by Clay Minerals at Reservoir Conditions

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Chemical Trapping of CO by Clay Minerals at Reservoir Conditions: Two Mechanisms Observed by In Situ High Pressure and Temperature Experiments Geoffrey M Bowers, John S. Loring, Herbert Todd Schaef, Sydney S. Cunniff, Eric D. Walter, Sarah D. Burton, Randolph K. Larsen, Quin R.S. Miller, Mark E. Bowden, Eugene S. Ilton, and R. James Kirkpatrick ACS Earth Space Chem., Just Accepted Manuscript • DOI: 10.1021/ acsearthspacechem.9b00038 • Publication Date (Web): 19 Apr 2019 Downloaded from http://pubs.acs.org on April 20, 2019

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ACS Earth and Space Chemistry

Chemical Trapping of CO by Clay Minerals at Reservoir Conditions: Two Mechanisms 2

Observed by In Situ High Pressure and Temperature Experiments

Geoffrey M. Bowers*1, John S. Loring2, H. Todd Schaef2, Sydney S. Cunniff1, Eric D. Walter3, Sarah D. Burton3, Randolph K. Larsen IV1, Quin R.S. Miller2, Mark E. Bowden3, Eugene S. Ilton2, R. James Kirkpatrick4 1

Department of Chemistry and Biochemistry, St. Mary’s College of Maryland, St. Mary’s City,

MD, USA, 20686 2

Pacific Northwest National Laboratory, Richland, WA, USA, 99354

3

Environmental Molecular Sciences Laboratory, Pacific Northwest National Laboratory,

Richland, WA, USA, 99354 4

Departments of Chemistry and Earth and Environmental Sciences, Michigan State University,

East Lansing MI, USA, 48824 Corresponding author email: [email protected]

KEYWORDS: hectorite, laponite, carbonation, clay, thin water film, fluid-solid interface.

ACS Paragon Plus Environment

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Abstract This paper presents the results of experiments performed in situ at temperature and pressure relevant to reservoir conditions (T = 323 K and Pfluid = 90 bar) to evaluate whether clay minerals can react with supercritical CO2 to produce carbonate phases by ion exchange-precipitation reactions and dissolution-reprecipitation reactions. The results show that both can occur on a timescale of hours under the conditions of our studies. The dissolution/reprecipitation mechanism was examined using Ca-, Cs-, and tetramethyl ammonium (TMA+) laponite, a synthetic smectite analogous to hectorite that has small particles (basal dimensions ~10 x 10 nm) and a high fraction of edge sites where only two of the usual three bridging oxygen atoms are shared with other tetrahedra in the silicate sheet (Q2 sites), making it an excellent case for examining the role of T-O-T edges. The ion exchange/precipitation mechanism was observed for a Pb-exchanged natural low-Fe smectite (hectorite). Novel X-ray diffraction and NMR and infrared spectroscopic tools provide in situ observation of these reactions in real time supported by a suite of ex situ analyses. The results demonstrate for the first time that

13

C NMR can

effectively characterize the amorphous and crystalline products of such reactions. For all three laponites, IR and NMR data show that HCO3- ion forms at water contents as small as ~5 H2O molecules/exchangeable cation. When the exchangeable cation is Ca2+, the IR data show the formation of carbonate anion at low water content as well, with the NMR spectra showing formation of amorphous calcium carbonate in vacuum-dried samples. For laponites equilibrated at 100% R.H. at atmospheric conditions and then exposed to scCO2,

13

C NMR shows the

presence of a greater number of more mobile HCO3- ions and a poorly crystalline or amorphous hydrous magnesium carbonate/bicarbonate phase that forms from Mg2+ released by clay dissolution. The 100% R.H. sample with exchangeable Ca2+ also forms calcite, vaterite and


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aragonite precipitates. Comparison of these and previously published results suggest that a high edge site Q2 fraction is crucial to the dissolution/reprecipitation process occurring on a short timescale. In the Pb-exchanged hectorite exposed to scCO2, once a critical humidity threshold of ~78% is reached, cerussite (PbCO3) forms rapidly concurrent with replacement of interlayer Pb2+ by H3O+ formed by reaction of CO2 with water on the clay surface. This type of reaction is not observed on a similar timescale with Ca- or Na-exchanged natural hectorite and other smectites, and the low solubility of cerussite appears to be the thermodynamic driving force for this process.

Introduction Clay minerals are major components of shales and other sedimentary rocks and also have many important industrial applications.1 In the context of geological carbon sequestration and enhanced and unconventional petroleum production, there has been significant recent interest in the interaction of CO2 with the expandable clay minerals known as smectites under elevated temperature and pressure conditions relevant to the geological subsurface. The results of experimental and computational molecular modeling studies have shown conclusively that CO2 can intercalate in the interlayer galleries of smectities under these conditions2-27 and that the amount of water associated with the clay (usually characterized by the relative humidity [R.H.] of the CO2, the number of water molecules/exchangeable cation, and the basal spacing of the clay) is critical in controlling CO2 intercalation. In contrast to the results for orthosilicates such as forsterite,28 however, there has been only limited evidence for the reaction of clay minerals with dry or variably wet supercritical CO2 (scCO2). It is commonly thought that the greater


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silicate polymerization of clay minerals and other phyllosilicates makes them less reactive than orthosilicates. Most previous experimental studies of the reaction of phyllosilicate minerals (e.g., smectites, vermiculites, and micas) with CO2 have used H2O-dominated fluids containing CO2 rather than CO2-dominated fluids.29-32 For example, Kang and Roh33 showed evidence of an ion exchange-precipitation mechanism at atmospheric pressure in a H2O-dominated fluid using Ca-exchanged vermiculite and continuous addition of NaCl to cause ion-exchange driven Ca-carbonate precipitation in a basic solution. The high-pressure infrared (IR) and X-ray photoelectron spectroscopic (XPS) study of Hur et al.34 reports formation of CO32- during reaction of a natural smectite (Wyoming montmorillonite) with CO2 at a pressure of 10.45 MPa and temperatures of 295 K and 343 K. They did not characterize the mechanism of reaction or detect neo-formation of any carbonate phase. However, most experimental studies of phyllosilicate interactions with CO2-dominated fluids have shown no evidence of carbonate forming reactions at elevated pressures and temperatures.2-8, 10-12, 26, 27 Loring et al. have recently proposed that even under very low water conditions smectite clay minerals can react to form carbonate/bicarbonate species with variably wet CO2-dominated fluids in two ways.26 One is by replacement of the interlayer cation by hydronium ion (H3O+) produced when CO2 dissolved in the adsorbed water film reacts to produce HCO3- or CO32-, followed by precipitation of the carbonate of the displaced exchangeable cation. The second is by dissolution of the clay and precipitation of carbonate or bicarbonate phases of the dissolved species. Miller et al. have recently demonstrated the reaction of CO2 in thin water films to form carbonic acid35, but as yet there is no clearcut experimental evidence of the ion exchangeprecipitation or dissolution-precipitation mechanisms in smectites exposed to scCO2.


