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Experimental study of Mg-rich silicates carbonation at 400 and 500 °C and 1 kbar
Chemical Geology 265 (2009) 79–87
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Chemical Geology j o u r n a l h o m e p a g e : w w w. e l s e v i e r. c o m / l o c a t e / c h e m g e o
Experimental study of Mg-rich silicates carbonation at 400 and 500 °C and 1 kbar F. Dufaud a, I. Martinez a,⁎, S. Shilobreeva a,b a b
Lab de Géochimie des isotopes stables et Centre de Recherches sur le Stockage Géologique du CO2, IPGP/TOTAL/SCHLUMBERGER/ADEME, 4 place Jussieu, 75252 Paris cedex 05, France Vernadsky Institute of Geochemistry and Analytical Chemistry of Russian Academy of Sciences, Moscow, Russia
a r t i c l e
i n f o
Article history: Received 6 February 2008 Received in revised form 29 January 2009 Accepted 30 January 2009 Editor: J. Fein Keywords: CO2 sequestration Mineral carbonation Ultrabasic rocks Carbonates
a b s t r a c t The reactivity of olivines, orthopyroxenes and serpentines (chrysotile) with CO 2 rich fluids was experimentally studied at 400–500 °C and 1 kbar (+/− NaCl). Levels of formation of solid carbonate phases were measured both by analysing the recovered gas phase and by step heating of the solid samples. Carbonation levels of several percents per hour (from 3 to 57% in 4 h depending upon experimental conditions) were measured with increasing efficiency in the order: orthopyroxene, chrysotile, olivine. In the case of olivine, a positive impact of water fugacity and salinity on carbonation levels was evidenced. Microstructures of the samples showed that the carbonation reaction proceeded by dissolution/precipitation. A coupling between solid carbonate production, mostly as magnesite, and olivine serpentinisation was demonstrated. Isotopic compositions of carbon in gas phases and carbonates were measured. They were found to be consistent with mass balance calculations, additionally suggesting the existence of an accessory carbon-bearing phase carrying negative 13C signatures. This phase was identified by transmission electron microscopy, as an ill ordered graphite phase, which would account for about 15% of carbon mineral under these conditions. © 2009 Elsevier B.V. All rights reserved.
1. Introduction With the increasing concentration of carbon dioxide in the atmosphere and the likely associated global climate change, there has been an increasing interest in carbon dioxide sequestration (e.g. Holloway, 1997a; Bachu, 2000; Oelkers and Schott, 2005), which consists of capturing and injecting CO2 into a target geological formation. Several geological formations are envisaged as possible options to store CO2: deep sedimentary formations and saline aquifers (e.g. Bachu and Adams, 2003; Kaszuba et al., 2003), depleted oil or gas reservoirs (e.g. Holloway, 1997b; Wildenborg and Lokhorst, 2005), basalts (e.g. Matter et al., 2007; McGrail et al., 2006) and ultrabasic rocks (Goff and Lackner, 1998). The main advantage of CO2 sequestration in basic and ultrabasic rocks is their high reactivity with CO2-rich fluids thus providing divalent cations such as Ca2+, Mg2+ and Fe2+, which may combine with the injected CO2 to form carbonate minerals leading to a very stable sequestration of CO2. Such a good reactivity is evidenced in basalt weathering which plays a major role on exchanges of CO2 between the solid Earth and the atmosphere. For instance, Dessert et al. (2003) showed that the annual CO2 consumption through basalt weathering represents 30–35% of the total flux of Na, K, Ca and Mg associated to silicate weathering (Gaillardet et al., 1999), although those rocks represent only 8% of the total rocks exposed at the Earth surface. Previous experimental studies investigated basalts dissolution levels and calcite precipitation both in laboratory experiments (e.g. Oelkers ⁎ Corresponding author. Tel.: +33 1 44 27 60 90; fax: +33 1 44 27 28 30. E-mail address: [email protected] (I. Martinez). 0009-2541/$ – see front matter © 2009 Elsevier B.V. All rights reserved. doi:10.1016/j.chemgeo.2009.01.026
and Gislason, 2001; McGrail et al., 2006) and injection sites (Matter et al., 2007) showing that water–rock interactions in these settings are in favour of permanent storage of CO2. Forsterite dissolution and magnesite precipitation under high partial pressure of CO2 (pCO2 = 1bar and 100 bar) at 30 and 95 °C were recently studied in batch reactors by Giammar et al. (2005) showing that nucleation of magnesite, requiring large critical saturation index, was a limiting factor for the global reaction. Experiments on olivine exposed to high CO2 pressure (185 bar) and high temperature (155 °C) showed that, under these conditions, a 78% carbonation level of the silicate was achieved after 30 min (O'Connor et al., 2002). Such experiments intended to demonstrate the feasibility of industrial processes of ex situ carbonation of ultrabasic minerals. In order to adequately simulate CO2 sequestration in basic and ultrabasic rocks, knowledge of the mechanisms, kinetics and thermodynamics of carbonate forming reactions with individual major mineral of these rocks is needed. In the present work, we chose to study the interactions between a supercritical H2O–CO2 fluid and three Mg-rich minerals: orthopyroxene, chrysotile and olivine, using both mineralogical and geochemical tools. For providing data on the high pressure–high temperature side of the phenomena, the total pressure (1 kbar) and temperature (400 °C and 500 °C) conditions were more elevated than in most envisioned CO2 storage projects or previous experiments. This has the interest to simulate high temperature events which might affect a CO2-storage site in basic and ultrabasic rocks on long time scales, as well as to provide interesting kinetic and mechanistic data associated to large carbonation levels.
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2. Materials and methods
Table 2 Experimental conditions.
2.1. Starting materials
Sample
Most of the experiments were carried out on (Mg,Fe)2SiO4 olivines which were extracted and separated from a peridotite from San Carlos, Arizona. In order to investigate the effect of the nature of the Mg-bearing mineral, (Mg,Fe)SiO3 orthopyroxene from the same San Carlos peridotite and Mg3Si2O5(OH)4 chrysotile collected in centimetric veins from Phillips Mine, Salt River Canyon (Arizona), were also used. Chemical composition of these starting materials was determined using electron microprobe and is reported in Table 1. The starting materials were also checked by powder X-ray diffraction and found to be pure within the precision of the X-ray diffraction determination. The minerals were crushed into powders with grain sizes ranging from 100 to 200 μm. Oxalic acid (H2C2O4) was the source of CO2. Three different batches of oxalic acid were used; two of them were anhydrous with δ13C = −8.20‰ and δ13C = −33.09‰, respectively; one was di-hydrated with an isotopic composition of δ13C = −20.08‰. When the temperature in the experiment exceeded approximatively 150 °C, depending upon pressure, oxalic acid decomposed through the following reactions (Holloway et al., 1968; Javoy et al., 1978): H2 C2 O4 →CO2 þ CO þ H2 O
ð1Þ
CO þ H2 O↔CO2 þ H2
ð2Þ
The complete decomposition can be written: H2 C2 O4 →2CO2 þ H2
ð3Þ
In the case of di-hydrated oxalic acid, the global decomposition equation is: H2 C2 O4 d2H2 O→2CO2 þ H2 þ 2H2 O
ð4Þ
At relatively low temperatures, complex decomposition schemes are occurring with formation of formic acid (Surdam et al., 1984), but they probably do not affect the high-temperature decomposition process schematized by reactions (3) and (4). 2.2. Experimental procedure Powdered samples (typically 100 mg) were loaded in a 4 cm-long platinum capsule (wall thickness of 0.2 mm) together with oxalic acid (typically 10 mg). In some runs, water or NaCl-rich water was added (2 to 5 mg, Table 2). Large salinities chosen in Seq 33 and Seq 37 (Table 2, 100 g l− 1 NaCl) indeed correspond to average pore aqueous
Table 1 Chemical compositions (wt.%) of silicates used in the high-pressure experiments. Olivine
Chrysotile
Orthopyroxene
n
17
8
17
SiO2 TiO2 Al2O3 Cr2O3 FeO MnO MgO CaO Na2O K2O NiO Estimated H2O Total
40.92 ± 0.54 0.02 ± 0.02 0.02 ± 0.02 0.03 ± 0.03 8.97 ± 0.24 0.13 ± 0.04 49.64 ± 0.39 0.08 ± 0.03 0.02 ± 0.02 0.01 ± 0.01 0.37 ± 0.08 0 100.21 ± 0.70
43.82 ± 0.76 0 0.10 ± 0.02 0.01 ± 0.01 1.05 ± 0.04 0 38.70 ± 0.56 0.06 ± 0.02 0.01 ± 0.01 0.00 ± 0.00 0.05 ± 0.06 16.16 83.84 ± 1.03
54.98 ± 0.70 0.05 ± 0.04 3.21 ± 0.06 0.58 ± 0.05 5.55 ± 0.18 0.11 ± 0.07 33.53 ± 0.80 1.06 ± 0.8 0.04 ± 0.03 0.01 ± 0.01 0.12 ± 0.07 0 99.27 ± 0.97
n: number of analysis.
Seq 9 Seq 10 Seq 11 Seq 13 Seq 14 Seq 15 Seq 16 Seq 31 Seq 32 Seq 33 Seq 34 Seq 35 Seq 36 Seq 37 Seq 38
T (°C) 400 400 400 500 500 500 500 500 500 500 500 400 400 400 400
Mineral Orthopyroxene Orthopyroxene Chrysotile Olivine Olivine Orthopyroxene Chrysotile Olivine Olivine Olivine Olivine Olivine Olivine Olivine Olivine
Silicate
Oxalic acid
Water
mg
mg
mg
200.19 200.50 195.77 100.30 100.10 100.43 100.38 100.30 101.22 100.02 49.80 100.80 99.15 99.90 48.57
19.80 (1) 19.99 (1) 20.27 (1) 10.69 (1) 10.05 (1) 10.79 (1) 10.91 (1) 10.34 (2) 9.84 (3) 10.25 (2) 5.27 (2) 10.23 (2) 10.35 (3) 10.75 (2) 5.02 (2)
0.00 2.14 0.00 0.00 1.58 0.00 0.00 0.00 0.00 4.59 3.44 0.00 0.00 3.70 3.04
All experiments were conducted at 1 kbar during 4 h. Isotopic composition of initial oxalic acid was: (1) δ13C = − 8.2‰; (2) δ13C = −33.09‰ and (3) di-hydrated oxalic acid with δ13C = −20.08‰. In experiments Seq 33 and Seq 37, a NaCl-rich water (with a concentration of 100 g l− 1 of NaCl) was added to the experimental charge.
