Magnesium Hydroxide Dehydroxylation/Carbonation Reaction Processes: Implications for Carbon Dioxide Mineral Sequestration

J. Am. Ceram. Soc., 85 [4] 742– 48 (2002)

journal

Magnesium Hydroxide Dehydroxylation/Carbonation Reaction Processes: Implications for Carbon Dioxide Mineral Sequestration Hamdallah Be´arat,†,‡ Michael J. McKelvy,†,‡ Andrew V. G. Chizmeshya,† Renu Sharma,†,‡ and Ray W. Carpenter†,‡ Center for Solid State Science, Arizona State University, Tempe, Arizona

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Science and Engineering of Materials Ph.D. Program, Arizona State University, Tempe, Arizona

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Gas-phase magnesium hydroxide carbonation processes were investigated at high CO2 pressures to better understand the reaction mechanisms involved. Carbon and hydrogen elemental analysis, secondary ion mass spectrometry, ion beam analysis, X-ray diffraction, and thermogravimetric analysis were used to follow dehydroxylation/rehydroxylation/carbonation reaction processes. Dehydroxylation is found to generally precede carbonation as a distinct but interrelated process. Above the minimum CO2 pressure for brucite carbonation, both carbonation and dehydroxylation reactivity decrease with increasing CO2 pressure. Low-temperature dehydroxylation before carbonation can form porous intermediate materials with enhanced carbonation reactivity at reduced (e.g., ambient) temperature and pressure. Control of dehydroxylation/ rehydroxylation reactions before and/or during carbonation can substantially enhance carbonation reactivity. I.

Mineral sequestration involves permanent CO2 disposal in the form of thermodynamically stable alkaline earth carbonates, namely magnesite and calcite, which already occur in vast quantities in nature.7 This procedure also offers the capacity to sequester all the CO2 that could be emitted from the world’s known coal reserves.11 Several factors favor magnesium-rich minerals (e.g., serpentine and olivine) as feedstock sequestration materials. These include the vast abundance of these minerals over several continents, their availability in readily minable deposits, and their relatively high magnesium contents.7 Because mineral sequestration inherently satisfies the sequestration process criteria of (i) disposing of CO2 permanently and (ii) being an environmentally benign process, its primary goal is economically viable process development. An essential component of viable process development is to enhance mineral carbonation reaction rates and yields to reduce process cost. Several mineral carbonation reaction processes are being investigated, including gas–solid and aqueous solution reactions for brucite, serpentine (chrysotile, lizardite, and antigorite), and olivine (e.g., forsterite).8 –10,12 In each case, CO2 is converted through an exothermic reaction to the stable mineral magnesite. However, the carbonation processes and their mechanisms are quite complex and relatively poorly understood. Developing a sound fundamental understanding of the mechanisms that govern these carbonation processes is essential to enhancing their reaction rates and reducing process cost. Magnesium-rich minerals containing hydroxide lamella (e.g., brucite and serpentine) are appealing candidate mineral sequestration feedstock materials. In particular, the serpentines are widely available and can be mined readily at low cost.7,11 In addition, dehydroxylation, which inherently accompanies carbonation, can disrupt the structure at the atomic level, with the potential to substantially enhance carbonation reactivity. Magnesium hydroxide (Mg(OH)2) was chosen as a prototype material to initiate investigation of the associated dehydroxylation/carbonation mechanisms because of (i) its chemical and structural simplicity, (ii) the interest in Mg(OH)2 gas–solid carbonation as a possible CO2 mineral sequestration process component,12,13 and (iii) its chemical and structural similarity to other, more complex, magnesiumrich lamellar hydroxide minerals (e.g., the serpentine-based minerals chrysotile, antigorite, and lizardite), whose carbonation reaction processes offer exciting low-cost potential.7,11,14 Although thermal decomposition of Mg(OH)2 has been extensively studied, its gas-phase carbonation has not.15 In the study conducted by Butt et al.12 on the kinetics of simultaneous dehydroxylation and carbonation of brucite, it was suggested that MgCO3 precipitates as a thin layer of nanocrystalline carbonate on the surface of disrupted Mg(OH)2 crystallites. It was concluded that the carbonate acts as a diffusion barrier, which can explain the observed difference between the kinetics of dehydroxylation and simultaneous dehydroxylation/carbonation processes. At 0.76 atm CO2, the carbonation reaction rate increased up to 375°C and decreased at higher temperatures. The decrease in the extent of carbonation at higher temperatures was attributed to exceeding the thermal stability of the product magnesite, which is known to