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The experiments described here were carried out using a novel suite of instruments in which the diffraction and spectroscopic data are acquired in situ at elevated temperatures and fluid pressures relevant to the geological subsurface (T = 323 K and PCO2 = 90 bars) while the reactions are proceeding. We investigated the potential for carbonation by the H3O+ exchange mechanism using a Pb-exchanged smectite (hectorite), because PbCO3 has a small solubility product and provides an end-member test example. To evaluate the possibility of dissolution reactions and the effects of different exchangeable cations, we used a synthetic hectorite (laponite) that we expected to be highly reactive, because it has very small particle sizes and a large fraction of Q2 sites. These Q2 sites correspond to silicate tetrahedra that share two rather than three of their oxygens with other tetrahedra in the tetrahedral sheet and are located on the layer edges. They will be referred to as “Q2 sites” or “Q2 broken edge sites” in the remainder of this paper. In both cases, the results presented here show the reactions occur on a timescale of hours under the conditions of these studies, suggesting that such reactions may take place in the pore systems of sedimentary rocks on longer timescales. Importantly, the results show for the first time that 13C NMR spectroscopy can detect and characterize both crystalline and amorphous phases produced by these reactions. Together, the data in this paper show that the nature of the reactions causing carbonation of phyllosilicates depends critically on the water content of the samples and the solubility products of the carbonates of the charge-balancing cations, highlighting the need for greater understanding of the structure, dynamics, and reactivity in thin water films. Smectite clay minerals are of particular value in studies of reactions at variable R.H.s because they have a wide range of compositions and physical and chemical properties, large external surface areas, and interlayer galleries that can contain many different exchangeable inorganic and organic


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cations and fluid species (e.g., H2O, CO2, and CH4). They are also excellent model materials for studying the fundamental behavior of 2-dimensional, slit-type pores bounded by solid phases with a permanent negative charge. Thus, the results address fundamental questions concerning ion transport and chemical reaction in nanoscale water films on surfaces and in confined spaces between them. The results also point the way to future research efforts to unravel the details of carbonation reactions involving not just clay minerals but other porous materials with exchangeable cations such as zeolites and metal organic frameworks. Methods Samples Experiments were performed with two different samples of the trioctahedral smectite clay mineral, hectorite, both of which have a low Fe content (< 0.3%) that makes them excellent materials for studying carbonation reactions using high resolution solid-state NMR. Clay reactivity was studied using a synthetic, high surface area hydroxy hectorite known as laponite. This sample has an average particle diameter parallel to the basal plane of the order of ~10 nm and contains ~10% Q2 Si sites on the broken edges of the T-O-T layers, as demonstrated by 29Si NMR (Figure S1). This particle size is much smaller than typical natural clay minerals, which normally have particle sizes of microns parallel to the basal plane. It is, thus, expected to be very reactive due to its small particle size in the basal dimension, and the consequent high fraction of Q2 sites, and its large surface area. The laponite has a structural formula of M+0.36[Li0.36,Mg5.64]Si8O20(OH)4.36 As with all hectorites, the charge is produced predominantly by isomorphic substitution of Li+ for Mg2+ in the octahedral sheet. Laponite exchanged with Ca2+, Cs+ and tetramethyl ammonium (TMA+: NC4H12+) were made by procedures previously used by our group.3, 37


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The experiments to evaluate the possibility of H3O+ exchange reactions used the natural San Bernardino hectorite (SHCa-1), which is available from the Source Clay Repository of the Clay Mineral Society, that had been exchanged with Pb2+ using similar procedures (see details in the Supporting Information). We have previously shown that the low Fe content of SHCa-1 minimizes line broadening and maximizes the information that can be extracted from NMR spectra.2, 3, 37-40 The sample used is the 60% than at lower R.H.s, and at these higher R.H.s the integrated intensity is continuing to increase at the end of each 6,000 second data collection run. In contrast, for Csand TMA-laponite at R.H.s of ~ 55% and above the integrated intensity increases rapidly up to ~500 – 1,000 seconds and then remains approximately constant. Note that the integrated intensities cannot be quantitatively compared between clay samples. This is because the absorbance was integrated in different spectral regions for different samples, the molar absorptivities of the different clay minerals are likely to be different, and the amount of clay on the ATR crystal is not the same in each experiment. Together, the transmission IR data for water content (Figure 1) and the ATR IR data for the laponite-associated CO32- and HCO3- (Figures 2 and S2) show that formation of carbonate species can occur at very low water contents, as low as ~5 H2O/exchangeable cation in the Csand TMA-samples at R.H.s < 10%. For the Ca-sample, higher H2O contents of at least ~ 50 – 60 H2O/Ca2+ at R.H.s > 60% appear to allow continuous formation of CO32-, whereas for the


11


Cs- and TMA samples, lower H2O contents of ~25 H2O/exchangeable cation or less in this R.H. range appear to allow relatively rapid formation of dominantly HCO3- with its abundance then reaching a steady state. .050

Relative Humidity (%) 0

Predominantly CO Stretches of Carbonates

.040

Absorbance

B 1.4

Ca-Laponite

HOH Bend of H2O

Integrated Absorbance

A

.030

.020

.010

0.000

6

1.2

14 23 31

1.0

39 47

.8

54 60 66

.6

71 76 81

.4 .2 0.0

1725

1625

1525

1425

1325

1225

0

1000

Wavenumber / cm-1

.080

.060

CO Stretch of Bicarbonate

.040

COH Bend of Bicarbonate

.020

1725

5000

6000

Relative Humidity (%) 0 9

.12

19 30 39

.10

47 55 62

.08

67 72

.06

76 80

.04 .02

1625

1525

1425

1325

1225

0

1000

Wavenumber / cm-1

3000

4000

5000

6000

.8

.6

Relative Humidity (%)

