Metal Carbonation of Forsterite in Supercritical CO2 and H2O Using

Feb 17, 2010 - The acidic dissolution kinetics of major rock-forming silicate minerals ..... Note that the formation of a near- surface amorphous regi...
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Metal Carbonation of Forsterite in Supercritical CO2 and H2O Using Solid State 29Si, 13C NMR Spectroscopy Ja Hun Kwak, Jian Zhi Hu,* David W. Hoyt, Jesse A. Sears, Chongming Wang, Kevin M. Rosso, and Andrew R. Felmy Pacific Northwest National Laboratory, P.O. Box 999, MS K8-98, Richland, Washington 99352 ReceiVed: January 6, 2010

Ex situ natural abundance magic angle spinning (MAS) NMR was used for the first time to study fundamental mineral carbonation processes and reaction extent relevant to geologic carbon sequestration (GCS) using a model silicate mineral forsterite (Mg2SiO4)+supercritical CO2 with and without H2O. Run conditions were 80 °C and 96 atm. With H2O but without CO2, 29Si MAS NMR reveals that the reaction products contain only two peaks of similar intensities located at about -84.8 and -91.8 ppm, which can be assigned to surface Q1 and Q2 species, i.e., SiO4 tetrahedra sharing one and two corners with other tetrahedra, respectively. Using scCO2 without H2O, no reaction is observed within 7 days. Using both scCO2 and H2O, the surface reaction products for silica are mainly Q4 species (-111.6 ppm) accompanied by a lesser amount of Q3 (-102 ppm) and Q2 (-91.8 ppm) species. No surface Q1 species were detected, indicating the carbonic acid formation and magnesite (MgCO3) precipitation reactions are faster than the forsterite hydrolysis process. Thus, it can be concluded that the Mg2SiO4 hydrolysis process is the rate limiting step of the overall mineral carbonation process. 29Si NMR combined with XRD, TEM, SAED, and EDX further reveals that the reaction is a surface reaction with the Mg2SiO4 crystallite in the core and with condensed Q2, Q3, and Q4 species forming highly porous amorphous surface layers. 13C MAS NMR unambiguously identified a reaction intermediate as Mg5(CO3)4(OH)2 · 5H2O, i.e., the dypingite. Introduction Burning fossil fuels such as coal, a major source of CO2 emissions to the atmosphere, appear likely to continue to provide a significant portion of total energy in both industrialized and developing countries. The widespread consensus in the scientific community is that increased levels of greenhouse gases such as CO2 are adversely affecting the global environment, as evidenced by trends in global warming and dramatic changes in weather patterns. These facts have made it critical to develop technologies that help offset emissions and stabilize CO2 in the environment.1 One proposed method for reducing CO2 emission is to store it underground in geologic repositories as a supercritical fluid.2-5 Main target repositories include saline aquifers and depleted oil reservoirs, but there also is interest in injection into alkaline rock formations rich in magnesium and iron.5 So-called geologic carbon sequestration (GCS) entails a number of possible longterm trapping mechanisms that can in principle keep the CO2 sequestered indefinitely. One such mechanism is trapping by mineralization of metal carbonate solids. In conditions where the availabilities of aqueous divalent cations such as Ca2+, Fe2+, and Mg2+ are sufficiently high, precipitation of metal carbonates can ensue by combining with dissolved CO2 species, yielding effectively permanent storage of the CO2. Little is known about rate-controlling reaction processes occurring during this transformation. It is predominantly viewed as occurring primarily in the aqueous phase, acidified to low pH due to the high CO2 pressure. The needed cations must be continually freed from mineral matrixes through dissolution to sustain precipitation of metal carbonates. * To whom correspondence should be addressed. Phone: (509) 371-6544. Fax: (609) 371-6546. E-mail: [email protected].

The acidic dissolution kinetics of major rock-forming silicate minerals has long been studied6-11 but only recently in the context of GCS. For example, the orthosilicate mineral forsterite (Mg2SiO4) has been used as a model material given its relatively simple crystal structure and relatively high aqueous reactivity. In these prior studies, conditions ranged from ambient to elevated temperatures over a range of pH values and typically involved very low ratios of mineral to solution.12-18 Recently, Giammar et al.19 were the first to investigate acidic aqueous forsterite dissolution and resulting magnesite (MgCO3) precipitation at temperatures and CO2 pressures relevant to GCS. These researchers used a batch reactor with mineral to H2O weight ratio of approximately 1:50. In their experiment, the reacted aqueous phase was analyzed by elemental analysis and equilibrium speciation modeling, and the water washed solids were studied by scanning electron microscopy, infrared spectroscopy, and X-ray diffraction. The results obtained by Giammar et al. offer considerable insights into the effects of pH, temperature, CO2 pressure on orthosilicate mineral dissolution, and subsequent carbonate precipitation. The relevant forsterite transformation reactions are illustrated in eqs 1-5.19 Four key reactions are proposed, including CO2 dissolution into the aqueous phase 1, carbonate anion formation 2, proton-promoted forsterite dissolution 3, and magnesite precipitation 4.