Introduction

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OSSIL fuels are the major worldwide energy source and will continue to be into the foreseeable future.1 As global energy use continues to grow exponentially, the climatic and environmental effects associated with increasing anthropogenic CO2 emissions are of increasing concern.2,3 Alternative energy sources can provide energy without CO2 emission. However, they are unlikely to satisfy increasing global energy requirements in the foreseeable future, as 85% of the world’s energy is currently generated from fossil fuels.2,3,4 A potentially viable alternative is CO2 sequestration.2,3,5 Several approaches are being considered, including geological, oceanic, and terrestrial sequestration, as well as CO2 conversion into useful materials. Although both terrestrial sequestration and useful material conversion can significantly impact the level of anthropogenic CO2 emitted, they do not have the sequestration capacity to address the magnitude of CO2 emissions expected in the long term. Alternatively, both geological and oceanic sequestration can address the magnitude of anthropogenic emissions via long-term storage. However, the permanence of such storage is of critical concern. In contrast, permanence is not an issue for CO2 mineral sequestration, an emerging candidate technology.5–10

A. W. Searcy—contributing editor

Manuscript No. 188006. Received January 17, 2001; approved October 23, 2001. Supported by the National Energy Technology Laboratory, U.S. Department of Energy, under Grant No. DE-FG26-98FT40112. Research conducted as part of CO2 Mineral Sequestration Working Group managed by the U.S. Department of Energy, consisting of members from Albany Research Center, Arizona State University, Los Alamos National Laboratory, the National Energy Technology Laboratory, and Science Applications International Corporation. † Center for Solid State Science, Arizona State University. ‡ Science and Engineering of Materials Ph.D. Program, Arizona State University.

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decompose to MgO ⫹ CO2 at higher temperatures.16 For these runs, the observed carbonation apparently occurred during cooling, as the temperature fell below the temperature at which carbonation can occur at 0.76 atm CO2 (i.e., 385°C).12 During dehydroxylation and simultaneous dehydroxylation/ carbonation of the lamellar brucite structure, a variety of physical and morphologic changes can occur.12,17–23 As dehydroxylation progresses at relatively low temperatures, several transformations can take place in the crystal: (i) translamellar cracking, (ii) delamination, and (iii) morphologic reconstruction after extensive dehydroxylation to form nanostructured materials. These transformations control the morphology of the dehydroxylating material. The first two factors can disrupt the material down to the submicrometer/nanoscale range (Fig. 1),17 whereas the third is responsible for morphologic reconstruction that yields the wellknown high-surface-area product, MgO.17–23 During simultaneous dehydroxylation/carbonation, process timing can substantially affect carbonation reactivity. For example, relatively rapid dehydroxylation may favor MgO sintering/ particle-size growth (which may be aided by the water vapor formed), inhibiting carbonation reactivity in the process. In addition, the formation of carbonate passivating layers may inhibit both carbonation and dehydroxylation processes, even for relatively slow dehydroxylation rates.12 In this paper, we present the results of a series of studies on the simultaneous dehydroxylation and carbonation of relatively pure, natural Mg(OH)2 single-crystal fragments. Carbon and hydrogen analyses were used to follow the extent of dehydroxylation and carbonation, as a function of CO2 pressure and temperature, for the first time. Compositional analysis as a function of crystal depth provided further insight into the simultaneous dehydroxylation/ carbonation process. Rehydroxylation/carbonation processes were also studied to provide broader insight into the dehydroxylation/ rehydroxylation/carbonation reaction mechanisms. II.

Experimental Procedures

(1) Materials Natural single-crystal brucite (Delora, Canada) was used as a starting material for this study. Elemental analysis by protoninduced X-ray emission and total carbon analysis showed the material contained 0.16% Mn, 0.06% C, 0.01% Cl, and 0.01% Si. Loss-on-ignition at 1500°C was 99.7% of the theoretical weight loss of Mg(OH)2. The crystal was structurally characterized by XRD. The unit cell parameters of the material were a ⫽ 3.147(1) and c ⫽ 4.765(1) Å, which are in good agreement with the known parameters for brucite, a ⫽ 3.147 and c ⫽ 4.769 Å.24 Singlecrystal fragments were cleaved from the starting crystal for the

Fig. 1. Morphology of Mg(OH)2 crystal illustrating translamellar and lamellar (delamination) cracking that occurs during dehydroxylation.