.7

0 9

CH3 Deformation

.040

F Integrated Absorbance

TMA-Laponite

HOH Bend of H2O

2000

Time / Seconds

.060 .050

4000

0.00

0.000

E

3000

D .14

Cs-Laponite

HOH Bend of H2O

2000

Time / Seconds


Absorbance

C

Absorbance

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

Page 12 of 37

.030

CO Stretch of Bicarbonate COH Bend of Bicarbonate

.020 .010

20 30 39 48

.5

55 62

.4

67 72 77

.3

81

.2 .1

0.000

0.0 1725

1625

1525

1425

1325

Wavenumber / cm-1

1225

0

1000

2000

3000

4000

5000

6000

Time / Seconds


12

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Figure 2. ATR IR spectra from the in situ IR spectroscopic titration of Ca-, Cs-, and TMAlaponite (A,C,E, respectively) with water at T = 323K and Pfluid = 90 bars in the HOH bending region of sorbed H2O, the C-O stretching region of carbonate and bicarbonate, and the COH bending region of bicarbonate as a function of R.H. Spectra were collected at ~4.5 minute intervals for a total of ~100 minutes at each R.H., and the spectra shown here are the last spectra after ~100 minutes of reaction. Panels B, D, and F are the integrated absorbance in the wavenumber regions 1330 to 1560 cm-1, 1230 to 1475 cm-1, and 1321 to 1383 cm-1 during the reaction of Ca-, Cs-, and TMA-laponite, respectively. Values are shown as a function of time at ~4.5 minute intervals at each R.H. value. The background spectrum for these data is the unreacted hectorite at 0% R.H. These plots demonstrate that carbonation rates for all three clay minerals dramatically increased beyond a R.H. value of ~55%. The spectra show that under these experimental conditions CO32- is the dominant carbonate species for Ca-laponite, whereas HCO3- is the dominant carbonate species for Cs- and TMAlaponite. The band near 1489 cm-1 is deformation the CH3 groups of TMA.46 CO2 - Laponite Interactions: 13C NMR Results The in situ 13C MAS NMR spectra of all the laponite samples in contact with scCO2 at 323 K and 90 bar are dominated by signal for bulk scCO2, but all also contain signal for HCO3-, and some contain signal for CO32-. For both the vacuum dried samples and those exposed to 100% R.H., the signal is dominated by a central peak for 13CO2 at 125.3 ± 0.1 ppm and two associated spinning sidebands (SSBs; Figure 3; Table 1). These features are very similar to those of pure scCO2 interacting with an otherwise empty rotor.3 For the vacuum dried samples, 97% of the total

13

C integrated intensity is associated with the CO2 molecules, while at 100% R.H. this

central peak and pair of sidebands accounts for 93.0% of the integrated intensity. The spectra all lack the broad, asymmetric spinning sideband (SSB) pattern indicative of CO2 confined in smectite interlayers.3 Thus, the NMR data are clear that bulk scCO2 experiences only brief interactions with the external surfaces of the laponite at both humidities. Table 1. Relative intensities (%; ±0.1%) of CO2, HCO3-, and laponite samples from 13C MAS NMR. Cation

TMA Cs

Equilibration RH CO2

vac vac

96.7 96.8

HCO3-, free 3.3 3.2

HCO3-/CO32-, amorph Mg(H)CO3

ACC

-


-

13

C carbonate minerals in

calcite vaterite

-

-

aragonite

-

13


Ca TMA Cs Ca Ca

vac 100% 100% 100% 100%, day 2

96.2 93.0 93.2 95.5 94.7

2.3 5.9 5.7 2.1 3.0

1.1 1.1 1.5 1.2

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1.5 -

0.3 0.4

0.2 0.3

0.4 0.5

Figure 3. 13C MAS NMR spectra of Ca-Laponite initially exposed to 100% R.H. (A) or vacuum dried (B). The blue region represents the HCO3- chemical shift range, and the red region the CO32- chemical shift range. Both (A) and (B) show an unexpanded spectrum dominated by the narrow resonance for scCO2 near 125 ppm, and a vertically expanded spectrum showing the resonances for HCO3- and CO32-. Stars mark the positions of spinning sidebands. All the vacuum dried and 100% R.H. samples yield a resonance between 160.7 and 161.6 ± 0.1 ppm representing HCO3-.47 For the vacuum-dried Cs- and TMA-laponite samples this resonance occurs at 160.7 ± 0.1 ppm, is quite broad (full width at half height [FWHH] = 4.9 ± 0.1 ppm) and contains ~3% of the total 13C signal (Figure 4, Table 1). For the vacuum-dried


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Ca-laponite this resonance occurs at 161.6 ± 0.1 ppm, is also quite broad (FWHH = 6.3 ± 0.1 ppm) and contains 2.3 ± 0.1% of the total signal intensity. For the 100% R.H. samples, the resonances for HCO3- occur at essentially the same chemical shift but narrow dramatically and increase in intensity. For the Cs- and TMA-laponite samples, the line widths decrease to 0.83 ± 0.1 ppm, and their relative intensities increase to 5.8 ± 0.1%, nearly twice the integrated intensity of the vacuum-dried samples (Figure 4, Table 1). For the 100% R.H. Ca-laponite the width of the resonance for HCO3- decreases to 1.15 ppm. Its relative intensity is 2.1% for the spectrum acquired immediately after exposure to scCO2 and 3.0% for the spectrum acquired a day later (Figure S3, Table 1), demonstrating that HCO3- continues to form over at least the first 24 hours while the samples are in contact with wet scCO2 in the NMR rotor.

Figure 4. 13C MAS NMR spectra of the carbonate/bicarbonate region of the Ca-, Cs-, and TMA-exchanged laponite samples initially vacuum dried or exposed to 100% R.H. at room pressure and temperature. The spectra were acquired at T = 323 K and Pfluid = 90 bar. The peak heights of the Ca-laponite are doubled relative to the others. The asterisks mark the


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location of the spinning sidebands for bulk scCO2, which are well-resolved from the peaks for HCO3- and CO32- at the MAS frequencies used. All three laponites exposed to 100% R.H. also yield a relatively broad peak centered near 165.7 ± 0.1 ppm that contains 1.1 to 1.5% of the total 13C integrated intensity (Figures 4 and S3; Table 1). Moore et al. have shown that the 13C chemical shifts of hydrated Mg-carbonate and bicarbonate phases occur near this value (nesquehonite [Mg(OH)(HCO3)·2H2O)] at 165.4 ppm, hydromagnesite

[4MgCO3·Mg(OH)2·4H2O)]

at

165.2

ppm,

and

dypingite

[4MgCO3·Mg(OH)2·5-8H2O] at 165.3 ppm).48 However, the widths of the resonances (FWHH = 2.0 ± 0.1 ppm) are much greater than for the crystalline phases studied by Moore et al. Thus, we assign this resonance to an amorphous Mg-carbonate/bicarbonate precipitate that is probably hydrous.49 In our experiments Mg2+ would be available to form Mg-carbonate/bicarbonate phases only if the laponite dissolves, since the only source of Mg2+ is the octahedral sheet of the clay.

