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Unique environmental conditions required for dawsonite formation: Implications from dawsonite synthesis experiments under alkaline condition Yutaro Takaya, Miao Wu, and Yasuhiro Kato ACS Earth Space Chem., Just Accepted Manuscript • DOI: 10.1021/ acsearthspacechem.8b00121 • Publication Date (Web): 16 Jan 2019 Downloaded from http://pubs.acs.org on January 20, 2019

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ACS Earth and Space Chemistry

Unique environmental conditions required for dawsonite formation: Implications from dawsonite synthesis experiments under alkaline condition

Authors: Yutaro Takaya1-4*, Miao Wu1, Yasuhiro Kato3,2,4

Affiliations: 1

Department of Resources and Environmental Engineering School of Creative Science

and Engineering, Waseda University, 3-4-1 Okubo, Shinjyuku, Tokyo 169-8555, Japan 2 Ocean

Resources Research Center for Next Generation, Chiba Institute of Technology,

2-17-1 Tsudanuma, Narashino, Chiba 275-0016, Japan 3

Frontier Research Center for Energy and Resources, School of Engineering, The

University of Tokyo, 7-3-1 Hongo, Bunkyo-ku, Tokyo 113-8656, Japan 4 Research

and Development Center for Submarine Resources, Japan Agency for Marine-

Earth Science and Technology (JAMSTEC), 2-15 Natsushima-cho, Yokosuka, Kanagawa 237-0061, Japan

*Corresponding author 1

ACS Paragon Plus Environment

ACS Earth and Space Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Yutaro Takaya Department of Resources and Environmental Engineering School of Creative Science and Engineering, Waseda University, 3-4-1 Okubo, Shinjyuku, Tokyo 169-8555, Japan Telephone: +81-3-5286-3318 E-mail: [email protected]

Keywords: Dawsonite, Hydrotalcite, Manasseite, CO2 geological storage, Mineral trapping, Synthesis experiment

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Abstract: Although many numerical simulation studies suggest the formation of dawsonite in CO2 reservoirs and its resulting contribution to secure geological carbon storage, dawsonite formation is not observed in experimental studies. In addition, the natural occurrence of dawsonite is scarce. The lack of certainty in “whether dawsonite forms in CO2 reservoirs” is a major concern for evaluating the security of geological carbon storage and has been discussed for decades. This study performed dawsonite synthesis experiments with coexisting elements (K, Ca, and Mg) and investigated the unique formation conditions of dawsonite. Our experiments clearly show that co-existing magnesium (MgCl2) inhibits dawsonite formation to form hydrotalcite and/or manasseite instead of dawsonite under alkaline conditions. As magnesium is widespread in the Earth’s crust, dawsonite could be formed only under extremely restricted conditions (Mg-poor condition) and is unlikely to form in CO2 reservoirs during the post-injection period. We also indicate that the discrepancy between numerical simulations and experiments arises from the incompleteness of thermodynamic databases. Our results significantly contribute to resolving the long-running controversy regarding the formation of dawsonite in CO2 reservoirs.

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INTRODUCTION Dawsonite is a hydrated sodium aluminum carbonate with a chemical formula NaAl(CO3)(OH)2. Although dawsonite is composed only of five elements that occur abundantly in the Earth’s crust, natural occurrences of dawsonite are rare, indicating unique formation conditions. Dawsonite is mainly formed by the dissolution of feldspathic, quartzofeldspathic, and arkosic rocks under high CO2 fugacity1-4, and it is also observed as an inclusion mineral of Au-quartz veins5-6 and quartz7, a marine carbonate in sedimentary rocks8, and a weathering product of zeolite9. However, the formation conditions related to these natural occurrences are not so unique in geological environments, and thus cannot fully explain the scarcity of dawsonite occurrences. This uncertainty in the formation conditions of dawsonite has been fueling a considerable controversy on the security of geologic carbon sequestration for decades10-11. In geologic carbon sequestration, injected CO2 dissolves and ionizes in the reservoir water (solubility trapping), and reacts with surrounding host rocks to form carbonate minerals such as CaCO3 and MgCO3 (mineral trapping). This series of chemical reactions, called geochemical trapping, transforms injected CO2 into a more stable phase (carbonate minerals) and is, therefore, recognized as an important factor ensuring the security of geologic carbon sequestration12-15. In this sense, CO2-water-rock interactions under CO2 4


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reservoir conditions have been investigated extensively with laboratory experiments16-23, field tests24-29, and numerical simulations30-35 to evaluate the potential of geochemical trapping. Although many numerical simulation studies have predicted the formation of dawsonite in CO2 charged reservoirs30-35, dawsonite formation is rarely observed in laboratory experiments under CO2 reservoir conditions. Wolff-Boenisch and Galeczka36 reported

the

formation

of

dawsonite-type

carbonate—NH4-dawsonite,

NH4AlCO3(OH)2—under post-injection conditions (after increased pH via matrix dissolution and the first phase of carbonate formation), but Na-dawsonite was not observed in the experiments. This discrepancy between the numerical simulations and laboratory experiments led to the question of “whether dawsonite will form in CO2 reservoirs and contribute to mineral trapping”, and resolving this discrepancy is important for predicting the long-term security of geologic carbon sequestration. Hellevang et al. 11 pointed out that the nucleation and growth rate of dawsonite could be significantly lower than predicted and that dawsonite growth could have been overestimated in these studies. However, the formation conditions remain unclear despite laboratory experiments focusing on the dissolution rate/characteristics and thermodynamic properties of dawsonite37-41 as well as many field observation studies42-43. In this study, we focused on the effect of co-existing elements on the formation 5



of dawsonite because this aspect of dawsonite formation has not been explored thus far. We conducted a series of dawsonite synthesis experiments and discussed 1) the unique (paleo-)environmental conditions for dawsonite formation where dawsonite is observed in the present day, and 2) the potential impact of dawsonite formation on geological CO2 storage.