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different types of experiments. Freshly cleaved crystal fragments were used to reduce the effect of structural defects (surface and internal), particle size, and surface impurities (e.g., MgO or MgCO3) on reactivity. (2) Reaction Setup and Conditions In each of the experiments presented here, several crystal fragments were placed in a quartz tube (⬃5 cm long and ⬃1 cm in diameter), which was then inserted in a high-pressure autoclave. The total mass of crystal fragments used for each run was held constant (⬃50 mg). The water vapor present was governed by Mg(OH)2 dehydroxylation. The reaction vessel was connected through a series of valves to a vacuum pump and to a dry CO2 gas cylinder (99.99%). The vessel was placed horizontally in a tube furnace. A thermocouple was attached to the outside of the vessel with its end located at the same position (furnace depth) as the crystals inside. The vessel was first evacuated and backfilled with CO2 three times. Then, CO2 pressure was brought to a preliminary pressure of about half the desired final pressure. The furnace was heated to reaction temperature and the CO2 pressure adjusted to its final value. The CO2 was in sufficient excess, so that the reaction pressure did not significantly change during carbonation. In some runs, samples were quenched from the reaction temperature and pressure. In this case, the reaction vessel was water-quenched from the reaction temperature, with the CO2 evacuated immediately after the temperature quench. All samples were then evacuated and stored in a helium-filled glove box (ⱕ1 ppm total O2 ⫹ H2O; M. Braun, Inc., Peabody, MA), before elemental analysis or other experiments, to avoid atmospheric CO2 or H2O uptake. (3) Analytical Techniques Thermogravimetric analysis (TGA) studies of in situ Mg(OH)2 dehydroxylation/rehydroxylation/carbonation processes were performed using a thermal analysis system (Model TG92, Setaram, Caluire, France). Sample weight changes were measured as a function of temperature and atmosphere for similarly sized brucite single-crystal fragments (three ⬃5 mg crystals/run) contained in alumina crucibles. Absolute weight sensitivity was 1 ␮g. Weight changes were observed to better than ⫾0.1%. The resulting TGA samples were evacuated and stored in the helium-filled glove box before analysis. Total hydroxide, carbonate, and oxide contents were determined using an analyzer (Model 2400 Series II CHNS Analyzer, Perkin– Elmer, Norwalk, CT). Comparative standards run before and after each sample gave a total weight percentage error of ⫾0.2% for carbon and ⫾0.3% for hydrogen. The total oxide present was determined by difference. Similarly sized (basal-plane surface area and thickness) crystals were used for comparative extent of dehydroxylation and carbonation studies. Powder XRD patterns were obtained using a diffractometer (Model D/MAX-IIB, Rigaku Co., Tokyo, Japan) with CuK␣ radiation. Samples were ground, adhered to a glass slide by spraying a vaseline-in-cyclohexane solution in N2 flow on the slide, and then the powder was spread uniformly onto it. Scans were done for 2␪ ⫽ 2° to 65° and with steps of 0.01°/s. Ion beam analyses were performed using a tandem ion accelerator system (1.7 MV; Model Tandetron, General Ionex Corp., Newburyport, MA). Magnesium, carbon, oxygen, and hydrogen concentrations were determined as a function of depth (using an ⬃1 mm diameter beam) from the basal plane for partially carbonated brucite single-crystal fragments with ⬃0.5 cm2 basal surface area and ⬃0.1 cm thickness. These results were integrated to give MgCO3, Mg(OH)2, and MgO concentrations as a function of depth from the sample basal plane. Magnesium was analyzed (⫾3 mol%) with normal Rutherford backscattering spectrometry (RBS) analysis using 2 MeV He2⫹ ions. Carbon and oxygen analyses were performed using nuclear reaction analysis. Carbon analysis (⫾1 mol%) used a C(␣,␣)C reaction that occurs at 4.65 MeV. Oxygen analysis (⫾2 mol%) was done using a similar reaction, which occurs at 3.05 MeV. Hydrogen analysis (⫾1 mol%) was performed using elastic recoil detection in a forward

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scattering mode, with the sample basal plane tilted at ␪ ⫽ 75° (with respect to the beam), using 2.8 MeV He2⫹ ions. Data simulation was done using RUMP (Rutherford Universal Manipulation Program).25 The surface areas of partially dehydroxylated, heat-treated Mg(OH)2 single-crystal fragments were determined using a dynamic sorption system (Model IGA-003, Hiden, Warrington, U.K.). Single-crystal fragments were heated under inert conditions, with subsequent BET measurements performed in situ to avoid atmospheric contact affecting the resulting sample surface area (e.g., via rehydroxylation), before surface area measurement. Secondary ion mass spectrometry (SIMS) analyses were performed with a spectrometer (Model 3f, Cameca IMS, Courbevoie, France). Analyses were conducted by rastering a 5 na Cs⫹ primary beam over the gold-coated (⬃20 –30 nm thick) sample, with secondary ions being accelerated normal to the sample surface. Negative ions from a 16 ␮m diameter circular area in the center of the 100 ⫻ 100 ␮m2 crater were allowed into the mass spectrometer. Positive charge buildup in the crater was alleviated using a normal-incidence electron gun. Secondary ions with 0 ⫾ 20 eV excess kinetic energy were detected. The species H⫺, 12C⫺, 18O⫺, and 26Mg16O⫺ were monitored by peak switching the magnetic field and were detected by an electron multiplier operated in pulse-counting mode. III.

