Environ. Sci. Technol. 2010, 44, 5999–6005
Dissolution and Precipitation of Clay Minerals under Geologic CO2 Sequestration Conditions: CO2-Brine-Phlogopite Interactions HONGBO SHAO, JESSICA R. RAY, AND YOUNG-SHIN JUN* Department of Energy, Environmental and Chemical Engineering, Washington University in St. Louis, St. Louis, Missouri 63130
Received March 31, 2010. Revised manuscript received June 2, 2010. Accepted June 4, 2010.
To ensure efficiency and sustainability of geologic CO2 sequestration (GCS), a better understanding of the geochemical reactions at CO2-water-rock interfaces is needed. In this work,bothfluid/solidchemistryanalysisandinterfacialtopographic studies were conducted to investigate the dissolution/ precipitation on phlogopite (KMg3Si3AlO10(F,OH)2) surfaces under GCS conditions (368 K, 102 atm) in 1 M NaCl. Phlogopite served as a model for clay minerals in potential GCS sites. During the reaction, dissolution of phlogopite was the predominant process. Although the bulk solution was not supersaturated with respect to potential secondary mineral phases, interestingly, nanoscale precipitates formed. Atomic force microcopy (AFM) was utilized to record the evolution of the size, shape, and location of the nanoparticles. Nanoparticles first appeared on the edges of dissolution pits and then relocated to other areas as particles aggregated. Amorphous silica and kaolinite were identified as the secondary mineral phases, and qualitative and quantitative analysis of morphological changes due to phlogopite dissolution and secondary mineral precipitation are presented. The results provide new information on the evolution of morphological changes at CO2-water-clay mineral interfaces and offer implications for understanding alterations in porosity, permeability, and wettability of pre-existing rocks in GCS sites.
Introduction Mitigation of climate change requires immediate actions to reduce anthropogenic CO2 emission and avoid the risks of future devastating effects. Capturing CO2 from large stationary sources such as power plants and storing it in deep saline aquifers is considered one of the more viable means to achieve this urgent goal (1, 2). In this process, CO2 is injected into formation rocks (usually sandstone) with high porosity and permeability, and above the formation rock, thick caprock layers with low permeability serve as a seal for the stored CO2. Injecting massive amounts of CO2 into deep geologic formations may cause a range of coupled thermal, hydrodynamic, mechanical, and chemical changes (3, 4). One change could be the alteration of the geochemical equilibrium between rock-forming minerals and formation waters. The interruption of the geochemical equilibrium could cause the * Corresponding author phone: (314)935-4539; fax: (314)935-7211; e-mail:
[email protected]; http://encl.engineering.wustl.edu/. 10.1021/es1010169
2010 American Chemical Society
Published on Web 06/29/2010
dissolution of rock minerals and/or precipitation of secondary mineral phases and thus change the morphology, porosity, permeability, and wettability of pre-existing rocks. These changes will then affect the transport of CO2 and the integrity of caprocks and formation rocks within the reservoirs. Therefore, understanding the kinetics and mechanisms of the interactions between minerals and fluids (formation water and supercritical CO2) is critical to achieve better prediction of the short- and long-term fate and transport of CO2 in GCS. In recent years, through microscopic and spectroscopic analysis on the solid and liquid phases, more understanding has been achieved concerning the alteration of rocks induced by CO2 injection (3) (and references therein). For example, the precipitation of secondary mineral phases on pre-existing rocks has been studied under GCS conditions (5-12). Most recently, Gilfillan et al. (13) reported that in seven gas fields with siliciclastic or carbonated reservoir lithologies, CO2 dissolution in formation water is the dominant sink for CO2. They also reported that a maximum of 18% of sequestered CO2 can form carbonate precipitation, which leads to their conclusion that “long-term anthropogenic CO2 storage models in similar geologic systems should focus on the potential mobility of CO2 dissolved in water”. However, one important point is missing from this report is that even small quantities of secondary mineral phases can cause significant changes in porosity and permeability (3, 4), depending on the size, shape, phase, and location of the precipitates. For example, if the precipitation occurs near the opening of pores in the formation rocks and the size is similar to the pore throats, then the permeability will decrease significantly. However, so far, no study