MATERIALS AND METHODS Dawsonite was synthesized with a batch type reactor, whose schematic is shown in Fig. 1 (MS Mini-reactor, OM Lab-Tech. Inc.). The autoclave is made of Hastelloy and is used with a Teflon inner vessel (80 cm3 in capacity). An Al(OH)3 suspension (colloid) was prepared by mixing AlCl3 solution (1M, 5 mL) and NaOH (3M, 5 mL) as the aluminum source for the synthesis of dawsonite. Then, NaHCO3 (1M, 40 mL) was added into the Al(OH)3 colloidal solution. A previous study revealed that superfluous MHCO3 (M=Na, K, NH4) is a requisite to the formation of single dawsonite-type compound (MAl(OH)2CO3) 37. Therefore, we set the molar ratio of NaHCO3/Al to be 8. After mixing this solution, 5 mL milliQ-water or 5 mL 1M Mx+Clx solution (M: K, Ca, and Mg as a coexisting element) was further added into the reaction system. Finally, 3M NaOH was added to adjust the initial pH (10.0 ± 0.1). The abovementioned pre-heating procedure 6


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was conducted within PE bottles. After pouring the mixed starting solution into the Teflon inner vessel, the autoclave was sealed and kept at pre-defined temperatures (80, 120 and 160℃) for 24 h. Synthesis experiments were also conducted with 0.01 M and 0.1 M Mx+Clx solution at 120℃. After cooling the autoclave to 50℃, the reacted solution was recovered, pH was measured, and the solution was separated into liquid and solid samples using a 0.45 µm filter. The liquid sample was diluted with 2 wt.% mixed acid (nitric acidhydrochloric acid-hydrofluoric acid (20:5:1)) for ICP-MS analysis (iCAP Q, Thermo Fisher Scientific Inc.). The solid sample was washed with milliQ water and dried at 70℃ for X-ray diffractometer (RINT-ULTIMA III, Rigaku Corporation) analysis and observed with SEM/EDS (SEM: Miniscope TM3000, Hitachi High-Technologies Corporation, EDS: Quantax70, Bruker Corporation). The XRD measurements were performed at a scanning speed of 2θ = 2°/min (step width 0.02°) over a range of 2θ = 3–70°. Each experiment was conducted twice to confirm the consistency of experiments (Table 1). As described above, we conducted the synthesis experiments under alkaline conditions. In CO2 geological storage, CO2-water-rock interactions proceed mainly under weak-acidic conditions. However, the pH condition of the reaction front may reach neutral to alkaline conditions. In addition, entire CO2 reservoir conditions approach neutral to alkaline conditions via the progress of geochemical trapping (i.e. post-injection 7



conditions).36

RESULTS Chemistry of the Reacted Solution: The water chemistry of the reacted solution determined by ICP-MS analysis are shown in Table 1. Remaining percentages of each element (100*[Concentration in reacted solution]/[Concentration in the starting solution]) are also shown in Table 1. The results showed that Al was completely consumed to form some sort of minerals in all the experiments, leaving no residual Al in the reacted solution. Ca and Mg were also mostly consumed to form minerals under all temperature conditions, leaving no residuals (Ca: 0.00–2.31 %, Mg: 0.16–4.08 %). In contrast, most of K remained in the reacted solution (75.08–100.00 %). The amount of residual Na in the reacted solutions decreased with increasing temperature (average 88.6 % at 80℃, 68.9 % at 120℃, 62.7 % at 160℃). In all cases except for MgCl2 experiments, the pH of the reacted solution remained unchanged throughout the reactions. On the contrary, the pH of the reacted solution decreased to 8.76–8.82 (80℃), 8.85–8.88 (120℃), and 9.52–9.76 (160℃) from 10 (pH of the starting solution) in experiments with 1M MgCl2 solution. XRD Analysis and SEM/EDS Observation of Solid Samples: X-ray diffractograms and electron micrographs of the solid samples are displayed in Fig. 2–4 and Fig. S1–18. 8


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In all the experiments without co-existing elements (Runs 01-02, 09-10, 29-30), the presence of dawsonite was confirmed through XRD (Figs. 2, Fig. S1, S5, and S15). Halite (NaCl), thermonatrite (Na2CO3·H2O), and trona (Na3H(CO3)2·2H2O) were also detected. No Al-bearing mineral other than dawsonite was detected in the solid samples, except for Run-29, in which quite weak peaks of bayerite (Al(OH)3) were detected in an X-ray diffractogram (Fig. S15). The intensity and the sharpness of dawsonite peaks in the diffractograms increased with increasing reaction temperature (Fig. 2), indicating that the crystallinity of dawsonite improved with increasing temperature. The SEM/EDS observation also demonstrated that the dawsonite crystal grew clearer at higher temperatures. Although fibrous dawsonite was observed in solid samples from the experiments at 120℃ and 160℃, dawsonite (Al-bearing materials in EDS) at lower temperatures (80℃) did not exhibit any clear shape (Fig. 2). Acicular/columnar shaped crystals (thermonatrite and trona) were also observed in the products from the lower temperature experiments. Experiments with K (KCl) (Runs 03-04, 11-16, 31-32) resulted in the same products as those without the coexisting elements (Figs. S2, S6-8, and S16). The concentration of K in the starting solution did not affect the types and combinations of the formed materials. However, calcite/aragonite (CaCO3) formed and precipitated in addition to dawsonite in the experiments with Ca (CaCl2) (Runs 05-06, 17-22, 33-34) (Fig. 9