Results and Discussion

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Fig. 2. Weight percentage of MgCO3 formed vs. CO2 pressure for similarly sized Mg(OH)2 single-crystal fragments reacted at 585°C (⫾5°C) for 16 h. (Œ) represent reactions in which autoclave was air-cooled (⬃30 min), while sample was still under CO2, to ambient temperature at end of each run. (f) represent reactions in which autoclave was water-quenched, immediately followed by CO2 evacuation at end of each run. Critical CO2 pressure for MgCO3 formation at 585°C (⬃53.7 atm) is shown.§

(1) Effect of CO2 Pressure on Simultaneous Dehydroxylation and Carbonation The extent of dehydroxylation and carbonation for samples reacted under different CO2 pressures at 585°C for 16 h is given in Table I. These compositions are based on carbon and hydrogen elemental analyses for similarly sized (basal-plane area and thickness), partially reacted, single-crystal fragments. The weight percentages of MgCO3, Mg(OH)2, and MgO present are calculated assuming that all detected carbon and hydrogen are present as MgCO3 and Mg(OH)2, respectively. More explicitly, the hydrogen is due to residual hydroxyl groups, which may be associated with Mg(OH)2 regions or lamellar oxyhydroxide intermediate formation, i.e., Mgx⫹yOx(OH)2y, which has recently been observed during Mg(OH)2 dehydroxylation.23 The weight percentage of carbonate formed and the extent of dehydroxylation (Mg(OH)2 converted to MgO or MgCO3) as a function of reaction pressure are shown in Figs. 2 and 3, respectively. The extent of carbonation increases with increasing CO2

Table I. Phase Composition of Mg(OH)2 Reacted for 16 Hours at 585°C as a Function of CO2 Reaction Pressure

Sample

P(CO2) (atm)

MgCO3 (wt%)

HB1 HB2 HB3 HB4 HB5 HB6 HB7 HB8 HB9 HB10 HB11 HB12 HB13 HB14 HB15 HB16† HB17†

1.4 5.5 10.2 13.6 16.3 20.4 20.6 21.1 27.2 43.6 54.5 64.0 106.2 140.9 183.8 10.2 73.4

5.1 7.3 7.0 5.6 6.8 10.7 20.5 21.6 26.1 26.0 16.5 11.2 4.1 1.8 1.4 1.4 8.3

Composition MgO (wt%)

88.4 82.6 81.7 87.4 81.8 77.1 63.9 62.5 55.3 57.9 20.2 24.7 15.4 7.0 14.7 95.8 44.6

Mg(OH)2 (wt%)

6.5 10.1 11.3 7.0 11.5 12.2 15.6 16.0 18.6 16.0 63.3 64.1 80.5 91.2 83.9 2.9 47.2

† All runs were air-cooled under CO2 to ambient temperature (⬃30 min) at the end of each run, except runs HB16 and HB17, which were water-quenched immediately followed by CO2 evacuation. Critical pressure for MgCO3 formation at 585°C is ⬃53.7 atm.26

Fig. 3. Percentage of dehydroxylation (wt% MgO ⫹ wt% MgCO3) vs. CO2 pressure for similarly sized Mg(OH)2 single-crystal fragments reacted at 585°C (⫾5°C) for 16 h. (Œ) represent reactions in which autoclave was air-cooled (⬃30 min), while sample was still under CO2, to ambient temperature at end of each run. (f) represent reactions in which autoclave was water-quenched, immediately followed by CO2 evacuation at end of each run. Critical CO2 pressure for MgCO3 formation at 585°C (⬃53.7 atm) is shown.§

pressure in the low-pressure region and decreases with increasing CO2 pressure in the high-pressure region, as shown in Fig. 2. This novel behavior suggests two different and intriguing processes: control carbonation in the low- and high-pressure regions. Carbonation is not expected to occur in the low-pressure region, as the carbonate is predicted to be thermodynamically unstable with respect to MgO(s) and CO2(g) at 585°C below ⬃53.7 atm of CO2(g).§ This suggests partially dehydroxylated samples are

§ Critical CO2 pressure for MgCO3, formation/decomposition as function of temperature is estimated from ln P(CO2) vs. 1/T(K) plot of the MgCO3 N MgO ⫹ CO2 decomposition/formation P vs. T data available in Refs. 12 and 16.