CO2(gaseous) ) CO2(aqueous)

(1)

CO2(aqueous) + H2O ) CO32- + 2H+

(2)

10.1021/jp1001308  2010 American Chemical Society Published on Web 02/17/2010

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Mg2SiO4(solid) + 4H+ ) 2Mg2+ + H4SiO4

(3)

Mg2+ + CO32- ) MgCO3(solid)

(4)

The combination of the four elemental reactions yields the net mineral trapping reaction 5

2CO2(gaseous) + 2H2O + Mg2SiO4(solid) ) 2MgCO3(solid) + H4SiO4

(5)

Although much is known about this model system, fundamental molecular-level processes controlling the rate, reaction extent, and overall behavior remain obscured. Furthermore, more generally, the majority of all previous research on mineral-fluid interaction for GCS has focused entirely on the aqueous phase, i.e., reactivity with aqueous solution or brines containing dissolved CO2 at essentially constant CO2 pressure. Very little work has explored what is perhaps the more relevant system over the long term: supercritical CO2 (scCO2) containing small amounts of dissolved water.4 Gradual displacement or dissolution of residual pore water into a buoyant plume of CO2 at the caprock could lead to intermediate- or oil-wet conditions in which minerals make direct contact and interact with a wet CO2 phase. Key questions include the following: What is the role of water if any beyond simply serving as the reactant-bearing solvent for catalyzing mineral transformation to metal carbonate? What is the structure and composition of reacted mineral surfaces during dissolution? What limits conversion to metal carbonate at the mineral/solution interface? Are there any important reaction intermediates? In GCS, the ratio of mineral solid to aqueous solution is expected to be much higher than what has been used in previous work, which may facilitate carbonate mineral precipitation. Equation 5 predicts the silica dissolution products are H4SiO4, which is oversimplified because oligomeric silica condensation inevitably occurs if the solid/solution ratio is high. Precipitation of silica to form amorphous silica networks is hypothesized as being important, but this has not yet been confirmed experimentally in GCS. Previous analytical tools employed were not capable of making this determination. For example, XRD detects primarily crystalline phases. Furthermore, washing of reacted solids would sacrifice the opportunity to detect amorphous silica. Nuclear magnetic resonance (NMR), a quantitative, nondestructive, element-specific probe of local structure with the capability of accessing buried surfaces/interphases regardless of their crystallinity is an ideal tool for addressing many of the scientific questions outlined above. In this work, 29Si and 13C magic angle sample spinning (MAS) NMR were used, for the first time, to investigate fundamental processes in metal carbonation of forsterite. High resolution TEM and XRD were also used to complement the NMR results. The purposes of this work were to investigate the reaction pathways and reaction extent, including phase evolution, in forsterite carbonation under a range of GCS fluid conditions and high solid to water ratio, and to demonstrate the power of NMR as an analytical tool for studying the mineralization process at a fundamental level. Pure H2O, dry scCO2, and H2O-scCO2 mixtures were used to sample extreme fluid conditions. For fluid mixtures, the mineral:H2O weight ratio of 1:1 was used, which yields a CO2-acidified H2O rich phase coexisting with a wet scCO2 rich phase at our run conditions. For the latter, the solubility of H2O in scCO2 is less than about 0.01 molar fractions near the critical point.20 The

Figure 1. Apparatus used for setting up the initial CO2 pressure and for carrying out the reaction. For condensation of CO2, container “e” was filled with liquid nitrogen. For carrying out the reaction, “e” was filled with H2O and its temperature was controlled by a thermometer and an electric part.

measurements were performed on forsterite powders in sealed batch reactors, and therefore, in these first experiments, NMR captures the ensemble average behavior of this mixed fluid phase reaction system. Evaluation of a single fluid phase reaction, i.e., a wet scCO2 rich phase, will be the subject of a subsequent investigation. Experimental Section Sample Preparation. Forsterite, i.e., Mg2SiO4, obtained from Alfa Aesar was used as received. High chemical purity and high crystallinity were confirmed by XRD (JPCDS card # 34-0189). The BET surface area of this as-received sample was about 1.0 m2/g. Unless otherwise specified, all of the reaction experiments were carried out using this sample as the reactant for Mg2SiO4. The device illustrated in Figure 1 was used to set up the initial CO2 pressure in the reactor. The operation principles are briefly described below. The output from the pressure gauge “p” of a CO2 cylinder, set at 5 atm, was connected to a 40 mL stainless steel container, “a”, via a three way valve, “b”. Once “a” was filled with 5 atm of CO2, “b” was switched to the channel connecting the reactor “d” through an on/off valve, “c”, followed by condensing CO2 in the reactor using liquid nitrogen. Reactor “d”, including the volume between “c” and “d”, had a volume of 11.2 mL before loading. After loading with 1 g of Mg2SiO4 mixed with 1 g of H2O, the net volume for CO2 storage in reactor “d” was about 10 mL once switch “c” was closed. The desired CO2 pressure at room temperature in the reactor was set using standard volumetric adsorption methods by condensing CO2 into the reactor multiple times using liquid nitrogen. For example, to reach a target CO2 pressure of 80 atm in the reactor, four condensing cycles were needed because the volume ratio of the calibrated volume “a” to the net volume of reactor “d” was 4 to 1 and the pressure in the calibrated vessel “a” was 5 atm. At the end of CO2 condensing, switch “c” was closed and the liquid nitrogen bath was replaced by a water bath at a constant temperature of 80 °C to carry out the reaction for 20 h, 4 days, and 7 days, with an initial CO2 pressure of 80 atm at room temperature (RT) or a CO2 pressure of about 96 atm at 80 °C. Using a CO2 critical density of 0.486 g/cm3 at a critical temperature of 31 °C and critical pressure of 73.8 bar21 to approximate the calculation, the molar ratio of mineral, H2O, and scCO2 in the current reaction system is about 1:7.8:15. At the end of each reaction at specified times, the residual pressure from the reactor was released and the solids were collected and flow dried for 20 h at ambient temperature using dry nitrogen purge. These moderately dried samples were used for the various ex situ analyses described below.