phases in some sedimentary basins (e.g. Hanor, 1994). The proportions of silicate, oxalic acid and water loaded in the capsule varied in order to investigate different water/solid ratios as well as different water/CO2 ratios. The capsule was welded on both ends and was observed carefully under a binocular to check for leakages. It should be emphasized that the platinum used as capsule material allows some H2 diffusion through the capsule walls and so that, reaction (2) is continuously displaced to the right. Four capsules per run were then loaded in an internally heated Ar pressure vessel, described in more details in Jendrzejewski et al. (1997). In a first step, the pressure was increased at room temperature up to 60% of the desired pressure; then temperature was increased up to 400 °C or 500 °C thus raising the pressure to its peak value. The duration of the experiments was 4 h. This duration was chosen as an empirical compromise between (1) allowing enough time for having significant advancement of the reaction necessary to investigate the carbonation by gas recovery and carbon isotopes and (2) not using too long durations which would correspond to a complete loss of hydrogen through the capsule walls, leading to more oxidized conditions than desired. Experiments were quenched by turning off the internal heater, thus achieving an approximately isobaric cooling (about 50 bar of pressure loss). Argon pressure was then released down to atmospheric pressure in less than 5 min. All capsules were weighed before and after experiments; the precision in weighing did not allow, however, to check for H2 losses. The analysis of the residual gas phase, presented later, showed that in most experiments only small amounts of CO and H2 were found in the capsule, meaning that reaction (2) had been completed to a finite advancement allowing to compute an effective oxygen fugacity and that large H2 loss occurred. The temperature conditions and the initial amount of the reactants are reported in Table 2. 2.3. Analyses After the experiments, the gaseous phases were first recovered by perforation of the capsule under vacuum; the liberated gases entered into an extraction line, used for carbon extraction, which allowed purification and quantification of the main gaseous components (CO2, CO, H2O and H2, see Pineau et al., 1976 for more details). Incondensable gases such as CO and H2 were first separated cryogenically from the condensable gases (CO2 and H2O) trapped into a liquid nitrogen reservoir. Step-heating of this trap allowed liberation of CO2 at approximately −145 °C and of water between −100 and −50 °C. Subsequent quantification of pure CO2 and pure water was then performed using a Toepler pump. The oxidation of the remaining CO, by passing over CuO furnace at 450 °C, allowed its transformation
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Table 3 Composition of the gases (in micromoles) present in the capsule after experiments and volatiles extracted from the solid after pyrolysis, normalized to 1 mg of solid. Sample Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq
9 10 11 13 14 15 16 31 32 33 34 35 36 37 38
Gas phase
Solid phase
nCO2(g)
nH2O(g)
nCO(g)
nH2(g)
nCO2(pyrolysis)
nH2O(pyrolysis)
381 (4) 387 (4) 384 (4) 167 (4) 145 (4) 224 (4) 194 (4) 193 (4) 128 (2) 138 (2) 87 (2) 171 (4) 111 (2) 101 (2) n.a.
4.8 (2) 111 (2) 68 (2) 49 (2) 75 (2) 81 (2) 133 (2) 8.6 (2) 86 (2) 196 (4) 198 (4) 2.8 (2) 121 (2) 102 (2) n.a.
45.3 (2) n.a. 13.5 (2) 0.5 (2) 0.3 (2) 1.0 (2) n.a. 8.2 (2) 2.7 (2) n.a. 0.6 (2) 3.2 (2) 3.0 (2) 2.1 (2) n.a.
39.3 (2) n.a. 36.7 (2) 23.2 (2) 0 16.8 (2) n.a. 13.2 (2) 4.7 (2) n.a. 2.8 (2) 8.4 (2) 2.0 (2) 4.6 (2) n.a.
0.17 (3) 0.23 (4) 0.39 (6) 0.55 (8) 0.75 (11) 0.07 (1) 0.34 (5) 0.27 (4) 0.42 (6) 0.83 (12) 0.17 (3) 0.47 (7) 0.68 (10) 1.17 (18) 0.38 (6)
0.11 (2) 0.13 (2) 6.75 (1.01) 0.11 (2) 0.11 (2) 0.04 (1) 6.56 (98) 0.08 (1) 0.01 (1) 0.18 (3) 0.29 (4) 0.22 (3) 0.14 (2) 0.33 (5) 0.18 (3)
r1
r2
3(1) 5(1) 12(1) 29(1) 38(1) 5(1) 11(1) 13(1) 16(1) 39(1) 25(1) 23(1) 31(1) 57(1)
7.7 (1.2) 10.7 (1.6) 17.6 (2.6) 24.1 (3.6) 35.6 (5.3) 2.80 (40) 14.3 (2.1) 12.1(1.8) 28.0(4.2) 39.1(5.9) 6.7(1.0) 21.4(3.2) 43.7(6.6) 54.4(8.2)
r1 and r2 (in %) are the estimated mole fraction of CO2 sequestered under the form of carbonate calculated from the analysis of the gas phase (r1) and from the CO2 obtained by pyrolysis of the solid (r2). n.a.: non analyzed.
into CO2, which was then trapped into liquid nitrogen and manometrically quantified using a Toepler system, yielding a precision of 5% on the volume. For carbon, the measured blanks for this procedure were 0.2 μmol, thus specifying the detection level. After such procedure, the residual gas is H2, which was also quantified using the Toepler pump. CO and CO2 were collected in glass tubes for further carbon isotopic measurements. Once the gas phases were analysed, the solid phases were recovered. Approximately 25 mg of each of these solids were used for estimating the carbon contents with the step-heating technique. For this purpose, the sample powders were loaded in glass tubes and heated slowly up to 900 °C in the extraction line. During their pyrolysis, solid carbonates were decomposed both into CO2 and CO, which was oxidized using a CuO furnace heated at 450 °C. Gaseous CO2 was continuously collected in a liquid nitrogen trap. After 30 min, the temperature was increased up to 1000 °C in order to burn any residual carbonates. The collected CO2, produced by decarbonation of the samples, was then quantified and analysed for its isotopic composition using a Finnigan Mat Delta E double collecting mass spectrometer. Water produced during this pyrolysis was also collected and quantified in the same way as CO2. Finally, few mg of the solid powder were mounted in epoxy resin and polished for mineralogical studies using Scanning Electron Microscopy (SEM), performed on a JEOL JSM 840A, equipped with an Energy Dispersive System PGT Sahara detector for semi quantitative analyses, working at 15–20 kV and 1–2 nA. Additional analyses were carried out by Analytical Transmission Electron Microscopy (ATEM) on finely ground powders resuspended in ethanol and let to evaporate on carbon coated copper grids. The TEM used was a JEOL 2100 FEG, operated at 200 kV, equipped with an energy dispersive X-ray analysis Si(Li) diode system from JEOL. 3. Results 3.1. Analyses of gas phases Compositions of the gas phases recovered from each experiment are reported in Table 3. They consist mainly of CO2, H2O, CO and H2 in various proportions depending upon silicate starting material, initial water content and salinity. In most experiments, however, only minor amounts of CO and H2 were analysed. The proportions of each gas were different from those predicted from simple oxalic acid decomposition (for example, considering 10 mg of oxalic acid decomposing through reaction (3), 222 μmol of CO2 and 111 μmol of H2 should be formed) because these gases were involved in several reactions that will be discussed later in the paper. Taking into account the small amount of CO measured, reaction
(2) was displaced to the right, thus leading to the formation of CO2 and H2. The small measured quantity of H2 relative to CO2, however, implies a large H2 diffusion through the capsule walls. Nevertheless, the presence of some CO and H2 in the end products ensures that oxygen fugacity in these experiments were low (see Discussion). Carbon isotopic compositions of CO and CO2 were measured in some experiments and are reported in Table 4 using the usual δ notation with the Pee Dee Belemnite standard. The δ13C values are variable and depend upon the isotopic composition of the source of CO2. The fractionation factor, Δ, which is not depending upon the isotopic composition of the source, is calculated and also given in the table. The calculated Δ values can be compared to theoretical values in order to discuss isotope equilibrium conditions. 3.2. Analyses of solid phases Quantities of CO2 and H2O produced by pyrolysis of aliquots of the solid samples are reported in Table 3. The values are given as normalized to a sampled solid mass of 1 mg. Scattering of analyses carried out on two different aliquots of a same sample suggests that large uncertainties of approximately 30% affect these measurements, due to heterogeneities in the solid phases. As will be noticed below, carbon in the solid phases is mainly present as solid carbonates whereas a small but significant amount is present as poorly ordered graphite. This complexity might contribute to sample heterogeneity. Carbon isotopic compositions of CO2 derived from the carbonates present in these solid samples were measured in selected experiments and are given in Table 4. As for the gases, the δ13C obtained on the carbonates reflects the isotopic composition of their source (i.e. oxalic acid). To avoid this, we calculated the fractionation factor Δ, between CO2 and carbonate, which will be used later in the paper to discuss the isotope equilibrium conditions. Table 4 Carbon isotopic composition of the carbon rich phases, expressed using the usual δ notation with the PDB standard (error bars: 0.1‰) and calculated isotopic fractionations (Δ). Sample T (°C) Mineral Seq 9 Seq 10 Seq 11 Seq 16 Seq 31 Seq 35 Seq 37
400 400 400 500 500 400 400
Orthopyroxene Orthopyroxene Chrysotile Chrysotile Olivine Olivine Olivine
δ13CCO2
δ13CCO
δ13Ccarbonate ΔCO2–carb
‰
‰
‰
n.a. − 4.95 − 4.53 − 6.82 − 32.70 − 30.97 − 29.33
n.a. n.a. − 35.60 n.a. −47.30 −59.70 −60.00
− 9.39 −9.53 −9.38 − 10.51 − 28.90 − 31.18 − 32.46
n.a 4.58 4.85 3.69 − 3.8 0.21 3.13
ΔCO2–CO n.a. n.a. 31.07 n.m 14.6 28.73 30.67
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The carbonation reactions were visualized using SEM. An example of a carbonated olivine powder is shown in Fig. 1. Olivine crystals (50 μm in size; Table 3, experiment Seq 31) were surrounded by reaction rims in which small white crystals are visible. Chemical analyses using the EDS of the SEM showed that these small crystals are Mg–Fe-rich and can be interpreted as being the newly formed iron bearing magnesite grains, surrounded by a silica-rich matrix in dark grey on the figure, as confirmed by TEM investigations (see below). It is interesting to observe in Fig. 1 that some olivine grains are showing two or three distinct reactions rims. The enhancement of carbonation in the presence of water, using di-hydrated oxalic acid (experiment Seq 32, Table 3), is clear from Fig. 2. This figure shows olivine grains surrounded by small magnesite crystals of approximately 2–5 μm in size, dispersed all over the sample. When sodium chloride was added to the experiment, carbonation levels were significantly increased (Table 3, experiments Seq 33 and Seq 37), which is also clear from SEM images. Fig. 3a shows microstructures typical of such samples (Seq 33) in which large globules of magnesite of approximately 10 μm formed during the reaction. EDX analyses of magnesite, small precipitates of NaCl and olivine are given in Fig. 3b. Comparison of carbonate abundance (Table 3) obtained after pyrolysis on samples reacted at 400 °C with 500 °C samples (initial content of water and salinity being equal) shows that carbonates are more abundant at the lowest temperature. For instance, Seq 31 (500 °C) and Seq 35 (400 °C) gave respectively 13(1) and 23(1) % of neo-formed carbonates (where numbers in parenthesis stand for the error bar); Seq 32 (500 °C) and Seq 36 (400 °C) gave 16(1) and 31(1) % of neo-formed carbonates and finally Seq 33 (500 °C) and Seq 37 (400 °C) gave 39(1) and 57(1) % of neo-formed carbonates. The largest amounts of CO2 transformed into solid carbonates were obtained from an olivine reacted at 400 °C with 100 g l− 1 of NaCl-containing water (Seq 37). ATEM studies on reacted powders showed the presence of euhedral magnesite crystals (Fig. 4), identified both by their EDX spectra (not shown here) and electronic diffraction pattern (inset in Fig. 4). Such crystals are present in all run products but not in the initial powders. Microprobe analyses on the largest magnesite crystals showed that they are iron-bearing with average Fe/Fe + Mg value of 10 atomic %, versus a Fe/Fe + Mg of initial olivine of 11 atomic %. Such chemical compositions are in agreement with previous experimental work on San Carlos olivine (O'Connor et al., 2002). In all samples, magnesites were systematically associated with a “proto-serpentine” phase, showing a fibrous microstructure (Fig. 5). This proto-serpentine is the silica rich cement surrounding the small Mg-rich carbonates grains observed by SEM. EDX spectra showed that the proto-serpentine also contained iron, with average
Fig. 1. Backscattered electron image of sample Seq 31 showing large olivine crystals in light grey surrounded by a magnesite (white spots) embedded in a silica rich layer (in dark grey). Several reactions rims are observed around olivine crystals showing that the reaction is surface controlled.