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formed without concurrent carbonation at 585°C in the lowpressure region. These samples then exhibit substantially enhanced carbonation reactivity on cooling, once the temperature is low enough for carbonation to occur in the low-pressure region. This behavior is confirmed by sample quenching from 585°C at pressures well into the low- and high-pressure regions. As seen in Fig. 2 and Table I, the extent of sample carbonation drops to just within experimental error of zero for the sample quenched and evacuated in the low-pressure region (10.2 atm CO2). On the other hand, the sample quenched and evacuated from the high-pressure region (73.4 atm) contains 8.3% carbonate, in good agreement with unquenched samples. Thus, carbonation primarily occurs at 585°C in the high-pressure region and during cooling and/or at ambient temperature in the low-pressure region. That the extent of carbonation increases with increasing pressure in the low-pressure region may be associated with (i) the greater temperature range over which carbonation can occur on cooling at higher CO2 pressures and (ii) the increased activity of CO2 present on cooling, when carbonation occurs. XRD analysis of the samples from the low- and high-pressure regions was also consistent with the presence of low- and hightemperature carbonation processes, respectively. Crystalline magnesite was found in the unquenched high-pressure samples, consistent with carbonation occurring at 585°C, where sufficient activation energy is present for crystallization to occur. However, magnesite was not found in the unquenched samples from the low-pressure region, indicating the carbonate formed is likely amorphous, consistent with carbonation occurring at lower temperatures on cooling. The occurrence of relatively rapid carbonation on cooling suggested that the partially dehydroxylated Mg(OH)2 formed at 585°C in the low-pressure region exhibits significantly enhanced carbonation reactivity on cooling. The extent of dehydroxylation observed for the quenched samples generally agrees with the trend observed for unquenched samples, decreasing with increasing pressure, indicating that dehydroxylation, as expected, primarily occurs at 585°C (Fig. 3). A significant decrease in dehydroxylation with increasing CO2 pressure occurs on crossing the critical pressure for MgCO3 formation. This can be attributed to passivating carbonate layer formation hindering the dehydroxylation process, consistent with previous dehydroxylation/carbonation studies performed at atmospheric pressure.12 The general decrease in dehydroxylation with increasing CO2 pressure observed in both the low- and high-pressure regions can be attributed to slower diffusion of H2O(g) away from the dehydroxylating lattice sites as CO2 pressure increases. The slower diffusion rates raise the partial pressure of H2O(g) present locally, slowing the dehydroxylation process. In the high-pressure region, where carbonation can occur at 585°C, surface carbonate formation can also inhibit the dehydroxylation process. However, the extent of dehydroxylation decreases with decreasing extent of carbonate formation (e.g., for potentially thinner passivating carbonate surface layers), suggesting that higher CO2 gas pressures primarily inhibit dehydroxylation by lowering the H2O(g) diffusion rate away from the reaction matrix. The relatively rapid carbonation observed on cooling in the low-pressure region is likely associated with lattice cracking, delamination, and morphologic reconstruction to form nanostructured materials during dehydroxylation at 585°C. This well-known process results in the formation of high-surface-area materials with low residual hydroxide concentrations during relatively slow/lowtemperature Mg(OH)2 dehydroxylation.26 Such high-surface-area materials are likely associated with the substantially enhanced carbonation reactivity observed on cooling in the low-pressure region of Fig. 2. To further explore this possibility, in situ TGA was combined with ex situ carbon and hydrogen elemental analysis to investigate the carbonation reactivity of partially dehydroxylated Mg(OH)2 during cooling. Similarly sized, freshly cleaved, 5 mg single-crystal brucite fragments were used as the starting material. Samples were initially dehydroxylated by heating under dry helium at 2°C/min to 375°C and holding them isothermal for 2 h (reaction position 1 in Fig. 4 and Table II), resulting in

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Fig. 4. In situ TGA studies of carbonation of low-temperature (375°C) calcined Mg(OH)2.23 After stabilizing at ⬃91 wt% dehydroxylation (position 1), atmosphere is changed from dry helium to humid (⬃80% humidity) and dry CO2 for the separate runs at 375°C, as shown. Slight carbonation (⬃2 wt%) occurs initially (by position 2), but does not continue (position 3). Cooling to ambient temperature (position 4) results in little reaction under dry CO2, whereas humid CO2 shows strong carbonation and rehydroxylation behavior, as confirmed by carbon and hydrogen elemental analysis. Slight dip between positions 3 and 4 is associated with instrumental/buoyancy effects.

high-surface-area materials that are ⬃91% dehydroxylated. (Samples similarly preheat-treated to ⬃91% dehydroxylation gave surface areas of 120 m2/g via in situ BET analysis.)26 The reaction gas was then changed to dry or humid (bubbled through a degassed, distilled water bubbler at 20°C) CO2. Both processes resulted in relatively rapid carbonation of ⬃2% of the sample, after which no further carbonation occurred at 375°C. (Similar runs changing from dry to humid helium at 375°C exhibited no change in weight, indicating the weight gain for humid CO2 results from carbonation.) Slight weight losses between positions 2 and 3 (Fig. 4 and Table II) were similarly assigned to dehydroxylation. Cooling to ambient temperature (23°C) under humid CO2 resulted in relatively rapid carbonation and rehydroxylation, while cooling to ambient temperature under dry CO2 resulted in substantially less carbonation. These observations indicated that substantial carbonation could occur during the cooling of, and at ambient temperature for, such partially dehydroxylated Mg(OH)2 materials, especially in the presence of H2O(g). In this regard, it is noteworthy that the dehydroxylation/carbonation studies described in Figs. 2 and 3 were conducted in a sealed autoclave. This allowed the H2O(g) formed during dehydroxylation to remain in contact with the samples during cooling, contributing to the substantial carbonation rates observed during cooling and/or at ambient temperature.

Table II. Sample Compositions for Reaction Positions Shown in Fig. 4 Composition (wt%) Humid CO2†

Dry CO2†

Reaction position

MgCO3

MgO

Mg(OH)2

MgCO3

MgO

Mg(OH)2

1 2 3 4‡

0.0 2.2 2.2 14.7

90.2 88.1 88.9 66.2

9.8 9.7 8.9 19.1

0.0 1.8 1.8 2.7

91.0 89.2 89.5 88.6

9.0 9.0 8.7 8.7

† Compositions are determined by weight change for positions 1–3, as described in text. ‡Final compositions for humid CO2 are determined by total carbon and hydrogen analysis, and those for dry CO2 are determined by TGA, assuming only carbonation occurs. Humid compositions agree well with final weight gain observed on cooling (between position 3 and 4; 7.8% TGA vs. 8.0% estimated from carbon and hydrogen elemental analysis). Dry compositions agree well with those from carbon and hydrogen analysis (2.4% MgCO3, 92.7% MgO, 4.9% Mg(OH)2). Slight increase in wt% of MgCO3 between positions 3 and 4 for dry CO2 run may be associated with traces of H2O left in TGA system during dehydroxylation.