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NMR Experiments. The single pulse 29Si and 13C magic angle spinning (MAS) experiments combined with high power 1 H decoupling were performed on a Varian-Chemagnetics 300 MHz Infinity spectrometer, corresponding to 1H, 13C, and 29Si Larmor frequencies of 299.982, 75.43, and 59.59 MHz, respectively. A commercial cross-polarization/MAS probe with a 7.5 mm outside diameter and 6 mm internal diameter pencil type spinner system was used. The sample cell resembles the commercial cell except that two solid Teflon plugs were made in such a manner that they can only be fully inserted into the zirconium cylinder after precooling using liquid nitrogen to seal ∼0.4 g of samples. The sample spinning rate used for all of the measurements was about 5.0 kHz. Tetrakis(trimethylsilyl) silane (TKS), [(CH3)3Si]4Si, was used as the secondary reference for 13 C (3.5 ppm) and 29Si (-9.8 ppm) relative to TMS (0 ppm for both 13C and 29Si).22 The pulse angle for acquiring both 13C and 29Si spectra was approximately 45°. The recycle times used were 5 s for 13C and 20 s for 29Si, respectively. All of the spectra were acquired at room temperature. XRD Analysis. XRD was carried out on a Philips PW3040/ 00 X’Pert powder X-ray diffractometer using Cu KR1 radiation (λ ) 1.5406 Å) in step mode between 2θ values of 15 and 75°, with a step size of 0.04°/s. Data analysis was accomplished using JADE (Materials Data, Inc., Livermore, CA) as well as the Powder Diffraction File database (2003 Release, International Center for Diffraction Data, Newtown Square, PA). High Resolution TEM Analysis. TEM was carried out on a JEOL JEM 2010 microscope with 200 keV operating voltage. The composition of the particles was analyzed by energydispersive X-ray spectroscopy (EDX). TEM specimens were prepared by dusting the powder particles onto a porous carbon film coated 200 mesh copper TEM grid. Results and Discussion 29

Si and 13C NMR Investigations. The 29Si MAS spectrum of as-purchased forsterite (Figure 2a) shows a single peak at -61.9 ppm arising from bulk Mg2SiO4, i.e., the magnesium isolated SiO4 tetrahedra in Mg2SiO4, consistent with a previous report.23 No other Si peak was observed except spinning sidebands. This result clearly demonstrates the chemical purity of forsterite used in the present work. The Role of H2O. To understand the role of H2O on CO2 uptake and mineral carbonation mechanisms, 29Si MAS spectra on samples collected after reaction at 80 °C and 96 atm of scCO2 without (Figure 2b) and with H2O (Figure 2c) for 20 h were acquired. Figure 2b is similar to Figure 2a, indicating no reaction occurred for the sample in dry scCO2 without H2O after 20 h. However, reaction product peaks, centered at about -102 and -111.6 ppm, are observed in Figure 2c on the samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C for 20 h. Using the integrated peak area, it was found that ∼8% of Mg2SiO4 was converted to products. To further clarify the role of H2O, 29Si MAS experiments on samples of 1 g of Mg2SiO4 + 1 g of H2O at 80 °C for 20 h and 7 days were carried out and the corresponding spectra are given in Figure 2d and e, respectively. Clearly, with pure H2O and in the absence of CO2, two peaks with approximately equal intensities (centered at about -84.8 and -91.8 ppm) are observed and conversion of Mg2SiO4 to products increases slightly from ∼1.6% (20 h reaction in Figure 2d) to ∼3.2% (7 day reaction in Figure 2e). Since for the pure water cases the system pressure was only the vapor pressure of H2O, there is concern of the effect of pressure in comparison to runs that include scCO2. To resolve this concern, the same reaction as those of Figure 2d and e but

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Figure 2. 29Si MAS NMR spectra. The spectra were vertically expanded by 32 times to highlight the reaction products. (a) The aspurchased Mg2SiO4; (b) 1 g of Mg2SiO4 + 96 atm of CO2 at 80 °C for 20 h; (c) 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C for 20 h; (d) 1 g of Mg2SiO4 + 1 g of H2O at 80 °C for 20 h; (e) 1 g of Mg2SiO4 + 1 g of H2O at 80 °C for 7 days; (f) 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of argon at 80 °C for 7 days. The number of accumulations was 4000 (a), 4000 (b), 3296 (c), 7824 (d), 8298 (e), and 26109 (f), respectively. The BET surface area of starting Mg2SiO4 was 1 m2/g.

in the presence of 96 atm of argon gas was carried out for 7 days and the corresponding 29Si MAS spectrum is plotted in Figure 2f. Again, two sharp reaction peaks of equal intensity centered at about -84.8 and -91.8 ppm are observed with similar area ratio to the overall peak area, i.e., ∼4%, as that of Figure 2e (3.2%) for a reaction time of 7 days. Although there are also broad peaks observed between -91.8 and -116 ppm, their peak intensities are much lower than those at -84.8 and -91.8 ppm. It has been reported by Xue et al.24 that the 1H-29Si CP/MAS NMR spectrum of the hydrated Mg2SiO4 sample contains two peaks of similar intensities near -86 and -93 ppm. These two CP/MAS peaks are very similar to the -84.8 and -91.8 ppm peaks observed in our current study. Thus, it can be concluded that the -84.8 and -91.8 ppm peaks in Figure 2 are surface silicon species that are close to protons. Note that the formation of a near- surface amorphous region enriched in silicon and hydrogen and depleted in magnesium has also been observed in dissolution of fosterite (refs 6 and 7) in aqueous solutions based on the results from Raman and elemental depth profile studies. Using the conventional symbols Qn for SiO4 tetrahedra, where the index n shows the number of other SiO4 tetrahedra sharing corner(s) with the tetrahedron under consideration, we can designate the type of silica species in this work as Q1, Q2, Q3, and Q4. On the basis of prior extensive literature reports on 29Si NMR chemical shifts,23,25-28 the most likely surface silica structures that give two resolved silica peaks with equal intensities would be Q1 (-84.8 ppm) and Q2 (-91.8 ppm) (see Figure 2). The peak centered at about -102 ppm is assigned to Q3 and likewise the peak centered about -111.6 ppm is assigned to Q4 in accordance with literature reports (refs 23 and 25-28). Note that, with the silica condensation increases, i.e., the increase of index n, approximately -10 ppm shift each time an increased silica condensation from Q2 to Q3 to Q4 is