3). The crystallinity of calcite/aragonite was greater than that of dawsonite, especially at lower temperatures (Fig. 3). Cubic calcite crystals were also observed in products at lower temperatures (80℃, 120℃) (Fig. 3). Although, other than dawsonite, Al-bearing minerals were not detected in the experiments with K (KCl) and Ca (CaCl2), hydrotalcite (Mg6Al2(CO3)(OH)16·4(H2O)) was observed in the experiments with Mg (MgCl2) (Runs 07-08, 23-28, 35-36) (Figs. 4, 5). Minor amount of manasseite, which has the same chemical composition as hydrotalcite, was also detected at 160℃ (Fig. 4). In these experiments, hydrotalcite (and dawsonite) did not exhibit any specific crystal shape. EDS mapping showed that hydrotalcite and dawsonite coexisted on the surface of the formed material (Fig. S19). No Mg-bearing minerals other than hydrotalcite and manasseite were detected in the solid samples.

DISCUSSION Effect of Coexisting Elements on the Formation of Dawsonite: In experiments without the coexisting elements, Al in the starting solution was completely consumed to form dawsonite (dawsonite formation efficiency: 100 %), because no Al-bearing mineral other than dawsonite was detected in the solid samples. While the remaining percentages of Na in the reacted solution were over 90% at 80℃, the percentages decreased to approximately 10


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60% at 160℃ with increased decomposition of NaHCO3 in the starting solution. Our results indicate that K did not affect dawsonite formation in our experiments. In the experiments with Ca, the removed Ca was precipitated as aragonite and/or calcite. In these experiments, Al-bearing minerals other than dawsonite were not observed in the reaction products, and therefore, the dawsonite formation efficiency was 100% even with Ca. If the amount of coexisting Ca is much larger than that of our experiments, Ca could indirectly affect dawsonite formation by forming Ca-carbonates by consuming carbonate or bicarbonate ions in the solution. In the experiments with Mg, the formation of hydrotalcite and manasseite were confirmed. Although the formation of dawsonite was also confirmed in a series of experiments, the formation efficiency of dawsonite was lowered by the production of Al-bearing hydrotalcite and manasseite. At 120℃, the peak intensity of hydrotalcite increased with increasing concentrations of MgCl2 in the starting solution, while the peak intensity of dawsonite clearly decreased (Fig. 5). In addition, no Mg-bearing mineral other than hydrotalcite and manasseite was observed in the solid sample. These results suggest that hydrotalcite and manasseite are preferentially formed over dawsonite and inhibit the synthesis of dawsonite by consuming Al under weakly alkaline conditions. In experiments with MgCl2 solution, the pH of the solution clearly decreased 11



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throughout the experiments. In our experimental system, the pH of the solution decreases with the formation of carbonate minerals because H+ is released, as shown in the following equation (cases of Ca-carbonate); 2NaHCO3 + CaCl2 → 2NaCl + CaCO3 + HCO3- + H+

(eq. 1)

On the other hand, when hydrotalcite/manasseite forms, more H+ is released (eq. 1); 2NaHCO3 + MgCl2 + 7/3H2O → 2NaCl + 1/6Mg6Al2(CO3)(OH)16·4(H2O) + 11/6HCO3- + 11/6H+ (eq. 2) Therefore, the solution pH appears to be lower after reaction with MgCl2 than in other experiments. Unique Environmental Conditions Required for Dawsonite Formation: Our results show that a specific condition required for dawsonite formation is an “extremely Mgpoor environment”. Because Mg is abundant in the Earth’s crust, the condition of “extremely Mg-poor environment” is very rare. Therefore, this environmental condition could be the primary reason of the scarce occurrence of dawsonite. Mg-carbonate has many

thermodynamically

metastable

phases

such

as

hydromagnesite

[4MgCO3·Mg(OH)2·4H2O], dypingite [4MgCO3·Mg(OH)2·(5-8)H2O], and nesquehonite [Mg(OH)(HCO3)·2H2O].44-45 Hydrotalcite/manasseite may also transform to another metastable phases or magnesite in response to the surrounding environmental conditions 12


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(groundwater chemistry, temperature, and pressure). Li et al.41 conducted dissolution experiments of dawsonite under acidic conditions (high PCO2 conditions). They showed that dissolution/decomposition of dawsonite proceeds rapidly at temperatures exceeding 120℃. This result also shows that non-exposure to acidic high-temperature fluids over 120℃ could also be a required environmental condition for dawsonite preservation in nature. Some large-scale natural occurrence sites of dawsonite have been observed, such as Green River formation in Colorado46 and Springerville-St. Johns Field (USA)1. In the Green River Formation, salt minerals trona (Na3H(CO3)2·2H2O), nahcolite (NaHCO3), dawsonite, halite and less abundant eitelite (Na2CO3·MgCO3), and wegscheiderite (Na5(CO3)(HCO3)3) were precipitated under an evaporitic alkaline environment.46 Dawsonite occurred with eitelite, but the amount of eitelite is minor. In the SpringervilleSt. Johns Field, on the other hand, dawsonite-bearing siltstone comprises quartz, plagioclase, K-feldspar, calcite, gypsum, dolomite, hematite, kaolinite, illite, muscovite, biotite, and dawsonite.1 The amount of dawsonite reaches 5–17 wt.%. Mg-bearing carbonate (dolomite) is observed in the same siltstone, but the textural relationships demonstrate that kaolinite and dawsonite are younger than the other carbonate minerals and grow on these minerals. Dawsonite is also found in large-scale in the Hailaer Basin, 13