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The above reaction processes follow from the formation of high-surface-area material, which is 91% dehydroxylated, at 375°C under helium. Changing the gas to CO2 at 375°C, under dry or humid conditions, results in only ⬃2% carbonation, after which the available reactive carbonation sites have been consumed. However, during cooling to ambient temperature, carbonation reactivity dramatically increases in the presence of humid CO2, even though the activation energy available to support the reaction is continuously decreasing. This indicates that new carbonation reactive intermediate reaction sites are forming on cooling. Because carbonation reactivity substantially increases on cooling in the presence of H2O(g) and rehydroxylation is observed to occur together with carbonation, it is likely that these new intermediate reaction sites are associated with rehydroxylation of the reaction matrix. (Similarly dehydroxylated, high-surface-area Mg(OH)2 materials have been observed to completely rehydroxylate to Mg(OH)2 under similar, but CO2-free, conditions.)26 Furthermore, we have recently discovered the formation of lamellar oxyhydroxide intermediate materials during Mg(OH)2 dehydroxylation,23 which may form during rehydroxylation as well. These intermediates have only slightly higher free energies of formation (⬃1–2 kcal/mol) than stoichiometrically equivalent amounts of MgO and Mg(OH)2. They likely play a significant role in the formation of new, more reactive, carbonation reaction sites during dehydroxylation and may play a similar role in the dramatically enhanced carbonation reactivity observed during rehydroxylation as well. Surface-adsorbed H2O that can form during cooling may also contribute to the enhanced carbonation reactivity observed, both as an avenue for rehydroxylation intermediate formation and for facilitating CO2 interactions with the reaction matrix surface. In the high-pressure region of Fig. 2, where carbonation can occur at 585°C, both the extent of carbonation and dehydroxylation were found to decrease with increasing CO2 pressure. As discussed above, we expect the rate of dehydroxylation to decrease because of the decreasing diffusion rate of H2O(g) away from the reaction matrix surface with increasing CO2 pressure, with passivating surface carbonate formation also affecting the process. The dramatic decrease in carbonation reactivity with increasing CO2 pressure in the high-pressure region similarly follows. Passivating carbonate layer formation inhibits both dehydroxylation and carbonation processes. However, it is not clear that this effect should increase with increasing CO2 pressure, as the extent of carbonation is observed to decrease with increasing CO2 reaction pressure. On the other hand, because increasing CO2 pressure decreases the diffusion rate of H2O(g) away from the cracks and pores formed during dehydroxylation, it also decreases the rate at which CO2 can diffuse to the necessary carbonation reaction sites, slowing carbonation. In addition, slowing the dehydroxylation process should slow the rate of reactive carbonation site formation, similarly contributing to the observed decrease in carbonation reactivity with increasing CO2 pressure.

(2) Effect of Temperature on Carbonation and Dehydroxylation Reactions The effect of temperature on the dehydroxylation/carbonation process is described in Fig. 5 for a CO2 pressure of 25.2 atm and a reaction time of 4 h. The resulting amounts of MgO, MgCO3, and Mg(OH)2 are plotted versus temperature. At this pressure, the critical temperature for MgCO3 formation is ⬃537°C, above which the carbonate is thermodynamically unstable with respect to MgO ⫹ CO2.§ Below the critical temperature, where carbonation can occur directly at reaction temperature, an exponential increase in the extent of carbonation, dehydroxylation, and MgO formation is observed with increasing temperature. Above the critical temperature, dehydroxylation is expected to be the primary process occurring at reaction temperature, with carbonation occurring during cooling and/or at ambient temperature, as discussed above. The maximum extent of carbonation is observed just above the critical temperature, consistent with relatively rapid carbonation (together with some potential rehydroxylation) occurring during