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Figure 3. 29Si MAS NMR spectra of samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C for 20 h (a), 4 days (b), and 7 days (c), respectively. The BET surface area of starting Mg2SiO4 was 1 m2/g. (d) 29Si MAS spectrum of a commercial 100% amorphous silica sample. The number of accumulations was 3296 (a), 4436 (b), 12400 (c), and 6000 (d), respectively.

observed. Simple silicic acid, H4SiO4/Si(OH)4 (Q0), is not observed in the present work probably because its concentration is too low to contribute to an observable solid state 29Si MAS NMR signal. However, it has been well established29 that in aqueous solutions monomeric silica species, i.e., Q0, H3SiO4-, and H2SiO42-, are always present in solution. Combining the results obtained from Figure 2 and those from a prior literature report (ref 19), the roles of H2O in the initial carbonation reaction of Mg2SiO4 + H2O + scCO2 become apparent. H2O has participated in two key initial stages of the reaction. One of them is the hydrolysis or dissolution of Mg2SiO4 via eq 6

Mg2SiO4(solid) + 4H2O ) 2Mg2+ + H4SiO4 + 4OH-

Figure 4. Deconvolution of the 29Si MAS spectra between -130 and -65 ppm: (a) 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C for 7 days; (b) the commercial 100% amorphous silica Davisil’s645 (pore volume ) 1.15 mL/g, surface area ∼300 m2/g, and average pore size ∼15 nm). In each case, the top trace (black) is the experimental spectrum; the simulated peak position, relative peak area, and line width using Gaussian line shape are (a) Q2 (-92 ppm, 1.1%, 308 Hz), Q3 (-102 ppm, 21.6%, 543 Hz), and Q4 (-111.6 ppm, 77.3%, 505 Hz); (b) Q2 (-92 ppm, 2.3%, 304 Hz), Q3 (-102 ppm, 26.4%, 450 Hz), and Q4 (-111.6 ppm, 71.2%, 515 Hz).

Figure 5. Conversion of Mg2SiO4 in the system of 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C as a function of the reaction time. Area bars are estimated using the possible experimental errors in loading the various reaction agents into the reactor.

(6) The results from Figure 2 indicate that the hydrolysis reaction reaches an equilibrium between the monomeric silica species in solution and the surface Q1 and Q2 silicates in the absence of CO2. Without CO2, the condensation of Q1 and Q2 silicates to Q3 and Q4 is hindered and is not observed over the course of reaction (Figure 2d-f). The second role of H2O is that of solvent for CO2 which enables carbonate anion formation as described by eqs 1 and 2. The Role of CO2. Figure 3a-c gives the 29Si MAS spectra acquired on the samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C and its dependence on reaction time. The reaction products consist of predominantly Q3 and Q4 species in addition to a very minor contribution from Q2 species. The existence of a very small amount of Q2 is evident from the line shape simulation of the 29Si MAS spectrum, as depicted in Figure 4a. The absence of Q1 and the presence of only a small amount of Q2 in the spectra at all reaction times indicates that the condensation of Q1 and Q2 to Q3 and Q4 occurs at a time scale faster than the hydrolysis reaction given in eq 6. In other words, hydrolysis appears to be one of the rate limiting steps responsible for the overall mineral carbonation reaction (eq 5). The conversion of Mg2SiO4 to Q2-Q4, estimated using

peak integrals in Figure 3, is plotted in Figure 5. About 8% of Mg2SiO4 is converted for a reaction time of 20 h, while about 47 and 67% of Mg2SiO4 are converted for reaction times of 4 and 7 days, respectively. An approximately linear relationship between the percentage of conversion and the reaction time is obtained. This result suggests that condensation of polymerized silica does not appreciably affect the overall reaction. The linear relationship is probably due to the rather slow reaction associated with the mineral carbonation process that is primarily determined by either the slow hydrolysis process or CO2 dissolution into the aqueous phase and subsequent carbonic acid formation and deprotonation (eqs 1 and 2), or a combination of both. Since H2O needs to penetrate through any reacted surface layer to continue the forsterite hydrolysis reaction of eq 6, apparently the presence of surface oligomeric silica still provides plenty of channels that H2O molecules can diffuse through to access the unreacted surface. A remaining question is whether the condensed silica is crystalline or amorphous. A lack of longrange order perhaps is more amenable to creation of channels to facilitate H2O diffusion. This hypothesis is confirmed by the 29 Si MAS spectrum in Figure 3d that was acquired on a highly porous 100% amorphous silica sample, Davisil’s-645, with a known pore volume of 1.15 mL/g, surface area of ∼300 m2/g,

