oversaturated with magnesite and approached or achieved equilibrium with nesquehonite and/or hydromagnesite (Table S1). The measured pH values were significantly higher than model predictions even at low aqMg which indicates a nonequilibrium behavior especially at early times. Immediately following magnesite nucleation, the aqueous data showed a gradual Mg2+ decrease with time that coincided with a consistent increase of %CO2 (in reacted solid), the latter 10741

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observed to occur at the 14−28 day interval, followed by the loss of hydromagnesite at 28−56 days range (Figure 5a). Possibly, equilibration with the slightly acidic bulk environment heightened the dissolution of these intermediate precipitates. The observation that complete carbonation occurred only at long time period, even though magnesite nucleation happened relatively early, also seemed to support the statement that magnesite formation and growth occurred gradually following metastable carbonates dissolution. The formation of magnesite in the aqueous systems at highly elevated CO2 partial pressures, therefore, seemed to take place via dissolution−recrystallization and did not appear to occur via solid phase transformation.34,59 The extensive rapid formation of hydrated carbonates upon brucite contact with scCO2 (Table S1 1day) probably hindered direct and immediate nucleation of magnesite. Furthermore, the early formation of hydromagnesite and/or nesquehonite, despite the fact that bulk solution was undersaturated with both solids (Table S1 in situ data), could be indicative that mineral− fluid interfacial processes controlled the hydrous carbonates precipitation. In fact, the formation/transformation of Mg carbonates during transitional and final stages of reaction at both 35 and 50 °C seemed to follow the basic concepts of interface-coupled dissolution−precipitation mechanism.52 Actually, it is highly plausible that similar interfacial mechanistic pathways described in detail by several papers52,60,61 were operating in the present study. From an earlier hypothesis, fast brucite dissolution resulted in localized interfacial regions that were saturated with respect to hydromagnesite and/or nesquehonite and led to their precipitation. Besides removing Mg ions from the interfacial solution, precipitation of hydrated carbonates slowed aqueous Mg release by increasing the surface passivation of the remaining brucite. Coupled with continuous Mg diffusion toward bulk solution, these processes led to an interfacial fluid undersaturated with respect to hydrated carbonates. Consequently, the earlier rapidly formed hydrous Mg-carbonates started to dissolve resulting in a fluid supersaturated with the stable magnesite which began to precipitate. Formation of small magnesite crystals that noticeably grew in quantity with time (although a slight advancement in crystallite size was observed as it was described earlier) could possibly indicate that numerous MgCO3 nucleation centers were provided by the dissolving solids. Considering that magnesite crystals grew distinctly and were not attached to the remaining solid indicates that the structure of metastable phases was not used as a growth template (Figure 2). Lastly, the dynamics of magnesite nucleation and growth were influenced by further increasing the partial pressure of CO2 as was shown in Figure 4 for 35 °C and 110 atm. One explanation we hypothesize is that increasing pCO2 promotes faster dissolution of metastable intermediates and therefore enhances magnesite formation. Environmental Implication. This study identifies a complex series of geochemical processes controlling magnesite precipitation from CO2 rich fluids at temperature and pressure conditions relevant to geologic sequestration. For example, the CO2 fluid temperatures at the Wallula Basalt Pilot Project62−64 varied with depth from 44 °C at the top of injection zone ∼850 m to 36 °C at formation−water interface. CO2 pressures at this injection site were controlled so as not to deviate higher than 10 atm above the reservoir conditions of ∼82 atm. Therefore, precipitation of magnesite (at 35 and 50 °C; 90 and 110 atm pCO2) in the present study highlights the importance of needing a better understanding of the carbonation pathways

Figure 5. Brucite transformation in scCO2, 50 °C and 90 atm: (a) ex situ data demonstrating formation of Mg-carbonates with reaction time; (b) in situ pH indicating slightly acidic conditions following pressurization; and (c) in situ data illustrating stages of reactivity: (i) solid dissolution determined from the increase in aqMg2+, (ii) magnesite formation determined from the gradual decline in aqMg2+, and (iii) a steady-state growth indicated by stable aqMg2+ concentration.

calculated from TGA data (Figure 5). Theoretically, if during the brucite reaction with scCO2, magnesite were the primary carbonate solid, then the CO2 content in the reaction product would increase up to 52%. However, a product with composition of 32 or 38% CO2 would be indicative of nesquehonite or hydromagnesite, respectively. The data in Figure 5a demonstrate that a combination of all phases was present at 7 days of reaction (CO2 (in reacted solid) = 41%), however, with time, magnesite grew as a major product of reaction at the expense of metastable intermediates (Figure 2). In fact, loss of the least stable phase at 50 °C, nesquehonite, was 10742

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Article

Environmental Science & Technology

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and processes expected to occur during subsurface storage of CO2. Although magnesite formation has been extensively examined, the present study is the first to report laboratory precipitation of magnesite at near ambient temperature from aqueous conditions. The inhibition of magnesite formation at low temperature and aqueous solutions has often been linked to the highly hydrated nature of Mg2+ ion and its unfavorable dehydration energetics.17,19 In this study we have demonstrated through a series of experiments that kinetically inhibited magnesite forms through the precipitation and dissolution of metastable precursor hydrated solids at conditions applicable to CO2 capture. Finally, even though formation of magnesite is slow, our findings support the need for inclusion of magnesite as an equilibrium phase in reactive transport simulations of CO2 sequestration in subsurface formations, given the long time scale required for effective disposal of greenhouse gases.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.5b02588. Additional information as noted in the text (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]; tel: +1-509- 371-6383; fax: +1-509-371-6354. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Dr. J. Hövelmann and the anonymous reviewers for their valuable suggestions and comments that improved the quality of the article. This work was supported by the Geosciences Research Program at PNNL supported by the U.S. Department of Energy, Office of Basic Energy Sciences, Division of Chemical Sciences, Geosciences & Biosciences, and the Office of Fossil Energy. Several of the experiments were performed using the Environmental Molecular Sciences Laboratory, a national scientific user facility sponsored by the U.S. Department of Energy’s (DOE) Office of Biological and Environmental Research, and located at PNNL. PNNL is operated for DOE by Battelle Memorial Institute under Contract DE-AC06-76RLO-1830. D.A.D. thanks the Robert Ramsay Fund of The University of Alabama for partial support.

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REFERENCES

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DOI: 10.1021/acs.est.5b02588 Environ. Sci. Technol. 2015, 49, 10736−10744