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Fig. 5. Weight percentage of MgO, MgCO3, and Mg(OH)2 present in similarly sized Mg(OH)2 single-crystal fragments reacted as function of temperature for 4 h under 25.2 atm CO2. Note: Critical temperature for MgCO3 formation under 25.2 atm CO2 is ⬃537°C.§

cooling and/or at ambient temperature, as observed in the lowpressure region of Fig. 2. Further increasing the reaction temperature increases the extent of MgO formation, as expected. However, the extent of carbonation decreases, indicating the carbonation reactivity of the dehydroxylated material on cooling has significantly decreased. This can be attributed to enhanced sintering of the highly dehydroxylated material formed at higher temperatures. Such sintering should decrease the high surface area of the intermediate dehydroxylated material formed, lowering both its carbonation and rehydroxylation reactivity on cooling. This is consistent with previous observations that dehydroxylated Mg(OH)2 materials prepared at higher temperatures exhibit lower rehydroxylation reactivity.26 (3) Dehydroxylation/Carbonation Reaction Progression as a Function of Distance from the Brucite Basal-Plane Surface The extent of simultaneous dehydroxylation and carbonation (above the critical pressure for carbonate formation) is followed as a function of distance from a single-crystal brucite basal-plane surface using SIMS and RBS. In the SIMS analysis, the reaction progression in a partially carbonated crystal fragment is followed to a depth of 4 ␮m, as shown in Fig. 6. The analyzed region is chosen to be crack-free via optical microscopic observation to minimize the effects of cracking on the observed sample interface. At ⬃0.5 ␮m into the sample, the SIMS signal stabilizes, revealing a region ⬃1.3 ␮m thick with relatively constant H/Mg, C/Mg, and O/Mg levels. This region is likely associated with passivating carbonate layer formation, based on the constant relatively high C/Mg and constant relatively low H/Mg signals observed. At a depth of ⬃2 ␮m, C/Mg begins to decrease, while H/Mg begins to increase, with depth from the crystal surface. This indicates the presence of a relatively sharp reaction front between the highly carbonated near-surface region and the largely unreacted crystal interior. This observation provides the first direct evidence of passivating carbonate layer formation, consistent with previous observations that carbonate formation can substantially inhibit the dehydroxylation process.12 However, such passivating layers are expected to be limited in their ability to inhibit the dehydroxylation/carbonation process. As the reaction progresses into the crystal interior, translamellar cracking, delamination, and the morphologic reconstruction of highly dehydroxylated regions associated with dehydroxylation can occur concurrent with carbonation. This provides new pathways for H2O(g) and CO2(g) diffusion away from and toward the reaction matrix, respectively, allowing the reaction to continue by circumventing passivating

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tion of MgO and MgCO3 decrease and the concentration of Mg(OH)2 increases with increasing depth. Two additional features draw particular attention. The first is the dramatic decrease in the extent of dehydroxylation with depth. As discussed above, this can be associated with H2O(g) escaping through the basal-plane surface as translamellar cracking, delamination, and morphologic reconstruction penetrate deeper into the sample as the reaction progresses. The second feature is the carbonate/oxide ratio, which is significantly greater in the crystal interior. This can be associated with the higher partial pressures of H2O(g) expected inside the reaction matrix, because of the longer diffusion paths needed for H2O(g) to escape the matrix. Such increased H2O(g) partial pressures inhibit dehydroxylation locally. Thus, the greater carbonate/oxide ratio observed in the crystal interior may be associated with the slower formation and longer life of local intermediate materials/reaction sites (e.g., oxyhydroxides) with higher carbonation reactivity. This allows more time for the carbonation of dehydroxylation intermediates to occur, before other competitive processes (e.g., further dehydroxylation and MgO sintering) can decrease their carbonation reactivity. Fig. 6. SIMS of extent of dehydroxylation and carbonation as function of depth into brucite single-crystal fragment partially reacted above critical carbonation pressure. H/Mg, C/Mg, and O/Mg signals are shown as function of depth from original brucite basal-plane surface. Sample was gold-coated, with overall crater depth being determined by surface profilometry. Carbon and hydrogen elemental analysis gave overall composition of 6.2 wt% MgCO3, 17.2 wt% MgO, and 76.6 wt% Mg(OH)2 for similarly sized and reacted single-crystal fragments.

carbonate layers. As SIMS analysis progresses into the crystal below a depth of 2 ␮m, the H/Mg and C/Mg signals simultaneously increase and decrease, respectively. Qualitatively, the H/Mg signal appears to increase with depth relatively faster than the C/Mg signal decreases over the same region. This suggests the presence of uncarbonated MgO material in the reaction interface together with MgCO3. The relatively constant O/Mg ratio across the reaction interface then follows from the O/Mg ratios for MgO, MgCO3, and Mg(OH)2, as there should be little change in going from a region containing similar amounts of MgCO3 and MgO into a region that primarily contains Mg(OH)2. This is generally consistent with our carbonation studies as well (e.g., Table I), which always show higher levels of dehydroxylation than carbonation for partially carbonated materials, indicating dehydroxylation generally precedes carbonation locally. Ion beam analysis was used to follow the simultaneous dehydroxylation/carbonation process deeper into more extensively reacted crystals. Crystal fragments were selected, partially carbonated, and mounted on glass slides. The elemental compositions of the outer few micrometers of the partially carbonated crystal were then analyzed. After analysis, the analyzed upper part of the crystal was cleaved using adhesive tape (removing ⬃0.05 mm of the partially carbonated crystal), the freshly exposed surface analyzed, and the procedure repeated. The results for a representative sample are given in Table III. The resulting analyses show the concentra-

Table III. Weight Percentage of Mg(OH)2, MgCO3, and MgO as a Function of Depth, from the Basal-Plane Surface of a Typical Brucite Single-Crystal Fragment Reacted above the Critical Carbonation Pressure† Component

Mg(OH)2 MgCO3 MgO MgCO3/MgO MgO/Mg(OH)2 MgCO3/Mg(OH)2

Content, as a function of depth (wt%) 0 mm ⬃0.05 mm ⬃0.10 mm

27 50 23 2.2 0.9 1.9

50 40 10 4.0 0.2 0.8

70 24 6 4.0 0.1 0.3

† Concentratiions are based on ion beam analysis of the total amount of magnesium, carbon, oxygen, and hydrogen present at each depth.