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and average pore size of ∼15 nm. It is known from the line shape simulation in Figure 4 that, although the relative percentage of the Q2, Q3, and Q4 species are somewhat different between the commercial amorphous sample and the samples from this work, the similarities in the chemical shifts and the peak line width confirm the amorphous silica nature of the reaction products from the mixed system of Mg2SiO4 + H2O + CO2. Since both Q2 (containing two -OH groups) and Q3 (containing one -OH group) are located at the silica surface, a rough estimate of the surface area of the silica products from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C for 7 days (Figure 4a) can be made using the relative ratio of (Q2+Q3)/(Q2+Q3+Q4). This ratio is 22.7% for Figure 4a and 28.7% for Figure 4b. Given the known surface area of 300 m2/g for Figure 4b, the estimated surface area corresponding to Figure 4a would be (22.7/28.7) × 300 ) 237 m2/g. Therefore, it can be concluded that silica products from the reaction system of 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C are both amorphous and highly porous. The approximate linear relationship between the percentage of Mg2SiO4 conversion and the reaction time in Figure 5 also suggests that the impact of any pressure drop that may have occurred upon reaction progression in the sealed reactor did not noticeably affect the overall reaction. It is known from eq 5 that conversion of 1 mol of Mg2SiO4 would require the consumption of 2 mol of CO2. The initial number of moles of CO2 is calculated as 0.033 mol based on the ideal gas law of PV ) nRT, where V ) 0.01 L or 10 cm3, P ) 80 atm, T ) 293 K, and R ) 0.0821 atm · L/(mol · K). The initial number of moles for Mg2SiO4 is 0.0071 mol. For 8, 47, and 67% conversion of Mg2SiO4, the moles reacted for Mg2SiO4 would be 0.000057, 0.0033, and 0.0048. The corresponding amount of CO2 consumed would be 0.00011, 0.0067, and 0.0095 mol. The corresponding nonreacted CO2 would be 0.033, 0.027, and 0.024 mol. At the reaction temperature of 80 °C (353 K), at the end of the reaction, the CO2 pressure in the reactor would be 96, 77, and 68.9 atm for 8, 47, and 67% Mg2SiO4 conversion, respectively. Note, at a pressure of 68.9 atm and temperature of 353 K, CO2 is in a gaseous (subcritical) state, since the critical point of CO2 is T ) 31 °C and P ) 73.8 atm. The experimentally obtained linear relationship between the percentage of Mg2SiO4 conversion and the reaction time over a pressure range from 68.9 to 96 ppm indicates that the mineral carbonation process is not very sensitive to the pressure change near the critical conditions. Direct spectroscopic evidence of magnesium carbonate precipitation is demonstrated in Figure 6 with 13C MAS NMR spectra acquired on the samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C at two reaction times, i.e., 20 h (Figure 6a and c) and 7 days (Figure 6b and d), respectively. Note that the BET surface area of the starting Mg2SiO4 powder for the corresponding spectra a and b was 1.0 m2/g, while the BET surface area of the starting Mg2SiO4 powder for the corresponding spectra c and d was 2.1 m2/g. In Figure 6a, a dominant peak located at about 170 ppm is observed, whose calibrated integrated peak area is increased by about 9-fold in Figure 6b when the reaction time is increased to 7 days, consistent with the result obtained from 29Si MAS NMR (Figures 3 and 5). The peak at 170 ppm is assigned to crystalline MgCO3. This assignment is confirmed by the 13C MAS spectrum (Figure 6f) acquired on a powder sample prepared from a single crystal MgCO3 (Magnesite) from Brumado, Bahia, Brazil. In Figure 6f, the peak center is located at exactly 170 ppm, although the peak is broadened by a trace amount of Fe3+ in the crystal. A

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Figure 6. (a-e) Natural abundance 13C MAS NMR spectra of samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C. The reaction time was 20 h for spectrum a and 7 days for spectrum b. In both spectra a and b, the BET surface area of the starting Mg2SiO4 material was 1.0 m2/g. Similarly, the reaction time was 20 h for spectrum c and 7 days for spectrum d. However, in both spectra c and d, the BET surface area of the starting Mg2SiO4 material was 2.1 m2/g. (e) 13C MAS NMR spectrum of the commercial Mg5(CO3)4(OH)2 · 5H2O sample. The number of accumulation numbers was 4000 (a), 4236 (b), 1744 (c), 17455 (d), and 4000 (e), respectively. The relative 13C peak areas between the chemical shifts of 160 and 175 ppm normalized to the same number of scans per unit sample weight are approximately 1 (a), 9.0 (b), 2.0 (c), and 7.6 (d), respectively. The area ratio of the 170 ppm peak to that of the combined 164.7 and 162.5 ppm peaks is approximately 1:1.1 in trace c. (f) Crystalline MgCO3 (Magnesite) from Brumado, Bahia, Brazil, Excalibur Mineral Company, Rare Minerals, Meteorites & Analytical Service. The peak center is at 170 ppm. Line broadening is caused by a trace of Fe3+ in the crystal.

careful evaluation of Figure 6a indicates two additional peaks located at about 164.7 and 162.5 ppm, but their corresponding peak intensities are very low. To confirm the existence of the 164.7 and 162.5 ppm peaks, the same reaction experiment as that of Figure 6a was carried out but using 99% 13C isotope enriched CO2. Clearly, the two peaks are observed (see Figure S1 in the Supporting Information). These two low intensity peaks completely disappear in Figure 6b at the long reaction time of 7 days. For the same reaction system but with the BET surface area of 2.1 m2/g for the starting Mg2SiO4 materials, the 13 C peaks located at 162.5 and 164.7 ppm are significantly increased at a reaction time of 20 h (Figure 6c). Again, these two peaks disappear when the reaction time is increased to 7 days (Figure 6d). It is also found that the total peak area of the 170, 164.7, and 162.5 ppm peaks in Figure 6c normalized to the same number of scans per unit sample weight is about 2.1 times that of the corresponding peak area in Figure 6a. Quantitative 29Si MAS NMR (Figure S2, Supporting Information) also indicates that about 18.7% of the Mg2SiO4 associated with Figure 6c has converted into amorphous silica. It is known from Figure 5 that only about 8% of the Mg2SiO4 has converted at a reaction time of 20 h in the case of a BET surface area of 1 m2/g for the Mg2SiO4. These results indicate that the reaction product at a reaction time of 20 h is approximately proportional to the surface area of the minerals. In other words, the reaction

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Figure 7. Left traces: XRD patterns of samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C and at a reaction time of 20 h, 4 days, and 7 days, respectively. The BET surface area of starting Mg2SiO4 was 1 m2/g. The bottom trace is the XRD acquired on a pure Mg2SiO4 sample. Right traces: The horizontally expanded XRD patterns of the corresponding left traces, highlighting the characteristic peaks of Mg2SiO4 and MgCO3.