As for the HCO3- resonance, the intensity of the resonance for amorphous Mg-

carbonate/bicarbonate in the reacting Ca-laponite increases with time as the sample is exposed to wet scCO2 in the NMR rotor (Figure S3). This result is consistent with continuing dissolution of the laponite structure occurring over at least 24 hours under the conditions of the NMR experiments. For the Ca-laponite, the 13C NMR spectra also contain resonances for Ca-carbonates, with the specific phases and abundances depending on the humidity of the fluid phase and thus the amount of adsorbed H2O. For the vacuum dried Ca-laponite, the CO32- resonance occurs at 168.3 ± 0.1 ppm, is quite broad (FWHH = 3.3 ± 0.1 ppm), and contains 1.5 ± 0.1% of the total intensity. The position and breadth of this peak are similar to those reported in the literature for amorphous calcium carbonate (ACC)50, 51, and we assign it to ACC as well. The 100% R.H. Ca-laponite


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sample yields three narrower resonances at 168.4 ± 0.1 ppm, 169.2 ± 0.1 ppm, and 170.8 ± 0.1 ppm (FWHH ~ 0.37 ppm for all three peaks) that together contain 0.9% of the signal in the spectrum acquired immediately after exposure to scCO2 and 1.2% of the signal intensity after 1 day. These peaks are readily assigned to the crystalline Ca-carbonate phases; calcite (168.4 ppm), vaterite (169.2 ppm) and aragonite (170.8 ppm).52 Simultaneous formation of these three phases has been previously observed in nanofluidic experiments involving reaction of Ca- and Mg-containing solutions at room conditions53, conditions in many ways similar to our experiments if the clay layers are dissolving and releasing Mg2+. The 13C NMR results are, thus, in good agreement with the ATR-IR data described above, showing the formation of both CO32- and HCO3- with the amounts and structural environments of both species dependent on the charge-balancing cation and the humidity. Both data sets show formation of CO32- in Ca-laponite and HCO3- in the Cs- and TMA laponites. The NMR, however, shows substantial amounts of HCO3- for Ca-laponite at both R.H.s, whereas the IR appears to show less. This difference may be due to overlap of the HCO3- signal and CO32- signal in the IR spectra, suggesting that the quantities detected by NMR and IR are more similar than they first appear. The difference may also be related to the different physical forms of the samples in the different experiments (bulk powder in the NMR and a thin film in the IR) or to differences in how H2O was introduced to the samples (the NMR experiments expose wet systems to dry scCO2 while the IR exposes dry samples to wet scCO2). We also note that the ATR-IR data show the presence of carbonate species under vacuum and before exposure to the scCO2. This carbonate must be present in the initial samples, but it is not detectable in the 13C NMR spectra because the natural abundance of 13C is only about 1.1%. Thus, all the observed 13C NMR signal for the carbonate and bicarbonate species must represent C originally in the 13C-enriched scCO2.


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The ambient condition XRD patterns of the laponite samples before and after reaction with H2O-saturated CO2 show the presence of only the laponite with no evidence of crystalline Caor Mg-carbonates or a silica phase (Figure S4). This observation is consistent with the proposed amorphous structure of some of the reaction products observed by 13C NMR and their overall low abundance. The ATR-IR and 29Si MAS NMR spectra show evidence of amorphous silica in the reacted samples, consistent with the idea of laponite dissolution. The ATR-IR data for Ca-laponite reacted with scCO2 at 81% R.H. clearly shows formation of amorphous silica (Figure S5). Similarly, the 29Si NMR spectrum of the Ca-laponite contains some signal above noise in the 110 ppm range that represents hydrous amorphous silica. The 29Si NMR spectra of the Cs- and TMA-laponites show no evidence for a silica phase produced by incongruent dissolution of the laponite (Figure S1). CO2 - Pb-Hectorite Interactions The HAADF STEM images of the Pb-exchanged SHCa-1 hectorite show redistribution of the Pb ions after exposure to 85% R.H. scCO at T = 323 K, P 2+

2

= 90 bar (Figure 5)

fluid

demonstrating that the vast majority of the initially charge-balancing Pb is converted to a Pb2+

rich reaction product. Before exposure to wet scCO , the charge-balancing Pb in the interlayer 2+

2

galleries is visible as the bright, parallel linear features on the left side of the clay particle in Figure 5a. In this image, the clay layers are oriented essentially perpendicular to the image plane, and thus the view is parallel to the basal surfaces. Images of the unreacted sample taken with the clay layers essentially parallel to the image plane shows a few ~2-3 nm Pb-rich domains and individual Pb ions which appear as isolated white dots (Figure 5b). The mottled appearance 2+

of this image is due to the wavy nature of the clay impacting the TEM focus for individual ions.


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Powder XRD (Figure 6) and TGA-MS (Figure S6) do not detect a carbonate phase in this unreacted sample, indicating that it is a minor impurity of the type expected by exposure to air and that it may be poorly crystalline, in agreement with the TEM images. For the sample exposed to 85% R. H. scCO at 323 K and 90 bar in the IR titration experiments discussed below, 2

the HAADF-STEM images do not show the bright lines indicative of interlayer Pb or the bright 2+

points indicative of individual Pb ions on the basal surface (compare Figures 5a to 5c). Instead, 2+

there are larger (100-200 nm) Pb-rich particles (Figure 5d) that HeIM shows are on the hectorite particle surfaces (Figure S7).

Figure 5. HAADF-STEM images of the Pb-SHCa-1 before (top row) and after (bottom row) exposure to 85% R.H. scCO2 at T = 323 K and Pfluid = 90 bar. The bright white lines in Figure 5a show the positions of interlayers containing Pb2+. The white points in Figure 5b are individual Pb2+ ions, and the bright ~2-3 nm patches in Figure 5b are nano-particles of a Pbrich phase, probably a Pb-carbonate. After exposure to 85% R.H. scCO2 the interlayer Pb2+ and individual Pb2+ is absent (Figure 5c), and larger ~100-150 nm Pb-rich domains are visible (Figure 5d). XRD and 13C NMR show these domains to be cerussite. The in situ XRD results at 323 K and 90 bars show the expected basal spacing expansion of smectite with increasing R.H. in scCO and the presence of cerussite (crystalline PbCO ) at R.H.s 3 2


3

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of 78.5% and higher (Figures 6 and S8). Exposure to anhydrous CO increases the basal spacing 2

of the clay (d001) from 10.45 Å to 11.96 Å, indicating that CO is intercalated into the interlayers 2

and that it plays an active role in expansion of Pb-exchanged smectite, as shown in previous reports.