has reported on the size, shape, phase, and location of secondary mineral phases and their transport under GCS conditions. Acquiring wellcontrolled experimental data on this aspect would help to construct effective prescreening criteria to select GCS sites for more sustainable CO2 sequestration projects. Clay minerals can be components of both formation rock and caprock in potential GCS sites (3, 14). Mica minerals, as an important component of clay, were found in the caprock of the Sleipner site in the North Sea (15) and in the Mount Simon sandstone at a potential CO2 sequestration site in the Illinois basin (16). Clay minerals that were originally present in sandstone showed ion-exchange reactions with saline fluids under GCS conditions, as reported by CzernichowskiLauriol et al (17). However, a holistic understanding of interactions of clay minerals and CO2, especially the surface morphology changes, is still pending. The objectives of this work were to investigate both the interactions between clay minerals and CO2 saturated saline water under GCS conditions and the potential influence of these interactions on the operation of GCS projects. Specifically, we aimed to understand the dissolution kinetics and surface morphology changes of a model clay mineral, phlogopite, in the presence of CO2 under high pressure (102 atm) and temperature (95 °C) in 1 M NaCl solution. To reach this goal, microscopic and spectroscopic measurements and fluid chemistry analyses were combined to qualitatively and quantitatively investigate changes in size, shape, and location of secondary mineral phases as well as those of dissolution pits.
Experimental Section Chemicals and Minerals. All chemicals used in this study were at least ACS grade. Deionized water was used after passing it through two 0.2 µM filters. Phlogopite (general formula: KMg3(Si3Al)O10(F,OH)2) was obtained from Ward’s VOL. 44, NO. 15, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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Natural Science, NY. The chemical composition of the phlogopite sample, analyzed by electron microprobe (EMP) and X-ray fluorescence (XRF), showed that the concentrations of K, Mg, Al, and Si was not significantly different in the bulk (XRF result) and µm range around surface (EMP result). The atomic ratio of K:Mg:Si:Al on the surface was 1:2.81:3.02: 1.24. Phlogopite specimens were cleaved to thin samples (∼20 µm, {001} cleavage surface), cut into rectangular flakes with dimensions of 2.5 cm × 0.8 cm, and cleaned with water, acetone, ethanol, and isopropyl alcohol to remove organic matter. High P/T Reaction System and Phlogopite Dissolution Experiment. In this study, experiments were conducted in a benchtop reactor constructed from HC alloy-276 (Parr Instrument, Moline, IL). In situ pH was measured with high pressure and temperature pH probes (Corr Instrument, TX). PTFE tubes containing phlogopite flakes and 4 mL 1 M NaCl solution were placed in the reactor. CO2 was pressurized to 102 atm in a syringe pump and introduced into the reactor, which was controlled at 368 K. The temperature, pressure, and NaCl concentration were chosen to simulate a GCS system. After the reaction (elapsed times were 3, 5, 8, 22, 43, 70, and 159 h), the aqueous phase was analyzed with an inductively coupled plasma-mass spectrometer (ICP-MS) (7500ce, Agilent Technologies, CA). The surface features of phlogopite were analyzed with AFM. AFM Sample Analysis. Phlogopite samples after reaction were washed with doubly filtered water and dried by flushing with N2. All height, amplitude, and phase contrast images of each phlogopite sample were collected simultaneously under ambient laboratory conditions in tapping mode. Probes were 125 µm long, with phosphorus (n) doped silicon tips (a nominal tip radius of 10 nm, MPP-11100-10, Veecoprobes). The images were collected with drive frequencies of 312-320 kHz, typical spring constants of 20-80 N/m, and scan rates of around 1-1.2 Hz. Identification of Secondary Mineral Phases. To identify secondary mineral phases, multiple techniques were utilized, including scanning electron microscopy equipped with energy-dispersive X-ray spectroscopy (SEM-EDX), X-ray diffraction (XRD), X-ray photoelectron spectroscopy (XPS), and nanosecondary ion mass spectroscopy equipped with Auger electron spectroscopy (Nano-SIMS-AES). However, these techniques failed for phlogopite flake samples for several reasons: the higher penetration depth of SEM-EDX compared to the small size of surface precipitates, the high concentration and crystallinity of secondary mineral phases required for XRD, and difficulties overcoming the high charging issue for Nano-SIMS analysis despite use of Au coatings. To facilitate the formation of secondary mineral phases, 0.4 g of phlogopite powder (passed through a 90 µm sieve) was added in 60 mL of 1 M NaCl solution and reacted with CO2 in the reactor equipped with pH probes for 159 h. After reaction, the suspension (including both liquid and solid phases) was centrifuged, and the solid phase was washed with water and then oven-dried at 50 °C. These samples were analyzed with a diffuse reflectance infrared Fourier transform spectroscope (DRIFTS -Thermo Scientific, Nicolet Nexus 470 instrument equipped with Praying Mantis) and XRD (Rigaku Geigerflex D-MAS/A, Japan). The supernatant was analyzed with high-resolution transmission electron microscopy (HRTEM, JEOL JEM-2100F field emission).