China.42 In this basin, dawsonite occurs as a cement filling the pores between detrital grains and as a replacement of feldspars with quartz overgrowths, microcrystalline quartz, ankerite and clay minerals in sandstone; the amount of ankerite (Mg-bearing carbonate) is very minor. These natural occurrences of dawsonite (the coexisting amount of Mg with dawsonite is small in Green River and Hailaer Basin, the generation timing of Mg-bearing carbonate differ with that of dawsonite in Springerville - St. Johns Field) suggest that the natural dawsonite formed under Mg-poor environmental conditions consistent with our experimental results. Contradiction Between Laboratory and Numerical Studies: Our experiments revealed that magnesium inhibits the formation of dawsonite by forming hydrotalcite and manasseite under alkaline conditions. We conducted some numerical calculations by using PHREEQC (ver. 3.4.0)47 with LLNL thermo database, which is a thermodynamic database provided by the Lawrence Livermore National Laboratory and is used in many numerical simulation studies, to simulate the results of our study. Here we calculated the saturation indices of minerals in the reacted solution. In our calculation, the amount of additional 3M NaOH for adjusting the starting pH is assumed to be 4 ml (the average value). Table 2 shows the calculated saturation indices in the cases of (a) 0.01 mol/L, 0.1 mol/L, and 1.0 mol/L MgCl2 solution added as a co-existing element under at 120℃ and 14


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(b) 1.0 mol/L MgCl2 solution at 80 ℃, 120℃, and 160℃. Minerals showing positive saturation indices are predicted to be saturated in the solution and then precipitate. In Table 2(a), the saturation indices of dawsonite show almost same values regardless of MgCl2 concentration in the starting solution, indicating no influence of Mg on the formation of dawsonite. In the cases of Mg-bearing minerals, saturation indices of the metastable

phases

of

magnesium

carbonates

such

as

hydromagnesite

(Mg5(CO3)4(OH)2·4H2O) as well as magnesite (MgCO3) definitely show positive values (Table 2(a), (b)). However, in the actual experimental system, magnesium precipitated as hydrotalcite/manasseite and inhibited the formation of dawsonite by consuming aluminum in the solution. This provides a clear explanation for the discordance between the experimental studies and the numerical calculations. Relevant phases such as hydrotalcite and manasseite are not included in the thermodynamic database and therefore, calculations using the database artificially stabilize dawsonite instead of showing the formation of these phases. Thus, extending the thermodynamic database to include phases such as hydrotalcite and manasseite would be the most important step toward improving the agreement between experiments and calculations.

CONCLUSIONS 15



Our experiments demonstrate that dawsonite formation is inhibited by coexisting magnesium, which forms hydrotalcite and manasseite, under alkaline conditions. Therefore, an “extremely Mg-poor environment” is required for dawsonite formation. As magnesium is one of the most common elements in the Earth’s crust, dawsonite is unlikely to form and contribute to mineral trapping. This required environmental condition can consistently explain the scarcity of the natural occurrences of dawsonite. However, at the conditions considered in this study, hydrotalcite and/or manasseite would form in CO2 reservoirs and might contribute to the security of CO2 storage instead of dawsonite. The discord between experimental and numerical studies on dawsonite formation is caused by the lack of thermodynamic data. Furthermore, the long-term safety of CO2 storage (mineral trapping of CO2) by dawsonite formation is highly unlikely. Therefore, the next step is to obtain thermodynamic and reaction kinetics data of hydrotalcite/manasseite as well as some other minor and meta-stable minerals for improving the numerical simulation and assessing the quality and safety for CO2 storage. It is also necessary to conduct dawsonite synthesis experiments with other elements as co-existing materials and under wider pH ranges (weakly acidic to neutral) to clarify the stability relations of dawsonite. Nevertheless, this study provides a basis for resolving the long-running controversy about the formation of dawsonite in CO2 reservoirs. 16


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Acknowledgments This work was partly supported by JSPS KAKENHI Grant Numbers 25820432, 14J09158 and 18K18213 for Y.T.

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Supporting Information This manuscript has a supporting information (Figs. S1-S19) which shows the X-ray diffractograms (Figs. S1-S18), and SEM images, EDS mapping image, and energy spectrum (Fig. S19) of solid samples from synthesis experiments.

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Study of Basaltic Glass–H2O–CO2 Interaction at 22 and 50 ℃: Implications for Subsurface Storage of CO2. Geochim. Cosmochimica Acta 2014, 126, 123–145. (20) Huq, F.; Haderlein, S. B.; Cirpka, O. A.; Nowak, M.; Blum, P.; Grathwohl, P. Flowthrough Experiments on Water–Rock Interactions in a Sandstone caused by CO2 Injection at Pressures and Temperatures Mimicking Reservoir Conditions. Appl. Geochem. 2015, 58, 136–146. (21) Takaya, Y.; Nakamura, K.; Kato, Y. Dissolution of Altered Tuffaceous Rocks under Conditions Relevant for CO2 Storage. Appl. Geochem. 2015, 58, 78–87. (22) Takaya, Y.; Nakamura, K.; Kato, Y. Long‐Term Reaction Characteristics of CO2– Water–Rock Interaction: Insight into the Potential Groundwater Contamination Risk from Underground CO2 Storage. Resour. Geol. 2018, 68(1), 93–100. (23) Menefee, A. H.; Giammar, D. E.; Ellis, B. R. Permanent CO2 Trapping through Localized and Chemical Gradient-Driven Basalt Carbonation. Environ. Sci. Technol. 2018, 52(15), 8954-8964. (24) Kharaka, Y. K.; Cole, D. R.; Hovorka, S. D.; Gunter, W. D.; Knauss, K. G.; Freifeld, B. M. Gas-Water-Rock Interactions in Frio Formation Following CO2 Injection: Implications for the Storage of Greenhouse Gases in Sedimentary Basins. Geology 2006, 34(7), 577–580. 22