3, 4, 6-8

Increasing R.H. of the scCO fluid causes progressively increasing d001 values, with 2

a maximum of 16.54 Å at ~100% R.H., showing intercalation of H O into the smectite interlayers 2

and development of a bilayer (2WL) hydration state. Cerussite is first detected by XRD at 54

55

~78.6% RH, and its diffraction intensity increases at higher R.H.s (Figure S8). When the XRD sample is returned to the vacuum condition, the basal spacing is 11.78 Å rather than the expected value of 10.45 Å, indicating ion exchange and/or retention of interlayer H O. Diffuse reflectance 2

IR spectra of unreacted and 85% R.H. scCO reacted Pb-SHCa-1 taken under 10 Torr vacuum -4

2

(Figure S9) show much more water in the reacted sample, consistent with the XRD results. These spectra also show greatly increased CO consistent with cerussite formation. Based on 23

the XRD and IR observations, we assign the Pb-rich particles in the HAADF-STEM images of the samples exposed to wet scCO to cerussite. This result clearly demonstrates reaction of the 2

Pb with the only source of CO , the scCO phase. Indeed, the TGA-MS data show the presence 2+

23

2

of ~12.3 wt.% cerussite decomposing at temperatures between 260 and 300°C (Figure S6). The mass changes observed by TGA combined with the hectorite and cerussite compositions show essentially complete conversion of the Pb in the sample to cerussite. 2+


20

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Figure 6. In situ powder X-ray diffraction patterns of Pb-SHCa-1 acquired after exposure to vacuum, dry scCO2, and scCO2 at 78% and 100% R.H. at T = 323 K and Pfluid = 90 bar. Reflections diagnostic to cerussite (PbCO3) appear at ~74% R.H. and grow in intensity with increasing R.H. The R.H. dependence of the in situ ATR-IR spectra of the Pb-SHCa-1 sample show two critical humidities associated with CO formation. Increase in the amount CO ion begins at 23

23

R.H.s as low as 3% (Figures 7 and S2), and the rate of CO formation increases above ~37% 23

R.H. The reaction did not reach steady state at higher R.H.s, even after 6,000 seconds of reaction. This behavior is generally similar to that of Ca-laponite (Figure 2). Unfortunately, the Pb-SHCa1 started to peel away from the ATR optics at higher RH, making the results at R.H.s > 85% unreliable.


21


.050

.040

Pb-SHCa-1

B 1.4

CO Stretch of Carbonate

Relative Humidity (%) 0 3

HOH Bend of H2O


A Absorbance

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

Page 22 of 37

.030

.020

.010

1.2

7 12 17

1.0

23 28 32

.8

38 42 46

.6

51 54

.4

58 62

.2 0.000

0.0 1725

1625

1525

1425

1325

1225

0

500

Wavenumber / cm-1

1000

1500

2000

2500

3000

Time / Seconds

Figure 7. A. ATR IR spectra in the HOH bending region of sorbed H2O and the asymmetric C-O stretching region of CO32- as a function of R.H. from the in situ IR spectroscopic titration of Pb-SHCa-1 with water at T = 323K and Pfluid = 90 bars. Spectra were collected at ~4.5 minute intervals for a total of ~45 minutes at each R.H., and the spectra shown here are the last spectra after ~45 minutes of reaction. B. The integrated absorbance of the asymmetric C-O stretching region of CO32- (1218 to 1515 cm-1) as a function of time at ~4.5 minute intervals at each R.H. value. The background spectrum for these data is the unreacted hectorite at 0% R.H. This plot demonstrates that carbonation dramatically increased beyond a R.H. value of ~46%. H, C, and Si MAS NMR spectra of the reacted and unreacted Pb-hectorite also demonstrate

1

13

29

the presence of cerussite in the reacted sample and provide insight into the mechanism of cerussite formation. The C MAS NMR peak in the reacted Pb-SHCa-1 sample is at the same 13

chemical shift as cerussite (Figure S10), although its width is larger, indicating greater structural disorder for the small particles formed in the clay. The Si MAS NMR spectra of the Pb29

exchanged hectorite acquired before and after exposure to 85% R.H. scCO at 323 K and 90 bars 2

are essentially identical (Figure S11), indicating no detectable dissolution of the hectorite T-OT structure on this timescale. This behavior contrasts with the observed dissolution of hectorite in acid digestion experiments at ambient conditions and with the dissolution of the laponite 56

discussed above. The H MAS NMR spectra of the reacted and unreacted Pb-hectorite samples 1


22

Page 23 of 37 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


are also essentially identical except for a resonance at 3.6 ppm representing H O in the reacted 3

+

sample (Figure 8). Previous molecular modeling studies show that the H chemical shift of H O 1

+

3

approaches that of interlayer adsorbed H O (~3.4 ppm) as the number of water molecules 2

hydrating it increases, and that tetrahydrated hydronium (H O ·4H O) has a H chemical shift of +

3

3.6 ppm, identical to that for the reacted sample here.

57, 58

1

2

Thus, we assign this resonance to H O

+

3

balancing the negative structural charge of the clay after exchanging onto it during removal of Pb to form cerussite. 2+

Figure 8. Ex situ 1H MAS NMR spectra of unreacted Pb-SHCa-1 (black), Pb-SHCa-1 exposed to scCO2 at 85% R.H. at T = 323 K and Pfluid = 90 bar in the I.R. titration experiments for 30 hours (red), and the difference spectrum with intensity increased by a factor of ~3 (blue). The peak at 3.6 ppm is diagnostic of the hydrated H3O+ ion.57 The peak near 1 ppm is readily assignable to structural OH- of the octahedral sheet.59 The peak at 5 ppm is probably 1H background from the probe head and some signal from adsorbed H2O. DISCUSSION Carbonate Speciation and the Critical Role of H O 2

The XRD, IR, and NMR results show that reactions involving scCO and the clay minerals 2

studied here must involve formation of HCO and CO from the scCO phase and that this can 3

23

2

occur in the presence of even small amounts of water. Formation of HCO and CO under these 3


23

23


Page 24 of 37

conditions inherently involves H O and the formation of H O , as described by the standard 2

3

+

reaction scheme: CO (sc) Û CO (aq) 2

(A)

2

CO (aq) + H O(l) ® H CO (aq) 2

2

2

(B)

3

H CO (aq) + H O(l) Û H O (aq) + HCO (aq)

(C)

HCO (aq) + H O(l) Û H O (aq) + CO (aq)

(D)

2

3

2

3

3

+

2

3

3

+

23

The XRD, IR, and NMR data for the laponites all show increasing amounts of HCO and CO at 3

23

high R.H.s, consistent with increased solvation of CO at higher water contents driving these 2

reactions toward their products. Likewise, the XRD, IR, and NMR data for the Pb-SHCa-1 show that these reactions can be critically dependent on R.H., with cerussite formation occurring only when the R.H. is high enough to cause formation of the bilayer hydrated state. This may be a kinetic effect, with the bilayer structure necessary to allow diffusion of Pb out of the interlayer 2+

galleries and H O into them at a rate detectable on the time scale of these experiments. 3