is a surface-controlled reaction. The peaks located at 164.7 and 162.5 ppm are assigned to Mg5(CO3)4(OH)2 · 5H2O, based on the same 13C chemical shifts from the spectrum (Figure 6e) acquired on a pure commercial Mg5(CO3)4(OH)2 · 5H2O powder sample. Because at long reaction times the peaks located at both 162.5 and 164.7 ppm disappeared, this indicates that Mg5(CO3)4(OH)2 · 5H2O has converted to crystalline MgCO3. This result suggests that the hydrated hydroxylated magnesium carbonate, i.e., Mg5(CO3)4(OH)2 · 5H2O, is a reaction intermediate and precursor to magnesium carbonate crystallization. Crystalline Mg5(CO3)4(OH)2 · 5H2O is the mineral dypingite that sometimes occurs in altered serpentenites.30 This mineral has been suggested to form under geological carbon sequestration sites upon injecting scCO2.31 The current results confirm the formation of dypingite as a reaction intermediate, providing a direct link between dypingite and the formation of real rocks, i.e., crystalline MgCO3. On the basis of the above findings, the role of CO2 can be summarized as follows. Without CO2, the hydrolysis reaction of Mg2SiO4 reaches an equilibrium state, i.e., the concentrations of the Mg2+, Q0, and other possible monomeric silica species in solution and the surface silica Q1 and Q2 species reach an equilibrium saturation state. In the presence of scCO2, or near the condition of scCO2, Mg2+ continuously reacts with CO32to form MgCO3 (eqs 1, 2, and 4). This reaction drives condensation of polymerized silica from Q0, Q1, and Q2 to highly porous and amorphous structures of Q3 and Q4. XRD and TEM Studies. XRD and TEM were used for validating the formation of MgCO3 and for investigating the morphological evolution and phase changes in the solids associated with magnesium carbonation. Shown in Figure 7 are the XRD patterns obtained on samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C as a function of the reaction time, where the XRD pattern of Mg2SiO4 is also included for comparison. Because of the nature of the XRD technique, quantification of Mg2SiO4 conversion is very difficult. Furthermore, it is very difficult to quantify the reaction products which have amorphous structures, due to their broad line shape. However, qualitative results obtained from the products with crystalline structure, MgCO3, show very consistent results with the 29Si MAS NMR results. For a reaction time of 20 h, a shoulder peak with low intensity was observed at 2Θ ) ∼32.65°. This peak, absent from the pure Mg2SiO4 spectrum (bottom trace in Figure 7 and its corresponding expansion at the right), can be assigned to MgCO3 according to the standard database (JCPDS card # 03-478). It is known that, at a reaction time of 20 h, 8% of Mg2SiO4 had already reacted based on 29Si NMR. In Figure 7, the relative peak intensity for the MgCO3

phase at 2Θ ) ∼32.65° increases drastically with reaction time. However, crystalline silica diffraction patterns of any kind were not detected even though ∼67% of Mg2SiO4 was hydrolyzed after 7 days of reaction. Only very weak and very broad amorphous peaks were detected. This result indicates that the silica species produced by hydrolysis are in a highly amorphous state, consistent with results drawn from our 29Si MAS NMR studies. A detailed evaluation of the XRD patterns reveals that the average crystallite size of the remaining Mg2SiO4 does not change, as revealed by the similar line width of the characteristic XRD peak for crystallite Mg2SiO4 located at about 2Θ ) ∼22.88° with the reaction time (see Figure S3 in the Supporting Information). Although this behavior seems against the hypotheses that the hydrolysis reaction starts on the surface of Mg2SiO4, the expectation is for smaller grain sizes to evolve. However, the finding could also be interpreted as preferential consumption of small sized Mg2SiO4 particles in the conversion to MgCO3 and amorphous silica, leaving behind more of the larger sized Mg2SiO4 particles. The existence of the crystallite Mg5(CO3)4(OH)2 · 5H2O is evident at 2Θ ) ∼30.75° in the XRD pattern of 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C and for a reaction time of 20 h using Mg2SiO4 powders with a BET surface area of 2.1 m2/g (Figure 8). However, the intensity of the 2Θ ) ∼30.75° peak is very low, approximately 1/10 that of the MgCO3 peak (see the corresponding expanded pattern at the right in Figure 8). In contrast, the peak area corresponding to Mg5(CO3)4(OH)2 · 5H2O (the 164.7 and 162.5 ppm peaks) in the 13 C MAS spectrum acquired on the same sample (Figure 6c) is 1:1 times the crystallite MgCO3 peak (170 ppm). These results clearly indicate that a majority of the Mg5(CO3)4(OH)2 · 5H2O species is in a highly amorphous state. The reason is that NMR detects both crystalline and amorphous structures while XRD detects primarily the crystalline structures. For reaction times of 4 and 7 days, the diffraction peak at 2Θ ) ∼30.75° completely disappears. This result is consistent with the 13 C MAS NMR-based conclusions discussed above that Mg5(CO3)4(OH)2 · 5H2O is a reaction intermediate and is converted to crystalline MgCO3 at long reaction times. In contrast, there is no 2Θ ) ∼30.75° in the XRD pattern of 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C and for a reaction time of 20 h using Mg2SiO4 powders with a BET surface area of 1.0 m2/g (the red trace in Figure 7). This is because the amount of intermediate, Mg5(CO3)4(OH)2 · 5H2O, is both relatively small and in a highly amorphous state when the surface area is low. To further confirm the concept of surface reaction and to explore morphology change, high resolution TEM images with

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Figure 8. Left traces: XRD patterns of samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C and at a reaction time of 20 h, 4 days, and 7 days, respectively. The BET surface area of starting Mg2SiO4 was 2.1 m2/g. The bottom trace is the XRD acquired on a pure Mg2SiO4 sample. Right traces: The horizontally and vertically expanded XRD patterns of the corresponding left traces, highlighting the characteristic peaks of Mg2SiO4, MgCO3, and Mg5(CO3)4(OH)2 · 5H2O.