Fig. 2. Backscattered electron image from sample Seq 32. Carbonation reaction is enhanced compared to Seq 31 because of the use of di-hydrated oxalic acid that increases the partial pressure of water in the experiment. Compared to sample Seq 31, magnesite crystals are larger (several microns in size) and dispersed out all over the samples.
Fe/Fe + Mg ratio of 7 atomic %. Determining the exact nature, speciation and oxidation state of iron in the proto-serpentine phase, which might be present as either Fe2+, Fe3+ or both, would require further high resolution analytical TEM studies. Finally, small quantities of poorly crystalline graphite were found in the reacted samples (Fig. 6). When orthopyroxene (Seq 9, 10 and 15) and serpentine (Seq 11 and 16) were the starting materials, smaller amounts of carbonates were formed; secondary proto-serpentine phases were observed as well. It is interesting to note that when NaCl was added to the experiments, the proto-serpentine phases contained some chlorine. 4. Discussion 4.1. Carbonation levels Carbonation levels were calculated using two methods. The first calculation is based on the analyses of the gas phase present in the capsule after the experiment. The difference between the initial number of moles of carbon introduced in the capsule as oxalic acid and the measured CO + CO2 is attributed to the formation of solid carbonates, thus giving the number of moles of CO2 sequestered. This number, reported as a percentage of the total equivalent quantity of CO2, is called r1 in Table 3. The second calculation used the equivalent quantity of CO2 measured by degassing the solid after the experiment, which is attributed to newly formed carbonates. This number, called r2, is reported in Table 3. The uncertainties on r2 are estimated to be of about 30% by comparing the results of pyrolysis experiments carried out on several aliquots from a same experimental run. In this respect, determination of CO2 mineral sequestration levels by analysis of the gases (r1) appears to be more precise in spite of possible changes in composition of the gas phase during quench. Moreover, since all major C-bearing carbon phases were analysed (i.e. CO and CO2), potential problems related to quench are indeed minor. Taking into account uncertainties, it is concluded that the carbonation levels measured by the two methods are mutually consistent except for samples Seq 32, 34 and Seq 36 in which the analysis of the solid by pyrolysis was presumably affected by strong heterogeneities. In all cases, it appears that the formation of solid carbonates is significant after 4 h at these pressure–temperature conditions and that mineral sequestration on the order of tens of percents are obtained. The carbonation levels increase from orthopyroxene, to chrysotile and then to olivine. From
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Fig. 3. (a) Backscattered electron image obtained from sample Seq 33, in which 100 g l− 1 NaCl aqueous phase was added in the capsule. Large spherules of magnesites (approximatively 10 μm in size) were formed all over the sample, in agreement with a precipitation process from a fluid phase. (b) Corresponding energy dispersive X-ray analyses of magnesite, sodium chloride (due to the small size of these precipitates, spectrum is contaminated by olivine) and olivine.
the different experimental conditions (temperature, salinity, dry or wet oxalic acid, water added) investigated with olivine, the largest CO2 mineral sequestration levels were obtained in experiments conducted in saline waters (100 g l− 1 NaCl), in agreement with previous experiments (O'Connor et al., 2002). Fig. 7 summarizes the measured carbonation levels: it appears clearly that carbonation is more efficient at 400 °C than at 500 °C, consistently with the microstructural observations shown above. The simplest way to rationalize this observation is to notice that equilibrium CO2 fugacity above magnesite is larger at 500 °C than at 400 °C and that the amount of CO2 stored as a solid is thus smaller at higher temperatures (see below). The low reactivity of CO2 at high temperature was also evidenced earlier by Bischoff and Rosenbauer (1996) in their high pressure study of rhyolite alteration in CO2 charged water at 200 °C and 350 °C. The greater alteration of the albitic plagioclase was partly explained in their study by the greater acidity of carbonic acid at 200 °C. Possible formation of solid carbonates upon pressure quench cannot be definitely excluded. However, the size of magnesites formed in our experiments (2–5 μm in most NaCl-free experiments and up to 10 μm in sample Seq 33) are large compared to those usually formed upon quench which are usually small euhedral crystals with typical sizes below 1 μm. Moreover, the CO2 fugacities calculated in equilibrium with magnesite at 400 °C and 500 °C (see below) are qualita-
tively in agreement with the measured CO2 pressures; this suggests a high pressure–high temperature equilibrium of CO2 with magnesite. 4.2. Mechanisms Carbonates were observed on silicate surface together with a Sirich layer (Fig. 1). TEM confirmed these observations, showing magnesite systematically associated to a fibrous proto-serpentine phase (Fig. 5). This leads to the conclusion that carbonation is coupled to a serpentinisation reaction. The newly formed magnesite crystals are euhedral and well faceted, as shown in Fig. 4. No crystallographical relationship between these magnesite crystals and the adjacent silicates was observed. Moreover, the microstructures observed by SEM, such as corrosion gulfs in olivine crystals (Fig. 8), strongly suggest a mechanism controlled at the fluid/silicate interface, most probably a dissolution/precipitation reaction such as described below. Dissolution of olivine yields a “proto-serpentine” phase and a magnesium-rich fluid. When carbonate saturation is achieved in the fluid, then magnesite precipitates. The dissolution reaction can be summarized as: 2þ
2Mg2 SiO4 þ 3H2 O þ 2CO2 ↔Mg
−
þ Mg3 Si2 O5 ðOHÞ4 þ 2HCO3
ð5Þ
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F. Dufaud et al. / Chemical Geology 265 (2009) 79–87
Fig. 6. ATEM image of poorly crystalline graphite in sample Seq 33 (associated diffraction pattern in the insert) forming small rounded shape spherules (see arrows), embedded within proto-serpentine fibers.
Fig. 4. Analytical Transmission Electron Microscopy image showing an example of a euhedral crystal of magnesite formed during the experiment Seq 31. The selected area diffraction pattern, indexed using the R-3C space group, is shown in the insert.
using a simplified form for the olivine formula, not taking into account Fe and minor elements, and an ideal stoichiometric chrysotile formula for proto-serpentine. Then the precipitation reaction occurs according to: 2þ
Mg
−
þ 2HCO3 ↔MgCO3 þ H2 O þ CO2
ð6Þ
which by summing Eqs. (5) and (6) leads to: 2Mg2 SiO4 þ 2H2 O þ CO2 ↔MgCO3 þ Mg3 Si2 O5 ðOHÞ4
ð7Þ
In turn, this proto-serpentine phase can be carbonated according to: Mg3 Si2 O5 ðOHÞ4 þ 3CO2 ↔3MgCO3 þ 2H2 O þ 2SiO2
ð8Þ
SiO2 is given here as an ideal end-member product of the reaction. A free silica phase is rarely observed in these experiments. It is likely
Fig. 5. ATEM image of a euhedral magnesite crystal surrounded by proto-serpentine fibers, in sample Seq 37, suggesting a relationship between carbonation and serpentinisation reactions.
Fig. 7. Carbonation levels, expressed as mole percent of sequestered CO2, as a function of temperature (light grey = 500 °C, dark grey = 400 °C) and fluid composition. Note that the highest carbonation level (57 mol%) was found at 400 °C with NaCl aqueous solution added to the capsule.
Fig. 8. Backscattered electron image of sample Seq 38 showing a large olivine crystal with corrosion gulfs suggesting dissolution of the silicate. Magnesites are smaller crystals dispersed in the surrounding matrix.
F. Dufaud et al. / Chemical Geology 265 (2009) 79–87
85
that the excess silicon is rather present in the proto-serpentine which is known for being able to accommodate variable (Mg + Fe)/Si ratios. The sum of Eqs. (7) and (8) leads to the complete carbonation of olivine according to:
Table 6 Molar fractions of CO2 (X), CO (Y), carbonate (Z) and calculated isotopic composition of oxalic acid using Eq. (11) in text. Sample T (°C) Mineral
X
Mg2 SiO4 þ 2CO2 ↔2MgCO3 þ SiO2
Seq 10 Seq 11 Seq 16 Seq 31 Seq 35 Seq 37
0.87 0.85 0.80 0.84 0.75 0.42
ð9Þ
Depending on the relative advancements of the processes (5) to (8), the leading mechanism will be a combination of Eqs. (7) and (8), the proportions of which can be estimated by considering the CO2 and H2O mole numbers obtained after pyrolysis of the solids and given in Table 3. If a and b are the advancement coefficients of Eqs. (7) and (8) respectively, the global reaction can be written as:
400 400 500 500 400 400
Orthopyroxene Chrysotile Chrysotile Olivine Olivine Olivine
(2) (2) (2) (2) (2) (2)
Measured Calculated δ13Coxalic acid δ13Coxalic acid
Y
Z
0.11 (2) 0.03 (2) 0.13 (2) 0.03 (2) 0.02 (2) 0.01 (2)
0.02 (4) −4.49 (8) 0.12 (4) −6.04 (8) 0.07 (4) −6.18 (8) 0.13 (4) −32.74 (8) 0.23 (4) − 31.45 (8) 0.57 (4) − 31.36 (8)
− 8.20 (5) − 8.20 (5) − 8.20 (5) − 33.09 (5) − 33.09 (5) − 33.09 (5)
The numbers in parenthesis stand for the error bars. The measured isotopic composition of oxalic acid is given for comparison.
equilibrium thermodynamics could explain that water has a positive effect on carbonation of olivine, as suggested by reaction (10).