Results and Discussion Dissolution of Phlogopite under CO2 Sequestration Conditions. The in situ pH of the 1 M NaCl-phlogopite system at 368 K dropped from around 5.68 ( 0.04 to 3.08 ( 0.05 within 1 h after CO2 was introduced, and it remained in that range during the following 158 h. In the aqueous phase, providing congruent dissolution occurs, the ratio of K:Mg:Si:Al should 6000
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FIGURE 1. Phlogopite dissolution in 1 M NaCl solution under 102 atm and 368 K. Error bars are the standard deviation of the means of three measurements for triplicate samples. The inset is an enlarged figure for Si, Mg, and Al. be 1:2.81:3.02:1.24, as we measured for phlogopite. In this study, however, we observed incongruent dissolution of phlogopite, with enhanced release of K over Mg, Si, and Al (Figure 1). At ambient pressure and temperature, the interlayer K in mica is dissolved through ion-exchange with other cations in solution (18). Therefore, the significantly high K concentration in the NaCl solution indicates that potassium dissolution from phlogopite under GCS conditions also occurs through ion exchange. The dissolution of Mg, Al, and Si from the phlogopite framework, however, was congruent. The calculated average ratio of Mg:Si in 1 M NaCl solution at 8, 22, 43, 70, and 159 h was 0.80 ( 0.26 (The standard deviation was calculated according to the propagation of uncertainties (19).), and the ratio of Al:Si was 0.38 ( 0.09. These ratios are not significantly different from those in the phlogopite framework (0.93 and 0.41, based on EMP results for unreacted phlogopite), suggesting that Al, Si, and Mg were proportionally dissolved from the framework of the phlogopite surface. A thermodynamic calculation with Geochemist’s Workbench (GWB, Release 8.0, RockWare, Inc.) was done to determine the presence of secondary mineral phases in the aqueous phase at the longest experimental reaction time (159 h). The results showed that the saturation indices for all the potential mineral phases provided by GWB were negative, indicating that within a 159 h period, the dissolution of the phlogopite was predominant and that no precipitation of new solid phases should be expected. To separate the effect of pressure from the pH effect of injected CO2, we conducted an experiment with N2 for 159 h under the same conditions as the CO2 experiments. The dissolution data under N2 indicated that except for K, the concentrations of different elements were less than 10% of those under CO2, suggesting that injecting CO2 into a deep saline aquifer will significantly interrupt the equilibrium between the saline water and rocks and cause rock dissolution. In addition, we observed that the presence of CO2 enhanced the ion-exchange of K by more than a factor of 2. While studying the replacement of interlayer K by Na+ in boiled 0.25 M NaCl, Newman found that there was a synergetic effect of protons on the cation exchange of K in phlogopite, which was explained by the incorporation of protons into the structure of phlogopite (20). This explanation can also be applied to our observation, because the dissolution of CO2 caused the pH in the reaction system to decrease to 3.08. Morphological Changes of Phlogopite Surfaces. After phlogopite reacted with CO2, both dissolution pits and nanoscale precipitates were observed on its surfaces. For all the samples at different reaction times, dissolution pits were
FIGURE 2. Formation of pits and channel-like features on phlogopite surfaces after reaction with CO2 in aqueous solution. Image A is shown in height mode (5 µm × 5 µm), after reaction in 1 N NaCl solution for 3 h under 102 atm of CO2 and 368 K. The white arrows show the position of dissolution pits; image B is shown in height mode with a scan size 10 µm × 10 µm, after reaction in 1 M NaCl for 159 h. The height cross sections under the images correspond to the white dotted line in the AFM images. The white column in image B shows the formation of a channel-like feature by pit connection.