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(25) Kharaka, Y. K.; Cole, D. R.; Thordsen, J. J.; Kakouros, E.; Nance, H. S. Gas–Water– Rock Interactions in Sedimentary Basins: CO2 Sequestration in the Frio Formation, Texas, USA. J. Geochem. Explor. 2006, 89(1-3), 183–186. (26) Matter, J. M.; Takahashi, T.; Goldberg, D. Experimental Evaluation of In Situ CO2‐Water-Rock Reactions during CO2 Injection in Basaltic Rocks: Implications for Geological CO2 Sequestration. Geochem. Geophy. Geosy. 2007, 8(2), doi: 10.1029/2006GC001427. (27) Gislason, S. R.; Wolff-Boenisch, D.; Stefansson, A.; Oelkers, E. H.; Gunnlaugsson, E., Sigurdardottir, H.; Sigfusson, B.; Broecker, W. S.; Matter, J. M.; Stute, M.; Axelsson, G.; Fridriksson, T. Mineral Sequestration of Carbon Dioxide in Basalt: A Pre-Injection Overview of the CarbFix Project. Int. J. Greenh. Gas Con. 2010, 4(3), 537–545. (28) Matter, J. M.; Stute, M.; Snæbjörnsdottir, S. Ó.; Oelkers, E. H.; Gislason, S. R.; Aradottir, E. S.; Sigfusson, B.; Gunnarsson, I.; Sigurdardottir, H.; Gunnlaugsson, E.; Axelsson, G.; Alfredsson, H. A.; Wolff-Boenisch, D.; Mesfin, K.; Taya, D. F. D.; Hall, J.; Dideriksen, K.; Broecker, W. S. Rapid Carbon Mineralization for Permanent Disposal of Anthropogenic Carbon Dioxide Emissions. Science 2016, 352(6291), 1312–1314. (29) Mito, S.; Xue, Z.; Ohsumi, T. Case Study of Geochemical Reactions at the Nagaoka CO2 Injection Site, Japan. Int. J. Greenh. Gas Con. 2008, 2(3), 309–318. 23



(30) Xu, T.; Apps, J. A.; Pruess, K. Reactive Geochemical Transport Simulation to Study Mineral Trapping for CO2 Disposal in Deep Arenaceous Formations. J. Geophys. Res.: Sol. Ea. 2003, 108(B2), doi: 10.1029/2002JB001979. (31) Xu, T.; Apps, J. A.; Pruess, K. Numerical Simulation of CO2 Disposal by Mineral Trapping in Deep Aquifers. Appl. Geochem. 2004, 19(6), 917–936. (32) Xu, T.; Sonnenthal, E.; Spycher, N.; Pruess, K. TOUGHREACT—a Simulation Program for Non-Isothermal Multiphase Reactive Geochemical Transport in Variably Saturated Geologic Media: Applications to Geothermal Injectivity and CO2 Geological Sequestration. Comput. Geosci. 2006, 32(2), 145–165. (33) White, S. P.; Allis, R. G.; Moore, J.; Chidsey, T.; Morgan, C.; Gwynn, W.; Adams, M. Simulation of Reactive Transport of Injected CO2 on the Colorado Plateau, Utah, USA. Chem. Geol. 2005, 217(3-4), 387–405. (34) Zerai, B.; Saylor, B. Z.; Matisoff, G. Computer Simulation of CO2 Trapped through Mineral Precipitation in the Rose Run Sandstone, Ohio. Appl. Geochem. 2006, 21(2), 223–240. (35) Okuyama, Y.; Todaka, N.; Sasaki, M.; Ajima, S.; Akasaka, C. Reactive Transport Simulation Study of Geochemical CO2 Trapping on the Tokyo Bay Model–With Focus on the Behavior of Dawsonite. Appl. Geochem. 2013, 30, 57–66. 24


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(36) Wolff-Boenisch, D.; Galeczka, I. M. Flow-through Reactor Experiments on Basalt(Sea) Water-CO2 Reactions at 90℃ and Neutral pH. What Happens to the Basalt Pore Space under Post-Injection Conditions? Int. J. Greenh. Gas Con. 2018, 68, 176–190. (37) Zhang, X.; Wen, Z.; Gu, Z.; Xu, X.; Lin, Z. Hydrothermal Synthesis and Thermodynamic Analysis of Dawsonite-Type Compounds. J. Solid State Chem. 2004, 177(3), 849–855. (38) Hellevang, H.; Aagaard, P.; Oelkers, E. H.; Kvamme, B. Can Dawsonite Permanently Trap CO2? Environ. Sci. Technol. 2005, 39(21), 8281–8287. (39) Hellevang, H.; Declercq, J.; Kvamme, B.; Aagaard, P. The Dissolution Rates of Dawsonite at pH 0.9 to 5 and Temperatures of 22, 60 and 77℃. Appl. Geochem. 2010, 25(10), 1575–1586. (40) Bénézeth, P.; Palmer, D. A.; Anovitz, L. M.; Horita, J. (2007). Dawsonite Synthesis and Reevaluation of its Thermodynamic Properties from Solubility Measurements: Implications for Mineral Trapping of CO2. Geochim. Cosmochimica Acta, 2007, 71(18), 4438–4455. (41) Li, F.; Cao, Y.; Li, W.; Zhang, L. CO2 mineral trapping: Hydrothermal Experimental Assessments on the Thermodynamic Stability of Dawsonite at 4.3 MPa pCO2 and Elevated Temperatures. Greenh. Gases Sci. Technol. 2018, 8(1), 77–92. 25