2aMg2 SiO4 þ ð2a−2bÞH2 O þ ða þ 3bÞCO2 ↔ða þ 3bÞMgCO3 þ ða−bÞMg3 Si2 O5 ðOHÞ4 þ 2bSiO2
ð10Þ 4.3. Isotopic measurements
The a and b values determined from the CO2 and H2O mole numbers given in Table 3 are shown in Table 5. From these data, a sequestration level (called r3 in Table 5) can be calculated using the mole number of carbonate, a + 3b, divided by the initial mole number of CO2 introduced. For most samples, sequestration levels r3 are similar to what is calculated with the mole number of CO2 obtained from pyrolysis (r2). This means that both the quantities of CO2 and H2O left in the sample are consistent with the proposed reaction (10) for interpreting the carbonation reaction, and that part of the serpentine formed during this process gave in turn some carbonate by reaction with CO2. As mentioned above, thermodynamic equilibrium is not achieved in the present experiments, especially because of possible continuous H2 loss from the capsules. Although equilibrium (9) is not achieved, it is still possible to obtain qualitative arguments from equilibrium thermodynamics. In particular, using molar volumes, enthalpies and entropies data for quartz, magnesite and forsterite (Holland and Powell, 1990) and assuming SiO2 (as quartz) activity of 1; equilibrium CO2 fugacities for Eq. (9) are 1.9 kbar at 500 °C and 0.26 kbar at 400 °C, respectively. Due to the different approximations and departure from equilibrium, we do not intend to mean that these values are realistic for predicting CO2 fugacities in these experiments, but we notice that these values are actually in the same order of magnitude as CO2 fugacities deduced from the measured CO2 mole numbers in the capsules and that the main temperature effects, i.e. strong increase in equilibrium fCO2 with temperature, will be likely retained in any carbonation process of this type. We thus suggest that carbonation levels are proportional to the difference between initial (after oxalic acid decomposition) and equilibrium CO2 fugacities in the capsule. A CO2 fugacity in equilibrium with magnesite lower at 400 °C than at 500 °C could thus explain faster carbonation rates at the lowest of these two temperatures. Similar correlations between kinetics and
Table 5 Results of a and b advancement coefficients for Eqs. (7) and (8), respectively (see text). Sample
a
b
r3
r2
Seq Seq Seq Seq Seq Seq Seq Seq Seq
0.179 0.229 0.098 0.109 0.275 0.151 0.2 0.223 0.416
0.124 0.174 0.058 0.104 0.185 0.006 0.09 0.153 0.251
24 32.6 11.8 18.5 36.1 7.4 20.6 29.3 50.8
24.1 (3.6) 35.6 (5.3) 12.1 (1.8) 28.0 (4.2) 39.1 (5.9) 6.7 (1.0) 21.4 (3.2) 43.7 (6.6) 54.4 (8.2)
13 14 31 32 33 34 35 36 37
A sequestration level calculated from the mole number of formed carbonates (a + 3b) is calculated and called r3. The numbers in parenthesis stand for the error bar. For comparison, r2, calculated directly from the CO2 obtained by pyrolysis of the solid, is also reported.
Isotopic compositions of CO2, CO and carbonates reported in Table 4 are variable, first of all because the isotopic composition of the carbon source is variable. If we then calculate the fractionation factor (Δ), written for example for ΔCO2–carbonates as the difference δ13CCO2 − δ13Ccarbonates, we can get rid of the variability in isotopic composition of the source. However, the measured values of ΔCO2–carbonates are also variable and do not show connection with isotopic equilibrium values (e.g. 4‰ at 400 °C and 5‰ at 500 °C) as measured by Appora-Gnekindy (1998). On the other hand, values of ΔCO2–CO, although showing also some variability, are close to equilibrium values (25‰ at 400 °C and 17‰ at 500 °C) as calculated in Richet et al. (1977). Indeed, conditions close to isotopic equilibrium are probably achieved in the gas phase, meaning that the quench did not affect the gas mixture. In the case of gas/solid isotope exchange, the heterogeneities might come from the fact that only aliquots of the solid phases were analysed. Moreover, as will be discussed below, the presence of a non analysed isotopically light phase, such as graphite, in various proportions in the solid, may affect strongly the isotopic composition of the companion carbonates. The major implication of these isotopic analyses is that knowing the isotopic composition of the source of carbon, oxalic acid, mass balance calculations can be done as: 13
13
13
13
δ Coxalic acid ¼ Xδ CCO2 þ Yδ CCO þ Zδ Ccarbonate
ð11Þ
where X, Y, and Z are molar fractions of the different phases and δ their isotopic composition. Results of this calculation are given in Table 6. The mass balance is consistent with the other determinations, but its imprecision, due to the aforementioned solid-state heterogeneities, makes it not very useful at this stage. An interesting observation, however, is that these calculations give total carbon isotopic compositions (calculated δ13Coxalic acid in Table 6) that are always less negative than the total carbon isotopic composition actually measured in the starting material (measured δ13Coxalic acid in Table 6). The more likely interpretation of this systematic difference is that a phase with negative δ13C signature is missing from the mass
Table 7 Molar fractions of CO2 (X), CO (Y), carbonate (Z) and reduced carbon (R) obtained from Eq. (12) by isotopic mass balance. Sample
T (°C)
Mineral
X
Y
Z
R
Seq 10 Seq 11 Seq 16 Seq 31 Seq 35 Seq 37
400 400 500 500 400 400
Orthopyroxene Chrysotile Chrysotile Olivine Olivine Olivine
0.68 (4) 0.73 (4) 0.69 (4) 0.79 (4) 0.64 (4) 0.37 (4)
0.08 (4) 0.03 (4) 0.11 (4) 0.03 (4) 0.01 (4) 0.01 (4)
0.02 (4) 0.10 (4) 0.06 (4) 0.12 (4) 0.20 (4) 0.49 (4)
0.22 (4) 0.14 (4) 0.14 (4) 0.06 (4) 0.15 (4) 0.13 (4)
The numbers in parenthesis are the error bars.
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balance. TEM observations show that a “graphite-type” phase was formed in association with carbonates (Fig. 6). Such reduced carbon could be this missing phase since graphite, which is not pyrolized together with carbonates, has usually values of δ13C more negative than carbonates (e.g. Des Marais, 2001). Using ΔCO2–graphite values calculated at this temperature (Bottinga, 1969), a putative molar fraction (called R in the following formula) of this reduced carbon phase can be estimated from: 13
13
13
13
13
δ Coxalic acid ¼ Xδ CCO2 þ Yδ CCO þ Zδ Ccarbonate þ Rδ C“graphite”
ð12Þ
and is provided in Table 7, showing that such an ill-ordered graphite phase could represent up to 22 mol% of the initial carbon injected in these experiments as oxalic acid. Also, the “graphite-type” phase present in the samples could contribute to the large isotopic heterogeneity of carbonates. Various proportions of graphite in the samples, having a very negative isotopic composition, close to the carbonates, will make the isotopic compositions of the carbonates very variable. Only the analysis of both content and isotopic composition of this “graphite-type” phase, which is not easy in such small amounts, will lead to a correct view of the isotopic equilibrium or non-equilibrium in these experiments. This result corroborates the predicted oxygen fugacity in the experiments, estimated using the after-equilibration gas composition from the reaction: CO þ 1=2O2 →CO2
ð13Þ
Using enthalpy and entropy data from Holland and Powell (1990), as well as a polynomial function for CO2 fugacity also from Holland and Powell (1990), assuming ideal behaviour of CO and ideal mixing of gases, average values of oxygen fugacities, fO2 (log units using a 1 bar pressure reference) were −32 ± 0.7 and − 25 ± 0.9 at 400 °C and 500 °C, respectively. These values are also close to that obtained using the empirical formula given in Floden et al. (1998). These fO2 values should be used only as qualitative indicator of the redox conditions of the experiments, not as equilibrium fO2 values because of H2 loss from the platinum capsules preventing equilibrium to be reached. These oxygen fugacities are close to the CCO buffer (C as graphite–CO, recalculated at 1 kbar) (Arculus et al., 1990). Such low values are likely due to the use of oxalic acid as a CO2 source, providing H2 during its decomposition. They can explain why we did not identify magnetite in the run products, which should normally form during serpentinisation reaction, and they are in agreement with the presence of a graphite phase in the solid products. 5. Conclusions Carbonation of powdered olivine, orthopyroxene and chrysotile has been studied at high pressure and high temperatures (1 kbar and 400–500 °C). As analysed quantitatively by post-experiment gas recovery, the fluid obtained by oxalic acid decomposition, complemented in some experiments with water and NaCl, is composed of CO2–H2O in most of the experiments and contains CO and H2 in smaller proportions. Due to possible continous H2 loss from capsules, thermodynamic equilibrium is not achieved in these experiments whose purpose is to provide new kinetic data on silicate carbonation as well as to study methods for measuring advancement of the carbonation process. In such conditions, olivine is the most reactive mineral with 57 mol% carbonation obtained at 400 °C, after 4 h experiment, in CO2 + additional NaCl-rich water. As previously described in literature, the reaction proceeds in two steps: silicate dissolution then followed by precipitation of magnesite. Presence of water and/or NaCl in the fluid is found to favour the reaction. Detailed mineralogical study reveals that serpentinisation of olivine is coupled to the carbonation reaction. Carbon isotopic measurements allowed to make mass balance calculations, which suggest that, under the