TABLE 1. Analysis of Surface Features on Phlogopite Surfaces after Reaction in 1 M NaCl under 102 atm and 368 K at Different Reaction Times dissolution pits time (h)
dominant feature
3
Da
5
D
8
D and Pb
22
D and P
43
D and P
70 159
D and P D and P
precipitates (nanoparticles
depth (nm)
width (nm)
volume (nm3/ pits)
height (nm)
width (nm)
0.6–0.9 0.9–1.3 2.0–2.8 1.0–2.2
20–50 20–80 80–150 80–200
(7.92 ( 5.14) × 102 (4.79 ( 2.53) × 103
1.2–1.5
200–300
(3.56 ( 2.52) × 10
0.8–1.5 1.5–2.9 1.5–2.8 5–10 5–8
20–40 30–80 20–80 40–80 140–180
1.2–1.5 3.7–6.4 6.0–15.5 1.0–8.9
200–500 500–700 350–600 280–600
(1.19 ( 0.68) × 105
20–45
-
1.51 × 107
– -
4.82 × 106 6.48 × 107
(1.88 ( 1.07) × 104
(1.41 ( 0.76) × 10 -
4
6
8–62 5–140
volume (nm3/particle) (4.63 ( 2.06) × 102 (7.42 ( 3.78) × 102 (7.51 ( 6.16) × 103
a “D” represents “dissolution”. b “P” represents “precipitation”. The dashed lines represent data that are not available due to particle aggregation or pit connection. The “volume” values are the averaged volumes of 20-30 isolated dissolution pits or particles except for particles in the 43, 70, and 159 h samples, where the particle volumes were calculated for the largest particles measured. The uncertainties are the standard deviation of 20-30 measurements.