Figure 9. TEM images of as-purchased Mg2SiO4 at different plot scales (a, b, and c). (d) SAED (selected area electron diffraction) pattern.

selected area electron diffraction (SAED) and energy-dispersive X-ray spectra (EDX) were performed. The TEM images of the starting Mg2SiO4 materials demonstrate a broad particle size distribution (Figure 9a and b). However, all of the particles show a highly crystalline phase as evidenced in the high resolution lattice image (Figure 9c) and SAED (Figure 9d) regardless of the particle size examined. TEM images of the samples with 8% conversion (Figure 10) show very interesting morphologies. A significant surface layer, which was confirmed as silica-based by EDX (bottom trace in Figure 10d), was detected with crystalline Mg2SiO4 in the core (Figure 10c). The TEM images clearly demonstrate that the hydrolysis reaction happened on the surface of Mg2SiO4 and only the magnesium ion leached out from the Mg2SiO4 surface and silica species remained as amorphous silica, consistent with 29 Si MAS NMR studies. Interestingly, each particle examined shows a different thicknesses of amorphous silica layers. Some particles almost completely converted to amorphous silica, but some particles have hardly any surface amorphous layers. Generally, larger crystals show thinner amorphous silica layers. This observation is consistent with the above explanation for why the average particle size obtained by XRD does not change

Figure 10. TEM images of samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C for a reaction time of 20 h (a, b, and c), demonstrating the surface amorphous silica layers. The BET surface area of starting Mg2SiO4 was 1 m2/g. (d) The energy-dispersive X-ray spectra (EDX) for spot analysis of this sample show characteristic spectra of the Mg2SiO4 crystal in the core (top trace) and the amorphous silica in the surface layer (bottom trace).

with increasing conversion. Thus, it can be concluded that, on average, larger size particles are less reactive in this system than smaller size particles. The reaction rate increases with increasing surface area, which apparently is primarily due to more facile release of magnesium into the aqueous solution in the case of smaller sized particles. This is particularly evident in samples at longer reaction times. The TEM images (Figure 11) collected from the sample after 4 days of reaction, i.e., with ∼47% conversion of Mg2SiO4 based on our 29Si MAS NMR, show highly heterogeneous morphology. For example, some of the Mg2SiO4 particles (Figure 11a) do not show noticeable amorphous silica layers on the surface. In contrast, some domains (Figure 11b) show only amorphous silica with essentially no magnesium (see top trace in Figure 11d). Over 5 µm sized perfect crystalline MgCO3 particles were observed (see Figure 11c). All of these results suggest that the smaller sized Mg2SiO4 particles tend to react faster than the larger sized particles.

Metal Carbonation of Forsterite

Figure 11. TEM images of samples from 1 g of Mg2SiO4 + 1 g of H2O + 96 atm of CO2 at 80 °C for a reaction time of 4 days at three selected spots to highlight the different phases: Mg2SiO4 crystal (a); amorphous silica (b); and MgCO3 crystal (c), respectively. (d) Top trace: EDX pattern corresponding to a spot in part b, where only the signal from silica is observed. Bottom trace: EDX corresponding to a spot in part c, where only the signal from crystalline MgCO3 is observed. The BET surface area of starting Mg2SiO4 was 1 m2/g.

Conclusions In this work, the rate and extent of reaction and associated phase changes of metal carbonation of the orthosilicate forsterite were studied in a mixture of Mg2SiO4+CO2+H2O with high mineral solid to water weight ratio of 1:1. A combination of solid state 29 Si and 13C MAS NMR, XRD, and TEM was used at or near the scCO2 condition and at a reaction temperature of 80 °C. 29Si NMR clearly shows, that in the absence of CO2 (i.e., the system of Mg2SiO4+H2O), the role of H2O is to hydrolyze Mg2SiO4, producing Mg2+ + Q0 + OH- and surface silica Q1 and Q2 species. The hydrolyzed surface silica products contain only Q1 and Q2 species with approximately 1:1 ratio. An equilibrium between Q0, Q1, Q2, and Mg2+ with a saturated concentration equivalent to less than about 3.2% of the Mg2SiO4 conversion is obtained at a reaction time of up to 7 days. In contrast, in the absence of H2O (i.e., the system of Mg2SiO4+CO2), no reaction product is observed, or the anticipated products have quantities too low to be detected by NMR. However, in the presence of both scCO2 and H2O (i.e., Mg2SiO4+CO2+H2O), the reaction products for silica are mainly Q4 species accompanied by a lesser amount of Q3 and Q2. No surface Q1 species were observed, indicating the carbonic anion formation and magnesite (MgCO3) precipitation reactions are faster than the Mg2SiO4 hydrolysis process. On the basis of these results, it can be concluded that the Mg2SiO4 hydrolysis process is the rate limiting step of the overall metal carbonation reaction in this model system. 29 Si NMR combined with XRD, TEM, SAED, and EDX further reveals that the reaction is a surface reaction with the Mg2SiO4 crystallite in the core and condensed Q2-Q4 species forming highly porous amorphous surface layers. 13C MAS NMR and XRD identified a reaction intermediate as Mg5(CO3)4(OH)2 · 5H2O. At long reaction times, only crystallite Mg2SiO4 phases are observed. Combining these new findings, the reaction pathway of metal carbonation of Mg2SiO4 in the presence of scCO2 and H2O can