reducing conditions of these experiments, a non negligible fraction (up to 22 mol% of the total sequestered carbon) can be found under the form of poorly crystalline graphite, also identified by transmission electron microscopy. Further experiments under such reducing conditions should concentrate on a more precise quantification of this reduced carbon phase, which can also be a potential source for carbon storage, mostly never taken into account in the models, and which is to be related to the abiogenic generation of hydrocarbons at lower temperatures (Berndt et al., 1996). These studies could also help to understand the redox sate of carbon during serpentinisation of oceanic crust. Acknowledgements We are grateful to Omar Boudouma and to François Guyot for their help during SEM and TEM studies of the samples. We particularly thank Françoise Pineau for several discussions on isotopic data. The manuscript was improved by insightful reviews by Olivier Vidal, Bill O'Connor and Pascale Bénézeth. This study was supported by the Centre de Recherches sur le Stockage Géologique du CO2, Institut de Physique du Globe de Paris/TOTAL/ SCHLUMBERGER/ADEME partnership. This is IPGP contribution 2383. References Appora-Gnekindy, I., 1998. Etude expérimentale du fractionnement isotopique du carbone et de l'oxygène dans les systèmes CO2–Carbonates liquides: Application aux contextes carbonatitiques. PhD. Thesis. Université Denis Diderot- Paris 7, IPGP, Paris, 351pp. Arculus, R.J., Holmes, R.D., Powell, R., Richter, K., 1990. Metal–silicate equilibria and core formation. In: Newsom, H.E., Jones, J.H. (Eds.), Origin of the Earth. Oxford University Press, New York, pp. 251–271. Bachu, S., 2000. Sequestration of CO2 in geological media: criteria and approach for site selection in response to climate change. Energy Convers. Manag. 41 (9), 953–970. Bachu, S., Adams, J.J., 2003. Sequestration of CO2 in geological media in response to climate change: capacity of deep saline aquifers to sequester CO2 in solution. Energy Convers. Manag. 44 (20), 3151–3175. Berndt, M.E., Douglas, E.A., Seyfried, W.E., 1996. Reduction of CO2 during serpentinization of olivine at 300 °C and 500 bar. Geology 24 (4), 351–354. Bischoff, J.L., Rosenbauer, R.J., 1996. The alteration of rhyolite in CO2 charged water at 200 and 350 °C: the unreactivity of CO2 at higher temperature. Geochim. Cosmochim. Acta 60 (20), 3859–3867. Bottinga, Y., 1969. Calculated fractionation factors for carbon and hydrogen isotope exchange in the system calcite–carbon dioxide–graphite–methane–hydrogen– water vapor. Geochim. Cosmochim. Acta 33 (1), 49–64. Des Marais, D., 2001. Isotopic evolution of the biogeochemical carbon cycle during the Precambrian. In: Valley, J.W., Cole, D.R. (Eds.), Stable Isotope Geochemistry. Reviews in Mineralogy and Geochemistry, vol. 43. Mineralogical Society of America, pp. 555–578. Dessert, C., Dupre, B., Gaillardet, J., François, L.M., Allegre, C.J., 2003. Basalt weathering laws and the impact of basalt weathering on the global carbon cycle. Chem. Geol. 202 (3–4), 257–273. Floden, A.M., Colson, R.O., Nermoe, M.K.B., Hendrickson, T.R., 1998. Effects of CO on the Activity of Nickel in a Simple Silicate Melt. Lunar Planet. Sci. Conf., XXIX, p. 1520. Gaillardet, J., Dupre, B., Louvat, P., Allegre, C.J., 1999. Global silicate weathering and CO2 consumption levels deduced from the chemistry of large rivers. Chem. Geol. 159 (1–4), 3–30. Giammar, D.E., Bruant, J.J., Robert, G., Peters, C.A., 2005. Forsterite dissolution and magnesite precipitation at conditions relevant for deep saline aquifer storage and sequestration of carbon dioxide. Chem. Geol. 217 (3–4), 257–276. Goff, F., Lackner, K.S., 1998. Carbon dioxide sequestring using ultramafic rocks. Environ. Geosci. 5 (3), 89–101. Hanor, J.C., 1994. Physical and chemical controls on the composition of waters in sedimentary basins. Mar. Pet. Geol. 11 (1), 31–45. Holland, T.J.B., Powell, R., 1990. An enlarged and updated internally consistent thermodynamic dataset with uncertainties and correlations: the system K2O– Na 2O–CaO–MgO–MnO–FeO–Fe2O3–Al2O3–TiO2–SiO2–C–H2–O2. J. Metamorph. Geol. 8, 89–124. Holloway, S., 1997a. An overview of the underground disposal of carbon dioxide. Energy Convers. Manag. 38, 193–198. Holloway, S., 1997b. Safety of the underground disposal of carbon dioxide. Energy Convers. Manag. 38 (Supplement 1), S241–S245. Holloway, J.R., Burnham, W.C., Millhollen, G.L., 1968. Generation of H2O–CO2 mixture for use in hydrothermal experimentation. J. Geophys. Res. 73 (20), 6598–6600. Javoy, M., Pineau, F., Liyama, I., 1978. Experimental determination of the isotopic fractionation between gaseous CO2 and carbon dissolved in tholeitic magma. Contrib. Mineral. Petr 67 (1), 35–39. Jendrzejewski, N., Trull, T.W., Pineau, F., Javoy, M., 1997. Carbon solubility in Mid-Ocean Ridge basaltic melt at low pressures (250–1950 bar). Chem. Geol. 138 (1–2), 81–92.
F. Dufaud et al. / Chemical Geology 265 (2009) 79–87 Kaszuba, J.P., Janecky, D.R., Snow, M.G., 2003. Carbon dioxide reaction processes in a model brine aquifer at 200 °C and 200 bars: implications for geologic sequestration of carbon. Appl. Geochem. 18, 1065–1080. Matter, J.M., Takahashi, T., Goldberg, D., 2007. Experimental evaluation of in situ CO2– water–rock reactions during CO2 injection in basaltic rocks: implications for geological CO2 sequestration. Geochem. Geophys. Geosyst. 8 (2), 1–19. McGrail, B.P., Schaef, H.T., Anita, M., Chien, Y.J., Dooley, J.J., Davidson, C.L., 2006. Potential for carbon dioxide sequestration in flood basalts. J. Geophys. Res. 111 (No B12), B12201. O'Connor, W.K., Dahlin, D.C., Rush, G.E., Dahlin, C.L., Collins, W.K., 2002. Carbon dioxide sequestration by direct mineral carbonatation: process mineralogy of feed and products. Miner. Metall. Process. 19 (2), 95–101. Oelkers, E.H., Gislason, S.R., 2001. The mechanism levels and consequences of basaltic glass dissolution: I. An experimental study of the dissolution levels of basaltic glass as a function of aqueous Al, Si and oxalic acid concentration at 25 °C and pH = 3 and 11. Geochim. Cosmochim. Acta 65 (21), 3671–3681.
87
Oelkers, E.H., Schott, J., 2005. Geochemical aspects of CO2 sequestration. Chem. Geol. 217, 183–186. Pineau, F., Javoy, M., Bottinga, Y., 1976. 13C/12C ratios of rocks and inclusions in popping rocks of the mid-aztlantic ridge and their bearing on the probleme of isotopic composition of deep-seated carbon. Earth Planet. Sci. 29, 413–421. Richet, P., Bottinga, Y., Javoy, M., 1977. A review of hydrogen, carbon, nitrogen, oxygen, sulphur and chlorine stable isotope fractionation among gaseous molecules. Annu. Rev. Earth Planet. Sci. 5, 65–110. Surdam, R.C., Boese, S.W., Crossey, L.J., 1984. The chemistry of secondary porosity. In: McDonald, D.A., Surdam, R.C. (Eds.), Clastic Diagenesis. AAPG Memoir, vol. 37, pp. 127–151. Wildenborg, T., Lokhorst, A., 2005. Introduction on CO2 geological storage. Classification of storage options. Oil Gas Sci. Technol. 60, 513–515.
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Chemical Geology j o u r n a l h o m e p a g e : w w w. e l s e v i e r. c o m / l o c a t e / c h e m g e o
Experimental study of Mg-rich silicates carbonation at 400 and 500 °C and 1 kbar F. Dufaud a, I. Martinez a,⁎, S. Shilobreeva a,b a b
Lab de Géochimie des isotopes stables et Centre de Recherches sur le Stockage Géologique du CO2, IPGP/TOTAL/SCHLUMBERGER/ADEME, 4 place Jussieu, 75252 Paris cedex 05, France Vernadsky Institute of Geochemistry and Analytical Chemistry of Russian Academy of Sciences, Moscow, Russia
a r t i c l e
i n f o
Article history: Received 6 February 2008 Received in revised form 29 January 2009 Accepted 30 January 2009 Editor: J. Fein Keywords: CO2 sequestration Mineral carbonation Ultrabasic rocks Carbonates
a b s t r a c t The reactivity of olivines, orthopyroxenes and serpentines (chrysotile) with CO 2 rich fluids was experimentally studied at 400–500 °C and 1 kbar (+/− NaCl). Levels of formation of solid carbonate phases were measured both by analysing the recovered gas phase and by step heating of the solid samples. Carbonation levels of several percents per hour (from 3 to 57% in 4 h depending upon experimental conditions) were measured with increasing efficiency in the order: orthopyroxene, chrysotile, olivine. In the case of olivine, a positive impact of water fugacity and salinity on carbonation levels was evidenced. Microstructures of the samples showed that the carbonation reaction proceeded by dissolution/precipitation. A coupling between solid carbonate production, mostly as magnesite, and olivine serpentinisation was demonstrated. Isotopic compositions of carbon in gas phases and carbonates were measured. They were found to be consistent with mass balance calculations, additionally suggesting the existence of an accessory carbon-bearing phase carrying negative 13C signatures. This phase was identified by transmission electron microscopy, as an ill ordered graphite phase, which would account for about 15% of carbon mineral under these conditions. © 2009 Elsevier B.V. All rights reserved.
1. Introduction With the increasing concentration of carbon dioxide in the atmosphere and the likely associated global climate change, there has been an increasing interest in carbon dioxide sequestration (e.g. Holloway, 1997a; Bachu, 2000; Oelkers and Schott, 2005), which consists of capturing and injecting CO2 into a target geological formation. Several geological formations are envisaged as possible options to store CO2: deep sedimentary formations and saline aquifers (e.g. Bachu and Adams, 2003; Kaszuba et al., 2003), depleted oil or gas reservoirs (e.g. Holloway, 1997b; Wildenborg and Lokhorst, 2005), basalts (e.g. Matter et al., 2007; McGrail et al., 2006) and ultrabasic rocks (Goff and Lackner, 1998). The main advantage of CO2 sequestration in basic and ultrabasic rocks is their high reactivity with CO2-rich fluids thus providing divalent cations such as Ca2+, Mg2+ and Fe2+, which may combine with the injected CO2 to form carbonate minerals leading to a very stable sequestration of CO2. Such a good reactivity is evidenced in basalt weathering which plays a major role on exchanges of CO2 between the solid Earth and the atmosphere. For instance, Dessert et al. (2003) showed that the annual CO2 consumption through basalt weathering represents 30–35% of the total flux of Na, K, Ca and Mg associated to silicate weathering (Gaillardet et al., 1999), although those rocks represent only 8% of the total rocks exposed at the Earth surface. Previous experimental studies investigated basalts dissolution levels and calcite precipitation both in laboratory experiments (e.g. Oelkers ⁎ Corresponding author. Tel.: +33 1 44 27 60 90; fax: +33 1 44 27 28 30. E-mail address: [email protected] (I. Martinez). 0009-2541/$ – see front matter © 2009 Elsevier B.V. All rights reserved. doi:10.1016/j.chemgeo.2009.01.026
and Gislason, 2001; McGrail et al., 2006) and injection sites (Matter et al., 2007) showing that water–rock interactions in these settings are in favour of permanent storage of CO2. Forsterite dissolution and magnesite precipitation under high partial pressure of CO2 (pCO2 = 1bar and 100 bar) at 30 and 95 °C were recently studied in batch reactors by Giammar et al. (2005) showing that nucleation of magnesite, requiring large critical saturation index, was a limiting factor for the global reaction. Experiments on olivine exposed to high CO2 pressure (185 bar) and high temperature (155 °C) showed that, under these conditions, a 78% carbonation level of the silicate was achieved after 30 min (O'Connor et al., 2002). Such experiments intended to demonstrate the feasibility of industrial processes of ex situ carbonation of ultrabasic minerals. In order to adequately simulate CO2 sequestration in basic and ultrabasic rocks, knowledge of the mechanisms, kinetics and thermodynamics of carbonate forming reactions with individual major mineral of these rocks is needed. In the present work, we chose to study the interactions between a supercritical H2O–CO2 fluid and three Mg-rich minerals: orthopyroxene, chrysotile and olivine, using both mineralogical and geochemical tools. For providing data on the high pressure–high temperature side of the phenomena, the total pressure (1 kbar) and temperature (400 °C and 500 °C) conditions were more elevated than in most envisioned CO2 storage projects or previous experiments. This has the interest to simulate high temperature events which might affect a CO2-storage site in basic and ultrabasic rocks on long time scales, as well as to provide interesting kinetic and mechanistic data associated to large carbonation levels.