the predominant features. The height cross sections in Figure 2A (3 h),B (159 h) show that, initially, isolated shallow pits appeared on the surface. As the reaction proceeded, pits became deeper and bigger, which caused some pits to connect with each other and finally form channel-like structures (Figure 2B). Within 159 h, the dissolution pits were as deep as 15.5 nm (Table 1). This observation is consistent with our chemical analysis result, which indicated that the concentrations of Mg, Al, Si, and K constantly increased during 159 h (Figure 1). However, contrary to our expectation of no precipitation, we did observe the formation of nanoparticles on the phlogopite surface after only 5 h (Figure 3A). In AFM analysis, phase imaging can detect variations in composition, adhesion, friction, viscoelasticity, and other properties (21). The significantly different colors of the newly formed particles in the phase mode images (the insets in Figure 3), compared with the background (phlogopite), indicate that the newly formed nanoparticles were secondary mineral phase(s). In our control experiment that was conducted under N2 (pH was 5.65 ( 0.05) for 159 h, dissolution pits were the predominant feature, but the depth of the pits
was generally less than 1 nm. This result indicates that the introduction of CO2, not temperature or pressure, is the main reason for the significant changes of morphology on phlogopite surfaces. Several processes could be responsible for the unexpected formation of nanoparticles on the phlogopite surface in our reaction system. First, the presence of a concentration gradient between the near surface and bulk solution can induce surface precipitation. Chemical analysis by ICP-MS measured the concentrations of aqueous species only in the bulk solution; however, the local concentrations of dissolved species near the surface could be higher than that in the bulk solution for a less-disturbed reaction system (22, 23). In this research the reaction system was not stirred throughout the experiments, which rendered the concentration gradient plausible. Second, the phlogopite surface may serve as a substrate and result in heterogeneous nucleation, which requires lower free energy of formation of a new phase than homogeneous nucleation (24). Third, the consumption of protons during phlogopite dissolution will cause a local pH increase at the near surface water layer, while the bulk VOL. 44, NO. 15, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Evolution of nanoscale precipitation on phlogopite surfaces after reaction with CO2 in 1 M NaCl under 102 atm of CO2 and 368 K. Images A, B, C, E, and F are for reaction times of 5, 8, 22, 43, and 159 h, respectively. Image D is enlarged from a small area of image B. All images are shown in height mode. The insets on the lower left corners are in phase mode. The insets in the right upper corners are schematic illustrations of particle position. The height cross sections below the images correspond to the white dotted lines in the height images. The arrows indicate the position of the particles on the edge of a dissolution pit. The dotted, curved line in image D indicates the evolution of the dissolution pits. solution pH remains unchanged. This local pH change could facilitate the precipitation of secondary minerals. Therefore, on the phlogopite surface, the dissolved species could be 6002
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supersaturated with respect to the secondary mineral phases, while the bulk solution remains undersaturated. Nucleation and growth from undersaturated liquid solution in the
FIGURE 4. DRIFTS spectra of phlogopite powder before and after reacting with CO2 in 1 M NaCl solution under 368 K and 102 atm for 159 h. presence of a foreign substrate have been observed for both semiconductors (25) and ionic compounds (26). In aqueous solution, Murdaugh et al. reported the precipitation of SrSO4 and PbSO4 on a BaSO4 substrate with saturation ratios (ion activity product/solubility product (Ksp)) lower than 0.1 (26). This phenomenon was explained by the favorable ion-surface interactions leading to ion enrichment and actual supersaturation in the two-dimensional interfacial zone, which may also be the case in this study. Location and Mobility of Nanoparticles. The location and transport of the newly formed nanoparticles on the phlogopite surface during the reaction showed interesting characteristics. Nanoparticles of a secondary mineral phase first appeared after 5 h (Figure 3A) and were located near the edges of dissolution pits. The newly formed particles were ellipses: the horizontal size was generally 20-50 times bigger than the vertical size. Three hours later (reaction time 8 h), we observed that particles had grown in both abundance and size (Figure 3B), but they were still located near the edge of the dissolution pits, either inside or outside (Figure 3D). However, after 22 h, the majority of the nanoparticles were observed on the phlogopite surface away from the pit edges (Figure 3C), suggesting the particles had migrated from the pit edges and relocated on the surface. After 43 h, only a few particles remained on the edges. The aggregation of particles