+

At atmospheric temperature and pressure in the absence of cations, reactions (A)-(D) lead to a slightly acidic aqueous solution with HCO as the dominant anion. At elevated temperatures 3

the K values for these reactions are displaced to lower pHs, but to our knowledge the equilibrium a

constants are not known at the temperature and pressure used here. Similarly, the effects of the amount of water on the surface of a clay (water film thickness) on the equilibrium constants are not known, and the (aq) designation above does not necessarily imply the presence of bulk water. Indeed HCO and CO form in the vacuum dried laponites in the presence of very thin water 3

23

films with a total of as little as 5 to 12 H O molecules/exchangeable cation. Note that this H O 2

2

includes that in the interlayer galleries and on the external basal and edge surfaces, and that it is not possible to differentiate between H O in these different structural environments with the 2


24

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current data set. We report the total H O content, because it is well characterized and to facilitate 2

comparison with our previously published data reported on this basis. The formation of solid phases containing CO drive the equilibria of reactions (C) and (D) 23

to the right for both laponite and Pb-SHCa-1. For example, cerussite formation in the SHCa-1 sample causes an ion exchange reaction that drives increased consumption of CO via reactions 23

(E) and (F). Hec-Pb + 2H O (aq) ® Hec2-H O + Pb (aq)

(E)

Pb (aq) + CO (aq) ® PbCO (s)

(F)

2+

3

2+

+

+

3

23

2+

3

The consumption of CO by cerussite formation shifts the equilibria of steps A-D towards 23

production of additional CO , particularly since CO is in excess in our experiments. The ion 23

2

exchange reaction ensures that additional Pb becomes available to consume additional CO 2+

23

(aq), which in turn consumes more scCO from the bulk fluid. For a bulk water phase, elevated 2

pressures and temperatures will favor much larger dissolved CO concentrations in Step A than 2

at atmospheric conditions (~1.1 M dissolved CO at 90 bar P and 323 K), but the solubility in 60

2

CO2

a thin water film is not known. The effectiveness of carbonate phase precipitation at driving reactions (A) – (D) via reaction (H) in the absence of ion exchange and (G) and (H) when ion exchange occurs is driven Hec-M + nH O (aq) ® Hec-nH O + M (aq)

(G)

M (aq) + n/2 CO (aq) ® M(CO ) (s)

(H)

n+

+

+

3

n+

3

23

3 n/2

n+

largely by the solubility product constant of the metal carbonate precipitate (in reactions G and H, M represents an exchangeable cation). Cerussite has a solubility product of 7.40x10 at n+

-14

25°C , roughly five orders of magnitude lower than the calcium carbonates. Thus, in addition to 61

driving Step (D) to the right, the low solubility of cerussite also keeps the activity of Pb (aq) 2+


25


Page 26 of 37

near the growing crystals very low, maintaining a chemical potential gradient to drive Pb

2+

diffusion out of the interlayers and resulting in a final product consisting of it and Hec-2H O

+

3

over a short reaction period. Once again, we also note that the R.H. must be high enough to form a 2-layer hydrate in the clay interlayer for this process to be rapid. For the laponites, even the presence of a small amount of water allows solvation of CO2 to proceed at least through reaction (C), forming HCO3- in all the samples irrespective of the charge balancing cation. We note, however, that for clay minerals these reactions might involve broken edge sites on the T-O-T layers as, for instance, H+ acceptors or donors. For the vacuum dried samples, in which observable dissolution of the clay structure and precipitation of hydrous Mgcarbonates did not occur, only the Ca2+ in Ca-laponite may have driven reactions (A) – (D) sufficiently to the right to produce a carbonate precipitate (ACC). The solubility products of even crystalline Ca-carbonates at ambient conditions are significantly larger than Pb-carbonate (calcite, 3.36x10-9; aragonite, 6.0x10-9) despite the similar hydration energies for Pb2+ and Ca2+ (1481 kJ/mol vs. 1577 kJ/mol, respectively),61 but the Ksp may be sufficient to allow Ca2+/H3O+ exchange reactions in this type of material. Thus, the absence of detectable precipitation of Cacarbonate from Ca-exchanged SHCa-1 hectorite or other smectites in experiments performed under the same conditions as those here2, 3, 7 may be due to the reduced thermodynamic driving force for secondary carbonate phase precipitation or to slower kinetics than the Pb2+ case due to a smaller chemical potential gradient driving diffusion and/or the larger particle sizes of these samples compared to laponite. All three would reduce the ability to experimentally observe any edge-site effects. Mg-carbonates have larger solubility products than Ca-carbonate (MgCO , 3

6.82x10 ; MgCO ·3H O, 2.38x10 ) , but because all the Mg in the laponite samples is -6

-6 61

3

2+

2

structurally held in the octahedral sheet, our results cannot be used to determine whether such


26

Page 27 of 37 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


phases would form under our experimental conditions with Mg as the exchangeable cation. The 2+

presence of exchangeable Cs and TMA in the Cs- and TMA-laponites did not drive reactions +

+

(A)-(D) and (G)-(H) sufficiently to the right to form CO in the absence of dissolution of the 23

clay. The solubility products of Cs- and TMA-carbonate are unknown, but they are surely even larger than those of the Mg-carbonates. For instance, of the alkali carbonates Li CO has a value 2

3

of 8.15x10 . Thus, formation of carbonate phases by simple ion exchange with more common -4 61

exchangeable alkali cations such as Na+ and K+ is not likely to occur. We note also that the amount of water in the sample can have a significant effect on the nucleation of carbonate phases. In the vacuum dried Ca-laponite only ACC formed, whereas in the 100% R.H. sample the crystalline phases calcite, vaterite and aragonite formed. This precipitation was not related to drying but occurred in situ in the pressurized NMR rotor and seems to be linked with dissolution of the clay structure, as discussed in the next section. Reaction of Clay Layers The small particle sizes of the laponites and the large number of potentially reactive Q2 edge sites compared to natural smectites and other clay minerals62 make this material an excellent model material to study the role that phyllosilicate edge site reactivity may play in carbonateforming reactions involving CO2. For the 100% R.H. laponite samples here, the clay T-O-T layers are clearly involved in the reactions, as demonstrated by the presence of Mgcarbonate/bicarbonate precipitates irrespective of the exchangeable cation, their growth over time while exposed to wet scCO (Figure 4), and the presence of amorphous silica as shown by 2

the IR and supported by the Si NMR results (Figures S1 and S5). As discussed above, the only 29

possible source of Mg for these precipitates is dissolution of the octahedral sheet of the clay. 2+