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2. Materials and methods
Table 2 Experimental conditions.
2.1. Starting materials
Sample
Most of the experiments were carried out on (Mg,Fe)2SiO4 olivines which were extracted and separated from a peridotite from San Carlos, Arizona. In order to investigate the effect of the nature of the Mg-bearing mineral, (Mg,Fe)SiO3 orthopyroxene from the same San Carlos peridotite and Mg3Si2O5(OH)4 chrysotile collected in centimetric veins from Phillips Mine, Salt River Canyon (Arizona), were also used. Chemical composition of these starting materials was determined using electron microprobe and is reported in Table 1. The starting materials were also checked by powder X-ray diffraction and found to be pure within the precision of the X-ray diffraction determination. The minerals were crushed into powders with grain sizes ranging from 100 to 200 μm. Oxalic acid (H2C2O4) was the source of CO2. Three different batches of oxalic acid were used; two of them were anhydrous with δ13C = −8.20‰ and δ13C = −33.09‰, respectively; one was di-hydrated with an isotopic composition of δ13C = −20.08‰. When the temperature in the experiment exceeded approximatively 150 °C, depending upon pressure, oxalic acid decomposed through the following reactions (Holloway et al., 1968; Javoy et al., 1978): H2 C2 O4 →CO2 þ CO þ H2 O
ð1Þ
CO þ H2 O↔CO2 þ H2
ð2Þ
The complete decomposition can be written: H2 C2 O4 →2CO2 þ H2
ð3Þ
In the case of di-hydrated oxalic acid, the global decomposition equation is: H2 C2 O4 d2H2 O→2CO2 þ H2 þ 2H2 O
ð4Þ
At relatively low temperatures, complex decomposition schemes are occurring with formation of formic acid (Surdam et al., 1984), but they probably do not affect the high-temperature decomposition process schematized by reactions (3) and (4). 2.2. Experimental procedure Powdered samples (typically 100 mg) were loaded in a 4 cm-long platinum capsule (wall thickness of 0.2 mm) together with oxalic acid (typically 10 mg). In some runs, water or NaCl-rich water was added (2 to 5 mg, Table 2). Large salinities chosen in Seq 33 and Seq 37 (Table 2, 100 g l− 1 NaCl) indeed correspond to average pore aqueous
Table 1 Chemical compositions (wt.%) of silicates used in the high-pressure experiments. Olivine
Chrysotile
Orthopyroxene
n
17
8
17
SiO2 TiO2 Al2O3 Cr2O3 FeO MnO MgO CaO Na2O K2O NiO Estimated H2O Total
40.92 ± 0.54 0.02 ± 0.02 0.02 ± 0.02 0.03 ± 0.03 8.97 ± 0.24 0.13 ± 0.04 49.64 ± 0.39 0.08 ± 0.03 0.02 ± 0.02 0.01 ± 0.01 0.37 ± 0.08 0 100.21 ± 0.70
43.82 ± 0.76 0 0.10 ± 0.02 0.01 ± 0.01 1.05 ± 0.04 0 38.70 ± 0.56 0.06 ± 0.02 0.01 ± 0.01 0.00 ± 0.00 0.05 ± 0.06 16.16 83.84 ± 1.03
54.98 ± 0.70 0.05 ± 0.04 3.21 ± 0.06 0.58 ± 0.05 5.55 ± 0.18 0.11 ± 0.07 33.53 ± 0.80 1.06 ± 0.8 0.04 ± 0.03 0.01 ± 0.01 0.12 ± 0.07 0 99.27 ± 0.97
n: number of analysis.
Seq 9 Seq 10 Seq 11 Seq 13 Seq 14 Seq 15 Seq 16 Seq 31 Seq 32 Seq 33 Seq 34 Seq 35 Seq 36 Seq 37 Seq 38
T (°C) 400 400 400 500 500 500 500 500 500 500 500 400 400 400 400
Mineral Orthopyroxene Orthopyroxene Chrysotile Olivine Olivine Orthopyroxene Chrysotile Olivine Olivine Olivine Olivine Olivine Olivine Olivine Olivine
Silicate
Oxalic acid
Water
mg
mg
mg
200.19 200.50 195.77 100.30 100.10 100.43 100.38 100.30 101.22 100.02 49.80 100.80 99.15 99.90 48.57
19.80 (1) 19.99 (1) 20.27 (1) 10.69 (1) 10.05 (1) 10.79 (1) 10.91 (1) 10.34 (2) 9.84 (3) 10.25 (2) 5.27 (2) 10.23 (2) 10.35 (3) 10.75 (2) 5.02 (2)
0.00 2.14 0.00 0.00 1.58 0.00 0.00 0.00 0.00 4.59 3.44 0.00 0.00 3.70 3.04
All experiments were conducted at 1 kbar during 4 h. Isotopic composition of initial oxalic acid was: (1) δ13C = − 8.2‰; (2) δ13C = −33.09‰ and (3) di-hydrated oxalic acid with δ13C = −20.08‰. In experiments Seq 33 and Seq 37, a NaCl-rich water (with a concentration of 100 g l− 1 of NaCl) was added to the experimental charge.
phases in some sedimentary basins (e.g. Hanor, 1994). The proportions of silicate, oxalic acid and water loaded in the capsule varied in order to investigate different water/solid ratios as well as different water/CO2 ratios. The capsule was welded on both ends and was observed carefully under a binocular to check for leakages. It should be emphasized that the platinum used as capsule material allows some H2 diffusion through the capsule walls and so that, reaction (2) is continuously displaced to the right. Four capsules per run were then loaded in an internally heated Ar pressure vessel, described in more details in Jendrzejewski et al. (1997). In a first step, the pressure was increased at room temperature up to 60% of the desired pressure; then temperature was increased up to 400 °C or 500 °C thus raising the pressure to its peak value. The duration of the experiments was 4 h. This duration was chosen as an empirical compromise between (1) allowing enough time for having significant advancement of the reaction necessary to investigate the carbonation by gas recovery and carbon isotopes and (2) not using too long durations which would correspond to a complete loss of hydrogen through the capsule walls, leading to more oxidized conditions than desired. Experiments were quenched by turning off the internal heater, thus achieving an approximately isobaric cooling (about 50 bar of pressure loss). Argon pressure was then released down to atmospheric pressure in less than 5 min. All capsules were weighed before and after experiments; the precision in weighing did not allow, however, to check for H2 losses. The analysis of the residual gas phase, presented later, showed that in most experiments only small amounts of CO and H2 were found in the capsule, meaning that reaction (2) had been completed to a finite advancement allowing to compute an effective oxygen fugacity and that large H2 loss occurred. The temperature conditions and the initial amount of the reactants are reported in Table 2. 2.3. Analyses After the experiments, the gaseous phases were first recovered by perforation of the capsule under vacuum; the liberated gases entered into an extraction line, used for carbon extraction, which allowed purification and quantification of the main gaseous components (CO2, CO, H2O and H2, see Pineau et al., 1976 for more details). Incondensable gases such as CO and H2 were first separated cryogenically from the condensable gases (CO2 and H2O) trapped into a liquid nitrogen reservoir. Step-heating of this trap allowed liberation of CO2 at approximately −145 °C and of water between −100 and −50 °C. Subsequent quantification of pure CO2 and pure water was then performed using a Toepler pump. The oxidation of the remaining CO, by passing over CuO furnace at 450 °C, allowed its transformation
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Table 3 Composition of the gases (in micromoles) present in the capsule after experiments and volatiles extracted from the solid after pyrolysis, normalized to 1 mg of solid. Sample Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq Seq
9 10 11 13 14 15 16 31 32 33 34 35 36 37 38
Gas phase
Solid phase
nCO2(g)
nH2O(g)
nCO(g)
nH2(g)
nCO2(pyrolysis)
nH2O(pyrolysis)
381 (4) 387 (4) 384 (4) 167 (4) 145 (4) 224 (4) 194 (4) 193 (4) 128 (2) 138 (2) 87 (2) 171 (4) 111 (2) 101 (2) n.a.
4.8 (2) 111 (2) 68 (2) 49 (2) 75 (2) 81 (2) 133 (2) 8.6 (2) 86 (2) 196 (4) 198 (4) 2.8 (2) 121 (2) 102 (2) n.a.
45.3 (2) n.a. 13.5 (2) 0.5 (2) 0.3 (2) 1.0 (2) n.a. 8.2 (2) 2.7 (2) n.a. 0.6 (2) 3.2 (2) 3.0 (2) 2.1 (2) n.a.
39.3 (2) n.a. 36.7 (2) 23.2 (2) 0 16.8 (2) n.a. 13.2 (2) 4.7 (2) n.a. 2.8 (2) 8.4 (2) 2.0 (2) 4.6 (2) n.a.
0.17 (3) 0.23 (4) 0.39 (6) 0.55 (8) 0.75 (11) 0.07 (1) 0.34 (5) 0.27 (4) 0.42 (6) 0.83 (12) 0.17 (3) 0.47 (7) 0.68 (10) 1.17 (18) 0.38 (6)
0.11 (2) 0.13 (2) 6.75 (1.01) 0.11 (2) 0.11 (2) 0.04 (1) 6.56 (98) 0.08 (1) 0.01 (1) 0.18 (3) 0.29 (4) 0.22 (3) 0.14 (2) 0.33 (5) 0.18 (3)
r1
r2
3(1) 5(1) 12(1) 29(1) 38(1) 5(1) 11(1) 13(1) 16(1) 39(1) 25(1) 23(1) 31(1) 57(1)
7.7 (1.2) 10.7 (1.6) 17.6 (2.6) 24.1 (3.6) 35.6 (5.3) 2.80 (40) 14.3 (2.1) 12.1(1.8) 28.0(4.2) 39.1(5.9) 6.7(1.0) 21.4(3.2) 43.7(6.6) 54.4(8.2)
r1 and r2 (in %) are the estimated mole fraction of CO2 sequestered under the form of carbonate calculated from the analysis of the gas phase (r1) and from the CO2 obtained by pyrolysis of the solid (r2). n.a.: non analyzed.