began after 22 h and continued to increase as the reaction proceeded. At 159 h, the volume of aggregated particles was 4 orders of magnitude higher than that at 5 h (Table 1). The selectivity of particle location is related to the reactivity of the different surfaces in phlogopite structure. Rufe and Hochella reported that the dissolution rate of phlogopite on {hk0} edges is 2 orders of magnitude faster than the bulk rates for the mineral (27). Assuming the phlogopite surfaces were well cleaved, then initially the dissolution of the phlogopite starts from the basal {001} surface. The formation of pits on phlogopite surfaces makes more {hk0} surfaces available, which facilitates a rapid dissolution around the pit edges, and allows the supersaturated mineral phase to precipitate in the area that is near the edges of the dissolution pits. Another feature observed during the early period of particle formation was that the dissolution pits usually extended around the nanoparticles. This phenomenon indicates that the presence of the secondary mineral phase (less soluble than phlogopite) on the phlogopite surface inhibited the dissolution by masking the surface. As this process proceeded, the nanoparticles and the phlogopite pillars under the nanoparticles were isolated from the surrounding area and appeared as specks inside the dissolution pits, as observed in Figure 3B,D. Further dissolution of the phlogopite pillars under the nanoparticles made the pillars thinner than the nanoparticles, and eventually the nanoparticles and the thin phlogopite pillars left the surface, leaving small cavities on the phlogopite surfaces, as shown in Figure 3C. Quantitative analysis of 27 cavities in Figure 3C indicates (data not shown) that these cavities were much smaller than the nanoparticles at both 22 and 8 h, which supports the presence of the thin pillars. The particles could move to the aqueous phase, but the extent of this process may not be significant. When the phlogopite was washed with water after reaction, the particles still remained on the surface (based on AFM observations), suggesting that the secondary mineral phases had greater affinity for the phlogopite surface than for the solutions. Therefore, the particles were more likely to remain on the phlogopite surface in our reaction system, where no fluid movement was present. However, in GCS sites, especially during the CO2 injection period, the relocation of particles due to supercritical fluid transport could be significant.
FIGURE 5. Schematic diagram of proposed reaction pathways for dissolution and precipitation on a phlogopite surface under geologic CO2 sequestration conditions. The top view shows the evolution of the particle on the edges of dissolution pits. The drawing does not reflect real dimensions. VOL. 44, NO. 15, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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Identification of Secondary Mineral Phases. While other techniques failed to identify secondary mineral phases on phlogopite flakes, DRIFTS spectra showed significant changes for phlogopite powder after reaction (Figure 4). The DRIFTS peak areas at the ranges of 400-550 cm-1 and 900-1200 cm-1 (attributed to Si-O bands) increased significantly compared with the peaks at 815, 730, and 655 cm-1 (attributed to Al-O bands) (28). This result suggests that the Si-O bands are more abundant in the secondary mineral phases, while Al-O bands do not change much, compared with unreacted phlogopite. Additionally, in the AFM analysis, the newly formed nanoparticles on phlogopite flake surfaces did not exhibit a euhedral shape (Figure 3). Therefore, we concluded that the majority of the secondary mineral phase is amorphous silica. However, we can also consider the possibility of the coexistence of other minor products. The HR-TEM analysis of the supernatant from the powdered phlogopite sample showed that nanoscale kaolinite (Al2Si2O5(OH)4) particles were also present in the aqueous phase. Although we observed kaolinite nanoparticles in the supernatant, the major product was amorphous silica for several reasons: (1) characteristic peaks of kaolinite did not appear in the XRD patterns of the solid phase after reaction, suggesting that the quantity of kaolinite was very small; (2) the intensities of Al-O bands in the DRIFTS spectrum (Figure 4) did not increase significantly over the bands of unreacted phlogopite; and (3) in HR-TEM analysis, a significant number of particles with no diffraction pattern were observed, suggesting an amorphous material. One caveat, however, is that the products for powdered phlogopite might be different from those in AFM images obtained from phlogopite flakes sample. Environmental Implications. Overall, the early period reactions between phlogopite and CO2 saturated saline water occur in the following steps (Figure 5). First, surface elements dissolve into the aqueous solution, and dissolution pits appear at the surface. Then, as the dissolution proceeds and dissolved species near the surface become supersaturated with respect to secondary