27


Page 28 of 37

Detectable dissolution does not occur under our experimental conditions for smectites with larger particle sizes, and the edge sites, thus, appear to be critical in this regard. Under acidic conditions, dissolution of silicates is thought to occur by electrophilic attack on the surface sites by H3O+.63 Such acid attack on hectorite is well known to occur at room temperature. However, to the best of our knowledge, no studies have examined the potential 56

effects of different protonation states of the oxygen sites (-O, -OH, or -OH2) on broken edges of phyllosilicate T-O-T layers on this attack. Similarly, possible effects of chemisorbed CO32- or HCO3- exchanged onto such sites is also not understood. For the vacuum dried laponite samples here, the lack of IR bands for amorphous silica show that acid attack was not observed at this R.H. Thus, it may not be necessary for formation of CO32-, since the Ca2+ for the ACC could have become available by ion exchange with H3O+, as discussed in the previous section. Regardless of whether pre-existing calcium carbonate or acid attack of the clay layers is involved in the formation of the observed ACC, mobilization of charge-balancing Ca is critically 2+

dependent on the presence and potentially the thickness of the adsorbed H O film. 2

The amount of surface-sorbed H2O thus appears to be an important factor in the dissolution reactions for the laponite samples. For phyllosilicates more generally, the fraction of broken edge sites on the T-O-T layers also appears to be a critical factor for the dissolution mechanism proposed by Loring et al.26 For example, the T-O-T layers of Pb-SHCa-1 are not detectably involved in dissolution reactions, consistent with the much lower Q2 fraction that minimize the contributions from broken edge sites; the fraction of broken edge sites may be so low that their effects cannot be detected with these techniques on the timescales of our experiments (Figure S11). Broader Implications


28

Page 29 of 37 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


The results of the experiments described here suggest that substantial amounts of carbonation reactions involving clay minerals are unlikely to occur on short (day to week) timescales in oil and gas reservoirs exposed to CO -dominated fluids but that such reactions may take place on 2

longer time scales. As discussed above, previous in situ laboratory experiments have not detected carbonation of Ca-, Na-, K-, NH - or Cs-exchanged montmorillonite (a dioctahedral 4

smectite) or hectorite under temperature, pressure, R.H., and time conditions similar to those here.

2-8, 10, 11, 26

Comparison of the results for Pb-hectorite and the laponite samples suggests that rapid

carbonate precipitation involving cation exchange with the H O produced in the formation of +

3

HCO and CO reactions requires a low K for the carbonate phase. Pb is not a geochemically 3

2+

23

sp

abundant charge balancing cation in the clay minerals of oil and gas reservoirs, and neither are other cations that form carbonates with low K , such as Zn, Fe(II) or Mn(II) (ZnCO ·H O, FeCO , sp

3

2

3

and MnCO have 298K K values between those of calcite and cerussite, approximately two 3

sp

orders of magnitude less than calcite on average ). However, such reactions may be possible 61

over much larger timescales, and additional, long-term studies will be needed evaluate this possibility. The net effect of these precipitates on the pore network in a non-conventional hydrocarbon reservoir is also unclear, since the carbonate phases have the potential to clog pores or induce additional fracturing as the precipitate grows.

64

Reactions involving the T-O-T structure of clay minerals and other phyllosilicates (e.g., smectite, illite, mixed layer illite-smectite, kaolinite, micas, or chlorite) may also occur on longer time scales, particularly in systems that contain greater amounts of H O. The observed 2

dissolution and reprecipitation reactions involving laponite on a 24 hour or shorter timescale here are clearly due to its small particle size and high edge site concentration allowing rapid attack of the T-O-T layer edge sites, especially at high fluid R.H.s or significant surface-


29


Page 30 of 37

adsorbed H O concentrations. Although such attack on natural clay samples with larger particle 2

sizes has not been observed on the laboratory time scale under our conditions, the laponite results suggest that it could occur on the multi-year lifetime of a petroleum reservoir or the even longer lifetime of a geological CO sequestration site. Similar dissolution/precipitation reactions would 2

result in the neoformation of dominantly Ca-, Mg-, and Fe-carbonates and possibly hydrous carbonates and bicarbonates. Note that Fe(II) is also an abundant structural species in clay minerals and Fe-oxyhydroxide phases and could be involved in precipitation of Fe-carbonates. The silica released during attack on the clay minerals would probably end up in quartz, and the aluminum in kaolinite, chlorite or more Al-rich illitic clay. The presence of amorphous calcium carbonate and magnesium carbonate/bicarbonate in the laponite samples suggests that precipitation of amorphous carbonate phases may be an important step in carbonate forming reactions in water-poor environments exposed to scCO . Their 2

formation and recrystallization and the subsequent growth of the crystalline phases will be a fruitful direction for further research and may provide new insight into the fundamental chemical transformation mechanisms of amorphous carbonates , particularly those that occur on mineral 65

surfaces. To the best of our knowledge, there have been no detailed, molecular-scale studies of the behavior of any amorphous carbonate at elevated temperature and P , perhaps because of CO2

the challenge of performing in situ time resolved spectroscopy and microscopy. Another general conclusion of this study is that ion exchange reactions in smectites can occur in the absence of bulk water as long as there is a sufficient thermodynamic driving force. The experimental observation of Pb exchange with H O occurs because the low solubility of Pb2+

+

3

carbonate and the subsequent concentration gradient that drives diffusive transport, at least once a bilayer of adsorbed H O is present. The observed ion exchange under these conditions suggests 2


30

Page 31 of 37 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


that such exchange could occur in thin water films driven by other thermodynamic factors. For example, scCO -induced dissolution of calcite in shales could produce a large enough 2

concentration gradient in the water films on the clay minerals to drive replacement of Na or K +

+

by Ca in the smectite component of the illite-smectite phase present in many shales. Such 2+

exchange could affect the intercalation behavior of CO and H O or perhaps the distribution of 3 2

2

pore sizes. Both mechanisms could impact porosity and permeability. Our group is currently exploring such ion exchange processes in systems containing smectite carbonate phases with thin water films. Beyond potential geochemical implications, this study also reveals a new mechanism to produce catalytically important, synthetic hydronium smectite that does not involve exposure to strongly acidic conditions. Recent work by Motokura and colleagues has shown that H O 3

+

smectites can be used to promote a variety of reactions involving simple alkenes, including allylsilylation, nucleophilic addition of sulfonamides and carboxamides, substitution of alcohols by amides and anilines, and nucleophilic addition of 1,3-dicarbonyl compounds.