into CO2, which was then trapped into liquid nitrogen and manometrically quantified using a Toepler system, yielding a precision of 5% on the volume. For carbon, the measured blanks for this procedure were 0.2 μmol, thus specifying the detection level. After such procedure, the residual gas is H2, which was also quantified using the Toepler pump. CO and CO2 were collected in glass tubes for further carbon isotopic measurements. Once the gas phases were analysed, the solid phases were recovered. Approximately 25 mg of each of these solids were used for estimating the carbon contents with the step-heating technique. For this purpose, the sample powders were loaded in glass tubes and heated slowly up to 900 °C in the extraction line. During their pyrolysis, solid carbonates were decomposed both into CO2 and CO, which was oxidized using a CuO furnace heated at 450 °C. Gaseous CO2 was continuously collected in a liquid nitrogen trap. After 30 min, the temperature was increased up to 1000 °C in order to burn any residual carbonates. The collected CO2, produced by decarbonation of the samples, was then quantified and analysed for its isotopic composition using a Finnigan Mat Delta E double collecting mass spectrometer. Water produced during this pyrolysis was also collected and quantified in the same way as CO2. Finally, few mg of the solid powder were mounted in epoxy resin and polished for mineralogical studies using Scanning Electron Microscopy (SEM), performed on a JEOL JSM 840A, equipped with an Energy Dispersive System PGT Sahara detector for semi quantitative analyses, working at 15–20 kV and 1–2 nA. Additional analyses were carried out by Analytical Transmission Electron Microscopy (ATEM) on finely ground powders resuspended in ethanol and let to evaporate on carbon coated copper grids. The TEM used was a JEOL 2100 FEG, operated at 200 kV, equipped with an energy dispersive X-ray analysis Si(Li) diode system from JEOL. 3. Results 3.1. Analyses of gas phases Compositions of the gas phases recovered from each experiment are reported in Table 3. They consist mainly of CO2, H2O, CO and H2 in various proportions depending upon silicate starting material, initial water content and salinity. In most experiments, however, only minor amounts of CO and H2 were analysed. The proportions of each gas were different from those predicted from simple oxalic acid decomposition (for example, considering 10 mg of oxalic acid decomposing through reaction (3), 222 μmol of CO2 and 111 μmol of H2 should be formed) because these gases were involved in several reactions that will be discussed later in the paper. Taking into account the small amount of CO measured, reaction
(2) was displaced to the right, thus leading to the formation of CO2 and H2. The small measured quantity of H2 relative to CO2, however, implies a large H2 diffusion through the capsule walls. Nevertheless, the presence of some CO and H2 in the end products ensures that oxygen fugacity in these experiments were low (see Discussion). Carbon isotopic compositions of CO and CO2 were measured in some experiments and are reported in Table 4 using the usual δ notation with the Pee Dee Belemnite standard. The δ13C values are variable and depend upon the isotopic composition of the source of CO2. The fractionation factor, Δ, which is not depending upon the isotopic composition of the source, is calculated and also given in the table. The calculated Δ values can be compared to theoretical values in order to discuss isotope equilibrium conditions. 3.2. Analyses of solid phases Quantities of CO2 and H2O produced by pyrolysis of aliquots of the solid samples are reported in Table 3. The values are given as normalized to a sampled solid mass of 1 mg. Scattering of analyses carried out on two different aliquots of a same sample suggests that large uncertainties of approximately 30% affect these measurements, due to heterogeneities in the solid phases. As will be noticed below, carbon in the solid phases is mainly present as solid carbonates whereas a small but significant amount is present as poorly ordered graphite. This complexity might contribute to sample heterogeneity. Carbon isotopic compositions of CO2 derived from the carbonates present in these solid samples were measured in selected experiments and are given in Table 4. As for the gases, the δ13C obtained on the carbonates reflects the isotopic composition of their source (i.e. oxalic acid). To avoid this, we calculated the fractionation factor Δ, between CO2 and carbonate, which will be used later in the paper to discuss the isotope equilibrium conditions. Table 4 Carbon isotopic composition of the carbon rich phases, expressed using the usual δ notation with the PDB standard (error bars: 0.1‰) and calculated isotopic fractionations (Δ). Sample T (°C) Mineral Seq 9 Seq 10 Seq 11 Seq 16 Seq 31 Seq 35 Seq 37
400 400 400 500 500 400 400
Orthopyroxene Orthopyroxene Chrysotile Chrysotile Olivine Olivine Olivine
δ13CCO2
δ13CCO
δ13Ccarbonate ΔCO2–carb
‰
‰
‰
n.a. − 4.95 − 4.53 − 6.82 − 32.70 − 30.97 − 29.33
n.a. n.a. − 35.60 n.a. −47.30 −59.70 −60.00
− 9.39 −9.53 −9.38 − 10.51 − 28.90 − 31.18 − 32.46
n.a 4.58 4.85 3.69 − 3.8 0.21 3.13
ΔCO2–CO n.a. n.a. 31.07 n.m 14.6 28.73 30.67
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The carbonation reactions were visualized using SEM. An example of a carbonated olivine powder is shown in Fig. 1. Olivine crystals (50 μm in size; Table 3, experiment Seq 31) were surrounded by reaction rims in which small white crystals are visible. Chemical analyses using the EDS of the SEM showed that these small crystals are Mg–Fe-rich and can be interpreted as being the newly formed iron bearing magnesite grains, surrounded by a silica-rich matrix in dark grey on the figure, as confirmed by TEM investigations (see below). It is interesting to observe in Fig. 1 that some olivine grains are showing two or three distinct reactions rims. The enhancement of carbonation in the presence of water, using di-hydrated oxalic acid (experiment Seq 32, Table 3), is clear from Fig. 2. This figure shows olivine grains surrounded by small magnesite crystals of approximately 2–5 μm in size, dispersed all over the sample. When sodium chloride was added to the experiment, carbonation levels were significantly increased (Table 3, experiments Seq 33 and Seq 37), which is also clear from SEM images. Fig. 3a shows microstructures typical of such samples (Seq 33) in which large globules of magnesite of approximately 10 μm formed during the reaction. EDX analyses of magnesite, small precipitates of NaCl and olivine are given in Fig. 3b. Comparison of carbonate abundance (Table 3) obtained after pyrolysis on samples reacted at 400 °C with 500 °C samples (initial content of water and salinity being equal) shows that carbonates are more abundant at the lowest temperature. For instance, Seq 31 (500 °C) and Seq 35 (400 °C) gave respectively 13(1) and 23(1) % of neo-formed carbonates (where numbers in parenthesis stand for the error bar); Seq 32 (500 °C) and Seq 36 (400 °C) gave 16(1) and 31(1) % of neo-formed carbonates and finally Seq 33 (500 °C) and Seq 37 (400 °C) gave 39(1) and 57(1) % of neo-formed carbonates. The largest amounts of CO2 transformed into solid carbonates were obtained from an olivine reacted at 400 °C with 100 g l− 1 of NaCl-containing water (Seq 37). ATEM studies on reacted powders showed the presence of euhedral magnesite crystals (Fig. 4), identified both by their EDX spectra (not shown here) and electronic diffraction pattern (inset in Fig. 4). Such crystals are present in all run products but not in the initial powders. Microprobe analyses on the largest magnesite crystals showed that they are iron-bearing with average Fe/Fe + Mg value of 10 atomic %, versus a Fe/Fe + Mg of initial olivine of 11 atomic %. Such chemical compositions are in agreement with previous experimental work on San Carlos olivine (O'Connor et al., 2002). In all samples, magnesites were systematically associated with a “proto-serpentine” phase, showing a fibrous microstructure (Fig. 5). This proto-serpentine is the silica rich cement surrounding the small Mg-rich carbonates grains observed by SEM. EDX spectra showed that the proto-serpentine also contained iron, with average
Fig. 1. Backscattered electron image of sample Seq 31 showing large olivine crystals in light grey surrounded by a magnesite (white spots) embedded in a silica rich layer (in dark grey). Several reactions rims are observed around olivine crystals showing that the reaction is surface controlled.
Fig. 2. Backscattered electron image from sample Seq 32. Carbonation reaction is enhanced compared to Seq 31 because of the use of di-hydrated oxalic acid that increases the partial pressure of water in the experiment. Compared to sample Seq 31, magnesite crystals are larger (several microns in size) and dispersed out all over the samples.
Fe/Fe + Mg ratio of 7 atomic %. Determining the exact nature, speciation and oxidation state of iron in the proto-serpentine phase, which might be present as either Fe2+, Fe3+ or both, would require further high resolution analytical TEM studies. Finally, small quantities of poorly crystalline graphite were found in the reacted samples (Fig. 6). When orthopyroxene (Seq 9, 10 and 15) and serpentine (Seq 11 and 16) were the starting materials, smaller amounts of carbonates were formed; secondary proto-serpentine phases were observed as well. It is interesting to note that when NaCl was added to the experiments, the proto-serpentine phases contained some chlorine. 4. Discussion 4.1. Carbonation levels Carbonation levels were calculated using two methods. The first calculation is based on the analyses of the gas phase present in the capsule after the experiment. The difference between the initial number of moles of carbon introduced in the capsule as oxalic acid and the measured CO + CO2 is attributed to the formation of solid carbonates, thus giving the number of moles of CO2 sequestered. This number, reported as a percentage of the total equivalent quantity of CO2, is called r1 in Table 3. The second calculation used the equivalent quantity of CO2 measured by degassing the solid after the experiment, which is attributed to newly formed carbonates. This number, called r2, is reported in Table 3. The uncertainties on r2 are estimated to be of about 30% by comparing the results of pyrolysis experiments carried out on several aliquots from a same experimental run. In this respect, determination of CO2 mineral sequestration levels by analysis of the gases (r1) appears to be more precise in spite of possible changes in composition of the gas phase during quench. Moreover, since all major C-bearing carbon phases were analysed (i.e. CO and CO2), potential problems related to quench are indeed minor. Taking into account uncertainties, it is concluded that the carbonation levels measured by the two methods are mutually consistent except for samples Seq 32, 34 and Seq 36 in which the analysis of the solid by pyrolysis was presumably affected by strong heterogeneities. In all cases, it appears that the formation of solid carbonates is significant after 4 h at these pressure–temperature conditions and that mineral sequestration on the order of tens of percents are obtained. The carbonation levels increase from orthopyroxene, to chrysotile and then to olivine. From
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Fig. 3. (a) Backscattered electron image obtained from sample Seq 33, in which 100 g l− 1 NaCl aqueous phase was added in the capsule. Large spherules of magnesites (approximatively 10 μm in size) were formed all over the sample, in agreement with a precipitation process from a fluid phase. (b) Corresponding energy dispersive X-ray analyses of magnesite, sodium chloride (due to the small size of these precipitates, spectrum is contaminated by olivine) and olivine.