mineral phases, nanoscale precipitates form on the pit edges. When the growing nanoparticles reach a critical size, they leave the surface and create small cavities from the detachment of phlogopite pillars underneath the particles. After that, nanoparticles aggregate to form larger precipitates on the surface. The morphological and transport behaviors of these nanoparticles are important because they can alter the permeability of pre-existing rocks at the GCS, although the porosity does not change significantly due to the small size of the precipitates. The permeability change will depend on the size, shape, locations, and phase of the new solid. Studying eight different sandstone samples, Kate and Gokhale reported that 22-81% of the total pore volume had pore radii less than 100 nm, and 2-36% of the total pore volume had pore radii less than 10 nm (29). Although they did not provide information on the size of pore throats, it is reasonable to estimate that the diameter of the throats could be well below 200 nm. In this work, we found that, after 22 h reaction in 1 M NaCl solution, the precipitates on phlogopite surface were mostly elliptical, with lengths in the range of 40-180 nm (Table 1). For longer reaction times, the largest dimension of the particles can be more than 1 µm (Figure 3E,F). Therefore, if those nanoparticles grow in the pore throats throughout the clay layers, or if they leave the surface and are transported with the fluid along the porous formation rocks, they may significantly decrease the permeability and consequently hinder the transport of CO2. An experimental study on water-rock interactions during CO2 flooding in the formation rocks from the Tensleep saline aquifer, Wyoming, found that although dissolution definitely occurred for all eight core samples, the permeability of seven samples 6004
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decreased 16-44% (30). These field site observations can be explained more by detailed micro- and nanoscale studies. In addition to their size and location, the different morphology and phases of newly formed nanoparticles can alter the wettability of the pre-existing rocks. The secondary mineral phase could cover the surface of the original minerals and change the surface wettability of the porous media for CO2 transport. From the AFM images of reacted phlogopite, the surface coverage of secondary mineral phase was 1.1, 5.9, 7.2, 19.3, 32.1, and 28.1% at 5, 8, 22, 43, 70, and 159 h, respectively. Therefore, nanoscale precipitation resulting from CO2 injection could significantly alter the surface properties of pre-existing rocks. Chiquet and Brosetal observed wettability changes of caprock minerals (mica and quartz) after introducing CO2 under GCS conditions (31). They proposed that wettability change could be caused by a pH decrease owing to the dissolution of CO2 into the aqueous phase. The current study provides another possible explanation, the formation of a second mineral phase within several hours after introducing CO2. The alteration of caprock wettability can influence CO2 breakthrough pressures across the caprock, which can be crucial for the safety of CO2 sequestration. In this work, we studied a single mineral system containing an aluminosilicate mineral. However, in GCS sites, other minerals, such as calcite and dolomite (4, 12), could serve as a pH buffer to compensate for the pH drop due to CO2 dissolution, thus causing different morphological and transport behaviors of secondary mineral phases. Also, in addition to our experimental conditions, brine in GCS sites may have higher salinity than 1 M NaCl and may contain other anions and cations (e.g., Ca2+ and SO42-), which can complicate the dissolution/precipitation processes. Furthermore, experimental studies on the effect of flue gas impurities (e.g., H2S and SO2), on CO2-water-mineral interactions need to be conducted to provide more experimental data for numerical modeling (32). These are all important topics in environmentally sustainable GCS operation for future investigations.
Acknowledgments This work is supported by the Consortium for Clean Coal Utilization, and Y.S.J. is partially supported by the Director, Office of Science, Office of Basic Energy Sciences, of the U.S. Department of Energy, in association with the Energy Research Frontier Center (EFRC) of Berkeley Lab, under Contract No. DE-AC02-05CH11231. We thank John Rupp for providing Indiana geological information. We thank Mark Henson and Peter Colletti for testing the stability of the construction reactor materials. Finally, we thank the three anonymous reviewers for critical review of the manuscript.
Supporting Information Available Descriptive texts, figures, and tables of mineral characterization (Figure S1, S2, Table S1), experimental setup (Figures S3 and S4), sample preparation, data analysis, AFM images (Figures S5-S7), and phlogopite dissolution results under N2 (Table S2), and HR-TEM results for secondary mineral phase identification (Figure S8). This material is available free of charge via the Internet at http://pubs.acs.org.
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