(42) Gao, Y.; Liu, L.; Hu, W. Petrology and Isotopic Geochemistry of Dawsonite-Bearing Sandstones in Hailaer Basin, Northeastern China. Appl. Geochem. 2009, 24(9), 1724– 1738. (43) Liu, N.; Liu, L.; Qu, X.; Yang, H.; Wang, L.; Zhao, S. Genesis of Authigene Carbonate Minerals in the Upper Cretaceous Reservoir, Honggang Anticline, Songliao Basin: A Natural Analog for Mineral Trapping of Natural CO2 Storage. Sediment. Geol. 2011, 237(3-4), 166–178. (44) Hänchen, M.; Prigiobbe, V.; Baciocchi, R.; Mazzotti, M. Precipitation in the MgCarbonate System—Effects of Temperature and CO2 Pressure. Chem. Eng. Sci. 2008, 63(4), 1012–1028. (45) Moore, J. K.; Surface, J. A.; Brenner, A.; Skemer, P.; Conradi, M. S.; Hayes, S. E. Quantitative Identification of Metastable Magnesium Carbonate Minerals by Solid-State 13C NMR Spectroscopy. Environ. Sci. Technol. 2014, 49(1), 657–664. (46) Dyni, J. R. Sodium Carbonate Resources of the Green River Formation. US Geological Survey Open-File Report, 1996, 729, 39. (47) Parkhurst, D. L.; Appelo, C. A. J., 2013, Description of input and examples for PHREEQC version 3 - A computer program for speciation, batch-reaction, onedimensional transport, and inverse geochemical calculations: U.S. Geological Survey 26


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Techniques and Methods, book 6, chap. A43, 497 p.

27



Figure and Table Captions: Figure Captions: Fig. 1 Schematic of the autoclave used in synthesis experiments. Fig. 2 X-ray diffractograms and electron micrographs of solid samples from the synthesis experiments without co-existing elements. (A) Run-01: 80℃, (B) Run-09: 120℃, and (C) Run-29:160℃. The small whitish crystals with a cruciform shape in the electron micrograph (B) are halite. Fig. 3 X-ray diffractograms and electron micrographs of solid samples from synthesis experiments with Ca (1M CaCl2 solution) as a co-existing element. (A) Run-06: 80 ℃, (B) Run-21: 120℃, and (C) Run-34: 160℃. The acicular/columnar shaped crystals in the electron micrograph (B) are trona. Rhombic shaped crystals are calcite. Fig. 4 X-ray diffractograms and electron micrographs of solid samples from synthesis experiments with Mg (1M MgCl2 solution) as a co-existing element. (A) Run-08: 80 ℃, (B) Run-23: 120℃, and (C) Run-35: 160℃. Reacted solid samples at 80℃ and 120 ℃ show hazy surfaces, without any clear crystal shape. EDS mapping shows that all elements composing dawsonite and hydrotalcite (Na, Mg, Al, C, and O) are included in this “hazy surface” (show the supplementary Fig. S19). 28


Page 28 of 38

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Fig. 5 X-ray diffractograms and electron micrographs of solid samples from synthesis experiments with Mg (MgCl2 solution) as a co-existing element. (A) Run-23: 0.01 M MgCl2, (B) Run-25: 0.1 M MgCl2, and (C) Run-27: 1.0 M MgCl2. The intensity of dawsonite decreases with increasing MgCl2 concentration and hydrotalcite intensity.

Table Captions: Table 1 Experimental conditions and water chemistry of the reacted solutions. The columns (%) next to each element (Na, Al, and Me) show the remaining percentage of each element (100*[Concentration in reacted solution]/[Concentration in the starting solution]) in the solution after reaction. Table 2 Saturation Indices of minerals in the reacted solutions with MgCl2 calculated by PHREEQC with LNLL thermo database. Positive values indicate supersaturation of each mineral. Calculated saturation indices of minerals in the cases of (a) 0.01 mol/L, 0.1 mol/L and 1.0 mol/L MgCl2 solution as a co-existing element at temperatures of 120 ℃, and (b) 1.0 mol/L MgCl2 solution added at 80 ℃, 120℃, and 160℃.

29



Page 30 of 38

Table 1 Experimental conditions and water chemistry of the reacted solutions. The columns (%) next to each element (Na, Al, and Me) show the remaining percentage of each element (100*[Concentration in reacted solution]/[Concentration in the starting solution]) in the solution after reaction. Run No.

Reaction Temp. [℃]

none (milliQ water)

01 02 03 04 05

1.0M KCl 80 1.0M CaCl2

06 07

1.0M MgCl2

08

none (milliQ water)

09 10 11

0.01M KCl

12 13

0.1M KCl

14 15

1.0M KCl

16 17 18 19

0.01M CaCl2 120 0.1M CaCl2

20 21

1.0M CaCl2

22 23

0.01M MgCl2

24 25

0.1M MgCl2

26 27

1.0M MgCl2

28

none (milliQ water)

29 30 31 32 33 34 35 36

x+

Me Clx Type

1.0M KCl 160 1.0M CaCl2 1.0M MgCl2

pH of the Start Solution

pH of the Reacted Soluiton

10.05

Water Chemistry of the Reacted Solution Na

Al

Me (K or Ca or Mg)

[mol/l]

[%]

[mmol/l]

[%]

[mmol/l]

[%]