66-68

However,

most published procedures for making H O -smectites involve high water/solid ratios at very low 3

+

pHs (pH < 2). Such conditions are known to damage the clay framework, potentially altering 56

the clay properties and their catalytic activity in an uncontrollable way. The synthesis of H O 3

+

smectite by reaction of the Pb -exchanged phase with variably hydrated scCO does not 2+

2

detectably attack the T-O-T structure of the clay (Figure S11). Although H O -smectite made by +

3

current methods do catalyze high yields in the chemical reactions listed above, study of the catalytic activity of these materials produced via a scCO -driven mechanism seems warranted. 2

Supporting Information


31


Page 32 of 37

Details of experimental methods and analysis; Si NMR of reacted laponite samples; ATR-IR 29

of the carbonate band region intensity as a function of R.H.; C NMR of Ca-laponite as a function 13

of time; ex situ XRD patterns of reacted laponites; ATR-IR spectra of reacted laponites; TGAMS of reacted vs. unreacted Pb-hectorite; helium ion microscope images of reacted vs. unreacted Pb-hectorite; in situ XRD patterns of Pb-hectorite showing cerussite growth; diffuse reflectance IR of reacted vs. unreacted Pb-hectorite; C MAS NMR of cerussite vs. reacted Pb-hectorite; Si 13

29

NMR of Pb-hectorite Acknowledgements This paper is based on work supported by the U.S. Department of Energy (DOE), Office of Science, Office of Basic Energy Sciences (BES), Chemical Sciences, Geosciences, and Biosciences Division through its Geosciences program at Pacific Northwest National Laboratory (PNNL) (JSL, ESI, QRSM), grant DE-FG02-08ER15929 at Michigan State University (RJK, P.I.), and a subaward from that grant to St. Mary’s College of Maryland (GMB, P.I.). HTS gratefully acknowledges support from the DOE Office of Fossil Energy at PNNL through the National Energy Technology Laboratory, Morgantown, West Virginia. GMB was also supported through the U.S. DOE, Office of Science, Office of Workforce Development for Teachers and Scientists (WDTS) under the PNNL Visiting Faculty Program (VFP) and SSC through the PNNL Science Undergraduate Laboratory Internship Program (SULI) operated out of the same DOE Office and Program. NMR, HeIM, and TEM data were collected at the Environmental Molecular Science Laboratory (EMSL), a DOE Office of Science User Facility sponsored by the Office of Biological and Environmental Research and located at PNNL. RJK and GMB also acknowledge support from the Michigan State University Foundation. Many thanks to Libor Kovarik (TEM) and Bruce Arey (HeIM) of PNNL for acquiring the microscopy


32

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data for this study. We also thank Katherine Huerta (SMCM) for her efforts to quantify the extent of Pb exchange in the Pb-SHCa-1 experiments. 2+

References 1. Ilgen, A. G.; Heath, J. E.; Akkutlu, I. Y.; Bryndzia, L. T.; Cole, D. R.; Kneafsey, T. J.; Milliken, K. L.; Pyrak-Nolte, L. J.; Suarez-Rivera, R., Shales at all scales: Exploring coupled processes in mudrocks. Earth-Science Reviews 2017, 166, 132-152. 2. Bowers, G. M.; Hoyt, D. W.; Burton, S. D.; Ferguson, B. O.; Varga, T.; Kirkpatrick, R. J., In situ C and Na magic angle spinning NMR investigation of supercritical CO incorporation in smectite-natural organic matter composites. J. Phys. Chem. C 2014, 118 (7), 3564-3573. 3. Bowers, G. M.; Schaef, H. T.; Loring, J. S.; Hoyt, D. W.; Burton, S. D.; Walter, E. D.; Kirkpatrick, R. J., Role of cations in CO adsorption, dynamics, and hydration in smectite clays under in situ supercritical CO conditions. J. Phys. Chem. C 2017, 121 (1), 577-592. 4. Loring, J. S.; Ilton, E. S.; Chen, J.; Thompson, C. J.; Martin, P. F.; Benezeth, P.; Rosso, K. M.; Felmy, A. R.; Schaef, H. T., In situ study of CO and H O partitioning between Namontmorillonite and variably wet supercritical carbon dioxide. Langmuir 2014, 30 (21), 61206128. 5. Loring, J. S.; Schaef, H. T.; Thompson, C. J.; Turcu, R. V.; Miller, Q. R.; Chen, J.; Hu, J.; Hoyt, D. W.; Martin, P. F.; Ilton, E. S.; Felmy, A. R.; Rosso, K. M., Clay hydration/dehydration in dry to water-saturated supercritical CO : Implications for caprock integrity. Energy Procedia 2013, 37, 5443-5448. 6. Loring, J. S.; Schaef, H. T.; Turcu, R. V. F.; Thompson, C. J.; Miller, Q. R. S.; Martin, P. F.; Hu, J.; Hoyt, D. W.; Qafoku, O.; Iton, E. S.; Felmy, A. R.; Rosso, K. M., In situ molecular spectroscopic evidence for CO intercalation into montmorillonite in supercritical carbon dioxide. Langmuir 2012, 28 (18), 7125-7128. 7. Schaef, H. T.; Loganathan, N.; Bowers, G. M.; Kirkpatrick, R. J.; Yazaydin, A. O.; Burton, S. D.; Hoyt, D. W.; Ilton, E. S.; McGrail, B. P.; Loring, J. S., Competitive coordination of interlayer cations by carbon dioxide and clay. ACS Appl. Mater. Interfaces 2017, 9, 3678336791. 8. Schaef, H. T.; Loring, J. S.; Glezakou, V.-A.; Miller, Q. R. S.; Chen, J.; Owen, A. T.; Lee, M.-S.; Ilton, E. S.; Felmy, A. R.; McGrail, B. P.; Thompson, C. J., Competitive sorption of CO and H O in 2:1 layer phyllosilicates. Geochim. Cosmochim. Acta 2015, 161, 248-257. 9. Schaef, H. T.; Ilton, E. S.; Qafoku, O.; Martin, P. F.; Felmy, A.; Rosso, K. M., In situ XRD study of Ca saturated montmorillonite (STx-1) exposed to anhydrous and wet supercritical carbon dioxide. Internat. J. Greenhouse Gas Contr. 2012, 6, 220-229. 10. Giesting, P.; Guggenheim, S.; Koster van Groos, A. F.; Busch, A., Interaction of carbon dioxide with Na-exchanged montmorillonite at pressures to 640 bars: Implications for CO sequestration. Internat. J. Greenhouse Gas Contr. 2012, 8, 73-81. 11. Giesting, P.; Guggenheim, S.; Koster van Groos, A. F.; Busch, A., X-ray diffraction study of K- and Ca-exchanged montmorillonites in CO atmospheres. Environ. Sci. Technol. 2012, 46, 5623-5630. 13

23

2

2

2

2

2

2

2

2

2

2+

2

2


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