IV.

Conclusions

Mg(OH)2 gas-phase carbonation is governed by several chemical, physical, and morphologic factors. First, CO2 pressure must exceed the critical pressure needed for MgCO3 formation at the reaction temperature for direct carbonation to occur. Further increasing CO2 pressure above the critical pressure slows both carbonation and dehydroxylation processes. Hence, under the reaction conditions observed, carbonation reactivity of Mg(OH)2 is greatest at the minimum pressure required thermodynamically for MgCO3 formation. At constant CO2 pressure, increasing the reaction temperature enhances carbonation reactivity until MgCO3 becomes thermodynamically unstable with respect to MgO and CO2. More generally, carbonate passivating layer formation can locally form physical barriers that inhibit further dehydroxylation and carbonation during simultaneous dehydroxylation/carbonation reaction processes. The effect of such passivating layers is countered by translamellar cracking, delamination, and the morphologic reconstruction associated with the dehydroxylation process. Translamellar cracking, delamination, and morphologic reconstruction, after extensive local dehydroxylation, are well-known to be associated with low-temperature Mg(OH)2 dehydroxylation, with the potential to produce materials with high surface areas (e.g., 100 –200 m2/g). Such dehydroxylation provides access to materials with excellent gas–solid contact areas to stimulate subsequent gas-phase carbonation processes. However, high surface area alone does not guarantee high carbonation reactivity. High concentrations of reactive sites for carbonation must be present as well. Transitory intermediate materials that form during Mg(OH)2 dehydroxylation/rehydroxylation can exhibit substantially enhanced carbonation reactivity. The ability of such intermediate formation to enhance carbonation reactivity is evidenced by dramatically enhanced gas-phase carbonation reactivity at ambient CO2 pressure and temperature during rehydroxylation of 91% dehydroxylated Mg(OH)2. Although the concentration of carbonation reactive intermediate sites may be low at any given intermediate hydroxide composition, simultaneous dehydroxylation/rehydroxylation processes can provide continuous access to fresh, carbonation-reactive, intermediate reaction sites. In this regard, Mg(OH)2 dehydroxylation has recently been discovered to be governed by lamellar nucleation and growth, which can result in a solid solution series of lamellar oxyhydroxide intermediate materials.23 Such intermediates provide a broad new range of potential carbonation reaction pathways, as shown in Fig. 7, which may generally enhance carbonation reactivity during dehydroxylation and contribute to the substantially enhanced carbonation reactivity observed during rehydroxylation. Integrating enhanced intermediate reactivity with the ability to form high-surface-area materials during dehydroxylation via translamellar cracking, delamination, and morphologic lattice reconstruction forming nanostructured

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Journal of the American Ceramic Society—Be´ arat et al.

Vol. 85, No. 4

Fig. 7. Possible Mg(OH)2 carbonation reaction pathways during dehydroxylation and/or rehydroxylation. Intermediate lamellar oxyhydroxide solid solution series, Mgx⫹yOx(OH)2y, is represented by nominal compositions of Mg3O(OH)4 and Mg3O2(OH)2.23 The light gray, dark gray, white, and black spheres correspond to magnesium, oxygen, hydrogen, and carbon atom positions, respectively. Whereas MgO and Mg(OH)2 can form by dehydroxylation and rehydroxylation, respectively (potentially cycling back and forth at low temperatures),26 carbonate formation is thermodynamically dictated as a one-way reaction (below MgCO3 decomposition temperature).

materials further extends this potential. The prototypical lamellar hydroxide mineral character of Mg(OH)2 suggests some of the above mechanisms may also apply to the carbonation of other magnesium-rich minerals that contain hydroxide lamella, such as the serpentines. An avenue of future research interest is the potential for low-level impurities to enhance mineral carbonation reactivity. Such enhanced reactivity has been previously observed for the reaction of CaO with SO2 in the presence of low levels of Fe2O3.27 Acknowledgments We thank the Center for Solid State Science for use of the Goldwater Materials Science Laboratories, including the Materials Facility, the Ion Beam Analysis Facility, the Secondary Ion Mass Spectrometry Facility, and the Goldwater Materials Visualization Facility. We thank the Department of Chemistry and Biochemistry for the use of the X-ray Diffraction Facility. We also wish to thank Dr. M. Mercer, of Hiden, Inc., for assistance with BET surface area determination, and Professor George Wolf for helpful discussions.

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