10.02

1.01

91.22

0.00

0.00

-

-

9.94

10.00

1.01

91.36

0.02

0.02

-

-

9.91

9.95

0.98

86.64

0.00

0.00

80.94

95.51

9.93

9.95

0.97

85.30

0.00

0.00

79.21

93.47

10.05

10.05

1.05

92.30

0.00

0.00

0.07

0.09

9.94

10.00

1.04

91.77

0.05

0.06

0.00

0.00

9.96

8.76

1.03

86.05

0.00

0.00

0.47

0.57

10.05

8.82

1.02

84.04

0.00

0.00

0.45

0.55

9.95

10.01

0.84

72.04

0.52

0.62

-

-

9.97

10.03

0.80

70.14

0.00

0.00

-

-

9.94

10.00

0.82

70.08

0.09

0.11

0.62

75.08

9.91

9.98

0.85

72.53

0.14

0.17

0.88

100.00

9.90

9.94

0.78

69.67

0.00

0.00

8.30

97.07

9.95

9.96

0.77

67.86

0.09

0.11

8.35

98.14

9.93

9.95

0.80

71.74

0.00

0.00

79.25

92.72

9.92

9.96

0.81

72.23

0.05

0.05

82.17

96.14

9.95

9.82

0.76

68.80

0.00

0.00

0.00

0.00

9.93

9.83

0.73

66.08

0.00

0.00

0.00

0.00

9.93

9.99

0.74

67.06

0.13

0.15

0.20

2.31

9.93

10.02

0.79

66.97

0.04

0.05

0.00

0.00

9.90

10.00

0.83

68.96

0.04

0.05

0.00

0.00

9.91

9.95

0.81

68.86

0.14

0.16

0.00

0.00

9.95

9.99

0.76

65.38

0.00

0.00

0.03

3.96

10.00

10.02

0.81

68.18

0.00

0.00

0.03

4.08

9.96

9.92

0.81

68.97

0.00

0.00

0.13

1.53

9.97

10.00

0.80

68.29

0.04

0.05

0.07

0.81

9.90

8.75

0.83

71.52

0.00

0.00

0.20

0.24

9.90

8.88

0.73

63.28

0.04

0.05

0.21

0.25

9.92

9.93

0.68

61.34

0.06

0.07

-

-

9.91

9.91

0.67

60.40

0.00

0.00

-

-

10.00

10.00

0.70

62.20

0.00

0.00

98.62

100.00

9.95

9.95

0.67

60.28

0.22

0.26

90.88

100.00

9.93

9.94

0.71

63.72

0.06

0.07

0.00

0.00

9.95

9.96

0.75

66.18

0.00

0.00

0.00

0.00

9.95

9.76

0.72

63.10

0.00

0.00

0.00

0.00

10.09

9.52

0.77

64.62

0.10

0.13

0.00

0.00


Page 31 of 38 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 2 Saturation Indices of minerals in the reacted solutions with MgCl2 calculated by phreeqc with LNLL thermo database. Positive values indicate supersaturation of each mineral. Calculated saturation indices of minerals in the cases of (a) 0.01 mol/L, 0.1 mol/L and 1.0 mol/L MgCl2 solution as a co-existing element at temperatures of 120 ℃. Saturation Index

Mineral

Chemical formula

Artinite

Mg2CO3(OH)2・3H2O

-2.72

-0.72

1.29

Bischofite

MgCl2・6H2O

-10.68

-9.68

-8.27

Boehmite

AlO2H

4.24

4.24

4.25

Brucite

Mg(OH)2

-2.21

-1.21

-0.18

Chloromagnesite

MgCl2

-20.82

-19.81

-18.40

Corundum

Al2O3

7.53

7.53

7.54

Dawsonite

NaAlCO3(OH)2

5.11

5.10

5.07

Diaspore

AlHO2

4.51

4.51

4.52

Gibbsite

Al(OH)3

3.73

3.73

3.73

Halite

NaCl

-2.58

-2.58

-2.39

Hydromagnesite

Mg5(CO3)4(OH)2・4H2O

-0.01

4.99

9.98

Lansfordite

MgCO3・5H2O

-2.86

-1.86

-0.88

Magnesite

MgCO3

1.84

2.84

3.83

Na2CO3

Na2CO3

-3.23

-3.23

-3.28

Na2CO3・7H2O

Na2CO3・7H2O

-6.12

-6.12

-6.17

Na2O

Na2O

-39.09

-39.09

-39.10

Nahcolite

NaHCO3

-1.15

-1.15

-1.19

Natron

Na2CO3・10H2O

-6.22

-6.22

-6.28

Nesquehonite

MgCO3・3H2O

-2.55

-1.55

-0.56

Oxychloride-Mg

Mg2Cl(OH)3・4H2O

-14.77

-12.76

-10.52

Periclase

MgO

-5.53

-4.53

-3.50

Spinel

Al2MgO4

2.73

3.73

4.77

Thermonatrite

Na2CO3・H2O

-3.63

-3.63

-3.68

o

120 C 0.01M MgCl2


120oC 0.1M MgCl2

120oC 1M MgCl2


Page 32 of 38

Table 2 Saturation Indices of minerals in the reacted solutions with MgCl2 calculated by phreeqc with LNLL thermo database. Positive values indicate supersaturation of each mineral. Calculated saturation indices of minerals in the cases of (b) 1.0 mol/L MgCl2 solution added at 80 ℃, 120℃, and 160℃. Saturation Index