J. Phys. Chem. C, Vol. 114, No. 9, 2010 4133 be briefly summarized as involving the following steps. Step 1: H2O reacts with the surface of Mg2SiO4 to generate the hydrolysis products according to eq 6. The hydrolysis process is heterogeneous and is likely caused by the overall difference in lattice defects on the Mg2SiO4 surface relative to particle size. Since smaller sized particles tend to have more lattice defects per unit weight than larger sized particles, a larger proportion of smaller sized particles would be expected to hydrolyze faster than larger sized particles. Simultaneous to hydrolysis, H2O reacts with CO2 to generate carbonate anion, i.e., CO32- + 2 H+, according to eq 2. Step 2: Released magnesium ion, Mg2+, diffuses away from the hydrolyzed Mg2SiO4 surface and into the aqueous phase and reacts with CO32- to form MgCO3 with Mg5(CO3)4(OH)2 · 5H2O as a reaction intermediate. The precipitation of MgCO3 effectively consumes the Mg2+ in the aqueous solution. This continuously drives the hydrolysis reaction of Mg2SiO4 toward the formation of MgCO3 crystals. Coincidentally, the silica products from the hydrolysis reaction of Mg2SiO4 polymerize and condense into Q2 to Q4 species, with Q4 as the dominant structure which is accelerated with the precipitation reaction of MgCO3. The condensed silica forms a highly porous amorphous layer around the Mg2SiO4 surface. Probably due to the required charge balance in the aqueous solution, Mg2+ ions are not noticeably observed inside the amorphous layers, meaning Mg2+ ions preferentially partition into the aqueous phase and participate in magnesium carbonate precipitation. Acknowledgment. This work was supported by the Carbon Sequestration Initiative and funded by Laboratory Directed Research and Development at Pacific Northwest National Laboratory (PNNL) and the U.S. Department of Energy (DOE), Office of Basic Energy Sciences, through a Single Investigator Small Group Research (SISGR) grant. All of the experiments were performed at the Environmental Molecular Science Laboratory, a national scientific user facility sponsored by the DOE Office of Biological and Environmental Research, and located at PNNL. PNNL is operated for DOE by Battelle Memorial Institute under Contract No. DE-AC06-76RLO-1830. Dr. Herbert T. Schaet is acknowledged for donating the MgCO3 crystal for the 13C MAS NMR given in Figure 6f. Supporting Information Available: Figures showing 13C MAS/NMR and 29Si MAS/NMR spectra and expanded XRD patterns. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Basic Research Needs For Geosciences: Facilitating 21ST Century Energy Systems. Work Shop Report, Office of Basic Energy Sciences, U.S. Department of Energy, February 21-23, 2007. (2) Bachu, S. Energy ConVers. Manage. 2000, 41, 953–970. (3) Bachu, S. Prog. Energy Combust. Sci. 2008, 34, 254–271. (4) Saripalli, P.; McGrail, P. Energy ConVers. Manage. 2002, 43, 185– 198. (5) Bruant, R. G., Jr.; Celia, M. A.; Guswa, A. J.; Peters, C. A. EnViron. Sci. Technol. 2002, 36, 240A–245A. (6) Casey, W. H.; Westrich, H. R.; Banfield, J. F.; Ferruzzi, G.; Arnold, G. W. Nature 1993, 366, 253–256. (7) Casey, W. H.; Swaddle, T. W. ReV. Geophys. 2003, 41 (2), 1008. (8) Matson, D. W.; Sharma, S. K.; Philpotts, J. A. Am. Mineral. 1986, 71, 694–704. (9) Magi, M.; Lippmaa, E.; Samoson, A.; Engelhardt, G.; Grimmer, A.-R. J. Phys. Chem. 1984, 88, 1518–1522. (10) Tsomaia, N.; Brantley, S.; Hamilton, J.; Pantano, C. G.; Mueller, K. T. Am. Mineral. 2003, 88, 54–67. (11) Alexander, G. B.; Heston, W. M.; Ller, R. K. J. Phys. Chem. 1954, 58, 453–455. (12) Blum, A.; Lasaga, A. C. Nature 1988, 331, 431–433.

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(13) Chen, Y.; Brantley, S. L. Chem. Geol. 2000, 165, 267–281. (14) Goff, F.; Lackner, K. S. EnViron. Geosci. 1998, 5, 89. (15) Oelkers, E. H. Chem. Geol. 2001, 175, 485–494. (16) Rosso, J. J.; Rimstidt, J. D. Geochim. Cosmochim. Acta 2000, 64, 797–811. (17) Westrich, H. R.; Cygan, R. T.; Casey, W. H.; Zemitis, C.; Arnold, G. W. Am. J. Sci. 1993, 293, 869–893. (18) Wogelius, R. A.; Walther, J. V. Geochim. Cosmochim. Acta 1991, 55, 943–954. (19) Giammar, D. E.; Bruant, R. G.; Peters, C. A. Chem. Geol. 2005, 217, 257–276. (20) Sabirzyanov, A. N.; II’in, A. P.; Akhunov, A. R.; Gumerov, F. M. High Temp. 2002, 40, 203–206. (21) Buhler, E.; Dobrynin, A. V.; DeSimone, J. M.; Rubinstein, M. Macromolecules 1998, 31, 7347–7355. (22) Hayashi, S.; Hayamizu, K. Bull. Chem. Soc. Jpn. 1991, 64, 685– 687.

Kwak et al. (23) Magi, M.; Lippmaa, E.; Samoson, A.; Engelhardt, G.; Grimmer, A. R. J. Phys. Chem. 1984, 88, 1518–1522. (24) Xue, X.; Kanzaki, M.; Fraser, D. G. Geochim. Cosmochim. Acta 2002, 66 (15A), A853–A853. Supplement: Suppl. 1. (25) Lippma, E.; Magi, M.; Samoson, A.; Tarmak, M.; Engelhardt, G. J. Am. Chem. Soc. 1981, 103, 4992–4996. (26) Cho, H.; Felmy, A. R.; Craciun, R.; Keenum, J. P.; Shah, N.; Dixon, D. A. J. Am. Chem. Soc. 2006, 128, 2324–2335. (27) Chung, I.; Maciel, G. E. J. Phys. Chem. B 1997, 101, 3052–3064. (28) Weir, M. R.; Kuang, W.; Facey, G. A.; Detellier, C. Clays Clay Miner. 2002, 50, 240–247. (29) Felmy, A. R.; Cho, H.; Rustad, J. R.; Mason, M. J. J. Solution Chem. 2001, 30, 509–525. (30) Raade, G.; Museum, M. Am. Mineral. 1970, 55, 1457–1465. (31) Frost, R. L.; Bahfenne, S.; Graham, J.; Martens, W. N. Thermochim. Acta 2008, 475, 39–43.

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