Mineral

Chemical formula

Artinite

Mg2CO3(OH)2・3H2O

-0.10

1.29

2.21

Bischofite

MgCl2・6H2O

-7.93

-8.27

-8.70

Boehmite

AlO2H

5.23

4.25

3.47

Brucite

Mg(OH)2

-1.53

-0.18

0.80

Chloromagnesite

MgCl2

-20.75

-18.40

-16.42

Corundum

Al2O3

8.49

7.54

7.06

Dawsonite

NaAlCO3(OH)2

6.70

5.07

3.64

Diaspore

AlHO2

5.55

4.52

3.69

Gibbsite

Al(OH)3

4.79

3.73

2.94

Halite

NaCl

-2.40

-2.39

-2.35

Hydromagnesite

Mg5(CO3)4(OH)2・4H2O

6.50

9.98

12.40

Lansfordite

MgCO3・5H2O

-0.62

-0.88

-1.24

Magnesite

MgCO3

3.31

3.83

4.19

Na2CO3

Na2CO3

-3.90

-3.28

-2.80

Na2CO3・7H2O

Na2CO3・7H2O

-5.08

-6.17

-7.44

Na2O

Na2O

-44.26

-39.10

-34.98

Nahcolite

NaHCO3

-0.97

-1.19

-1.37

Natron

Na2CO3・10H2O

-5.19

-6.28

-7.37

Nesquehonite

MgCO3・3H2O

-0.35

-0.56

-0.91

Oxychloride-Mg

Mg2Cl(OH)3・4H2O

-9.96

-10.52

-11.20

Periclase

MgO

-5.47

-3.50

-2.00

Spinel

Al2MgO4

4.43

4.77

5.00

Thermonatrite

Na2CO3・H2O

-4.06

-3.68

-3.43

80oC, 1M MgCl2


120oC, 1M MgCl2

160oC, 1M MgCl2

Page 33 of 38 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


relief valve pressure gauge

thermocouple

autoclave

heater

temparature controller

Fig. 1 Schematic of the autoclave used in synthesis experiments.


ACS Earth and Space Chemistry 120,000

(A)

Intensity (cps)

100,000

Run-01 (80oC)

dawsonite (NaAlCO3(OH)2) halite (NaCl) thermonatrite (Na2CO3·H2O)

80,000 60,000 40,000

10 μm

20,000 0 120,000

(B)

Intensity (cps)

100,000

Run-09 (120oC)

trona (Na3H(CO3)2·2H2O)

80,000 60,000 40,000

10 μm

20,000 0 400,000

Run-29 (160oC)

300,000

(C)

Intensity (cps)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47

Page 34 of 38

200,000 100,000 0

10 μm 20

2θ (o)

40

60

Fig. 2 X-ray diffractograms and electron micrographs of solid samples from the synthesis experiments without co-existing elements. ACS Paragon Plus Environment (A) Run-01: 80℃, (B) Run-09: 120℃, and (C) Run-29:160℃. The small whitish crystals with a cruciform shape in the electron micrograph (B) are halite.

Page 35 of 38


dawsonite halite calcite aragonite

Run-06 (80oC)

(A)

Intensity (cps)

300,000 200,000 100,000

20 μm

0

20

140,000

(B)

Intensity (cps)

120,000

2θ ( )

40

60

o

Run-21 (120oC)

100,000 80,000 60,000 40,000

10 μm

20,000 0

20

400,000

2θ ( )

40

60

o

Run-34 (160oC) 300,000

(C)

Intensity (cps)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47

200,000 100,000 0

10 μm 20

2θ (o)

40

60

Fig. 3 X-ray diffractograms and electron micrographs of solid samples fromACS synthesis with Ca (1M CaCl2 solution) as a co-existing element. Paragonexperiments Plus Environment (A) Run-06: 80 ℃, (B) Run-21: 120℃, and (C) Run-34: 160℃. The acicular/columnar shaped crystals in the electron micrograph (B) are trona. Rhombic shaped crystals are calcite.


140,000

(A)

Intensity (cps)

120,000

Run-08 (80oC)

dawsonite hydrotalcite

100,000 80,000 60,000 40,000

20 μm

20,000 0 400,000

20

40

60

Run-27 (120oC)

(B)

Intensity (cps)

300,000

200,000

100,000

0 600,000 500,000

(C)

Intensity (cps)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47

Page 36 of 38

20 μm 20

40

Run-35 (160oC)

60

manasseite

400,000 300,000 200,000

10 μm

100,000 0

20

2θ (o)

40

60

Fig. 4 X-ray diffractograms and electron micrographs of solid samples from synthesis experiments with Mg (1M MgCl2 solution) as a co-existing element. ACSsolid Paragon Plus Environment (A) Run-08: 80 ℃, (B) Run-23: 120℃, and (C) Run-35: 160℃. Reacted samples at 80℃ and 120 ℃ show hazy surfaces, without any clear crystal shape. EDS mapping shows that all elements composing dawsonite and hydrotalcite (Na, Mg, Al, C, and O) are included in this “hazy surface” (show the supplementary Fig. S19).

Page 37 of 38


Run-23 (0.01 M MgCl2)

dawsonite

(A)

Intensity (cps)

150,000

100,000

50,000

0 140,000

(B)

Intensity (cps)

120,000

20

40

Run-25 (0.1 M MgCl2)

60

dawsonite hydrotalcite

100,000 80,000 60,000 40,000 20,000 0 400,000

20

40

Run-27 (1.0 M MgCl2)

(C)

60

dawsonite halite hydrotalcite

300,000 Intensity (cps)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47

200,000

100,000

0

20

40

60

2θ ( ) Fig. 5 X-ray diffractograms and electron micrographs of solid samples from synthesis experiments with Mg (MgCl2 solution) as a co-existing element. (A) Run-23: 0.01 M MgCl2, (B) Run-25: 0.1 M MgCl2, andACS (C)Paragon Run-27:Plus 1.0Environment M MgCl2. The intensity of dawsonite decreases with increasing MgCl2 concentration and hydrotalcite intensity. o


100,000

50,000

0 140,000

decreasing intensity

Intensity (cps)

150,000

Increasing intensity

200,000

dawsonite with 0.01 M MgCl2 (NaAl(CO3)(OH)2) hydrotalcite (Mg6Al2(CO3)(OH)16·4(H2O))

20

40

Intensity (cps)

60

with 0.1 M MgCl2

120,000 100,000 80,000 60,000 40,000 20,000 0 400,000

20

40

60

with 1.0 M MgCl2 300,000 Intensity (cps)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47

Page 38 of 38

200,000

100,000

0

20

40

2θ ( ) ACS Paragon Plus Environment o

TOC

60

ACS Earth and Space Chemistry - ACS Publications - American

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