the different experimental conditions (temperature, salinity, dry or wet oxalic acid, water added) investigated with olivine, the largest CO2 mineral sequestration levels were obtained in experiments conducted in saline waters (100 g l− 1 NaCl), in agreement with previous experiments (O'Connor et al., 2002). Fig. 7 summarizes the measured carbonation levels: it appears clearly that carbonation is more efficient at 400 °C than at 500 °C, consistently with the microstructural observations shown above. The simplest way to rationalize this observation is to notice that equilibrium CO2 fugacity above magnesite is larger at 500 °C than at 400 °C and that the amount of CO2 stored as a solid is thus smaller at higher temperatures (see below). The low reactivity of CO2 at high temperature was also evidenced earlier by Bischoff and Rosenbauer (1996) in their high pressure study of rhyolite alteration in CO2 charged water at 200 °C and 350 °C. The greater alteration of the albitic plagioclase was partly explained in their study by the greater acidity of carbonic acid at 200 °C. Possible formation of solid carbonates upon pressure quench cannot be definitely excluded. However, the size of magnesites formed in our experiments (2–5 μm in most NaCl-free experiments and up to 10 μm in sample Seq 33) are large compared to those usually formed upon quench which are usually small euhedral crystals with typical sizes below 1 μm. Moreover, the CO2 fugacities calculated in equilibrium with magnesite at 400 °C and 500 °C (see below) are qualita-
tively in agreement with the measured CO2 pressures; this suggests a high pressure–high temperature equilibrium of CO2 with magnesite. 4.2. Mechanisms Carbonates were observed on silicate surface together with a Sirich layer (Fig. 1). TEM confirmed these observations, showing magnesite systematically associated to a fibrous proto-serpentine phase (Fig. 5). This leads to the conclusion that carbonation is coupled to a serpentinisation reaction. The newly formed magnesite crystals are euhedral and well faceted, as shown in Fig. 4. No crystallographical relationship between these magnesite crystals and the adjacent silicates was observed. Moreover, the microstructures observed by SEM, such as corrosion gulfs in olivine crystals (Fig. 8), strongly suggest a mechanism controlled at the fluid/silicate interface, most probably a dissolution/precipitation reaction such as described below. Dissolution of olivine yields a “proto-serpentine” phase and a magnesium-rich fluid. When carbonate saturation is achieved in the fluid, then magnesite precipitates. The dissolution reaction can be summarized as: 2þ
2Mg2 SiO4 þ 3H2 O þ 2CO2 ↔Mg
−
þ Mg3 Si2 O5 ðOHÞ4 þ 2HCO3
ð5Þ
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Fig. 6. ATEM image of poorly crystalline graphite in sample Seq 33 (associated diffraction pattern in the insert) forming small rounded shape spherules (see arrows), embedded within proto-serpentine fibers.
Fig. 4. Analytical Transmission Electron Microscopy image showing an example of a euhedral crystal of magnesite formed during the experiment Seq 31. The selected area diffraction pattern, indexed using the R-3C space group, is shown in the insert.
using a simplified form for the olivine formula, not taking into account Fe and minor elements, and an ideal stoichiometric chrysotile formula for proto-serpentine. Then the precipitation reaction occurs according to: 2þ
Mg
−
þ 2HCO3 ↔MgCO3 þ H2 O þ CO2
ð6Þ
which by summing Eqs. (5) and (6) leads to: 2Mg2 SiO4 þ 2H2 O þ CO2 ↔MgCO3 þ Mg3 Si2 O5 ðOHÞ4
ð7Þ
In turn, this proto-serpentine phase can be carbonated according to: Mg3 Si2 O5 ðOHÞ4 þ 3CO2 ↔3MgCO3 þ 2H2 O þ 2SiO2
ð8Þ
SiO2 is given here as an ideal end-member product of the reaction. A free silica phase is rarely observed in these experiments. It is likely
Fig. 5. ATEM image of a euhedral magnesite crystal surrounded by proto-serpentine fibers, in sample Seq 37, suggesting a relationship between carbonation and serpentinisation reactions.
Fig. 7. Carbonation levels, expressed as mole percent of sequestered CO2, as a function of temperature (light grey = 500 °C, dark grey = 400 °C) and fluid composition. Note that the highest carbonation level (57 mol%) was found at 400 °C with NaCl aqueous solution added to the capsule.
Fig. 8. Backscattered electron image of sample Seq 38 showing a large olivine crystal with corrosion gulfs suggesting dissolution of the silicate. Magnesites are smaller crystals dispersed in the surrounding matrix.
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that the excess silicon is rather present in the proto-serpentine which is known for being able to accommodate variable (Mg + Fe)/Si ratios. The sum of Eqs. (7) and (8) leads to the complete carbonation of olivine according to:
Table 6 Molar fractions of CO2 (X), CO (Y), carbonate (Z) and calculated isotopic composition of oxalic acid using Eq. (11) in text. Sample T (°C) Mineral
X
Mg2 SiO4 þ 2CO2 ↔2MgCO3 þ SiO2
Seq 10 Seq 11 Seq 16 Seq 31 Seq 35 Seq 37
0.87 0.85 0.80 0.84 0.75 0.42
ð9Þ
Depending on the relative advancements of the processes (5) to (8), the leading mechanism will be a combination of Eqs. (7) and (8), the proportions of which can be estimated by considering the CO2 and H2O mole numbers obtained after pyrolysis of the solids and given in Table 3. If a and b are the advancement coefficients of Eqs. (7) and (8) respectively, the global reaction can be written as:
400 400 500 500 400 400
Orthopyroxene Chrysotile Chrysotile Olivine Olivine Olivine
(2) (2) (2) (2) (2) (2)
Measured Calculated δ13Coxalic acid δ13Coxalic acid
Y
Z
0.11 (2) 0.03 (2) 0.13 (2) 0.03 (2) 0.02 (2) 0.01 (2)
0.02 (4) −4.49 (8) 0.12 (4) −6.04 (8) 0.07 (4) −6.18 (8) 0.13 (4) −32.74 (8) 0.23 (4) − 31.45 (8) 0.57 (4) − 31.36 (8)
− 8.20 (5) − 8.20 (5) − 8.20 (5) − 33.09 (5) − 33.09 (5) − 33.09 (5)
The numbers in parenthesis stand for the error bars. The measured isotopic composition of oxalic acid is given for comparison.
equilibrium thermodynamics could explain that water has a positive effect on carbonation of olivine, as suggested by reaction (10).
2aMg2 SiO4 þ ð2a−2bÞH2 O þ ða þ 3bÞCO2 ↔ða þ 3bÞMgCO3 þ ða−bÞMg3 Si2 O5 ðOHÞ4 þ 2bSiO2
ð10Þ 4.3. Isotopic measurements
The a and b values determined from the CO2 and H2O mole numbers given in Table 3 are shown in Table 5. From these data, a sequestration level (called r3 in Table 5) can be calculated using the mole number of carbonate, a + 3b, divided by the initial mole number of CO2 introduced. For most samples, sequestration levels r3 are similar to what is calculated with the mole number of CO2 obtained from pyrolysis (r2). This means that both the quantities of CO2 and H2O left in the sample are consistent with the proposed reaction (10) for interpreting the carbonation reaction, and that part of the serpentine formed during this process gave in turn some carbonate by reaction with CO2. As mentioned above, thermodynamic equilibrium is not achieved in the present experiments, especially because of possible continuous H2 loss from the capsules. Although equilibrium (9) is not achieved, it is still possible to obtain qualitative arguments from equilibrium thermodynamics. In particular, using molar volumes, enthalpies and entropies data for quartz, magnesite and forsterite (Holland and Powell, 1990) and assuming SiO2 (as quartz) activity of 1; equilibrium CO2 fugacities for Eq. (9) are 1.9 kbar at 500 °C and 0.26 kbar at 400 °C, respectively. Due to the different approximations and departure from equilibrium, we do not intend to mean that these values are realistic for predicting CO2 fugacities in these experiments, but we notice that these values are actually in the same order of magnitude as CO2 fugacities deduced from the measured CO2 mole numbers in the capsules and that the main temperature effects, i.e. strong increase in equilibrium fCO2 with temperature, will be likely retained in any carbonation process of this type. We thus suggest that carbonation levels are proportional to the difference between initial (after oxalic acid decomposition) and equilibrium CO2 fugacities in the capsule. A CO2 fugacity in equilibrium with magnesite lower at 400 °C than at 500 °C could thus explain faster carbonation rates at the lowest of these two temperatures. Similar correlations between kinetics and
Table 5 Results of a and b advancement coefficients for Eqs. (7) and (8), respectively (see text). Sample
a
b
r3
r2
Seq Seq Seq Seq Seq Seq Seq Seq Seq
0.179 0.229 0.098 0.109 0.275 0.151 0.2 0.223 0.416
0.124 0.174 0.058 0.104 0.185 0.006 0.09 0.153 0.251
24 32.6 11.8 18.5 36.1 7.4 20.6 29.3 50.8
24.1 (3.6) 35.6 (5.3) 12.1 (1.8) 28.0 (4.2) 39.1 (5.9) 6.7 (1.0) 21.4 (3.2) 43.7 (6.6) 54.4 (8.2)
13 14 31 32 33 34 35 36 37
A sequestration level calculated from the mole number of formed carbonates (a + 3b) is calculated and called r3. The numbers in parenthesis stand for the error bar. For comparison, r2, calculated directly from the CO2 obtained by pyrolysis of the solid, is also reported.
Isotopic compositions of CO2, CO and carbonates reported in Table 4 are variable, first of all because the isotopic composition of the carbon source is variable. If we then calculate the fractionation factor (Δ), written for example for ΔCO2–carbonates as the difference δ13CCO2 − δ13Ccarbonates, we can get rid of the variability in isotopic composition of the source. However, the measured values of ΔCO2–carbonates are also variable and do not show connection with isotopic equilibrium values (e.g. 4‰ at 400 °C and 5‰ at 500 °C) as measured by Appora-Gnekindy (1998). On the other hand, values of ΔCO2–CO, although showing also some variability, are close to equilibrium values (25‰ at 400 °C and 17‰ at 500 °C) as calculated in Richet et al. (1977). Indeed, conditions close to isotopic equilibrium are probably achieved in the gas phase, meaning that the quench did not affect the gas mixture. In the case of gas/solid isotope exchange, the heterogeneities might come from the fact that only aliquots of the solid phases were analysed. Moreover, as will be discussed below, the presence of a non analysed isotopically light phase, such as graphite, in various proportions in the solid, may affect strongly the isotopic composition of the companion carbonates. The major implication of these isotopic analyses is that knowing the isotopic composition of the source of carbon, oxalic acid, mass balance calculations can be done as: 13
13
13
13
δ Coxalic acid ¼ Xδ CCO2 þ Yδ CCO þ Zδ Ccarbonate
ð11Þ
where X, Y, and Z are molar fractions of the different phases and δ their isotopic composition. Results of this calculation are given in Table 6. The mass balance is consistent with the other determinations, but its imprecision, due to the aforementioned solid-state heterogeneities, makes it not very useful at this stage. An interesting observation, however, is that these calculations give total carbon isotopic compositions (calculated δ13Coxalic acid in Table 6) that are always less negative than the total carbon isotopic composition actually measured in the starting material (measured δ13Coxalic acid in Table 6). The more likely interpretation of this systematic difference is that a phase with negative δ13C signature is missing from the mass
Table 7 Molar fractions of CO2 (X), CO (Y), carbonate (Z) and reduced carbon (R) obtained from Eq. (12) by isotopic mass balance. Sample
T (°C)
Mineral
X
Y
Z
R
Seq 10 Seq 11 Seq 16 Seq 31 Seq 35 Seq 37
400 400 500 500 400 400
Orthopyroxene Chrysotile Chrysotile Olivine Olivine Olivine
0.68 (4) 0.73 (4) 0.69 (4) 0.79 (4) 0.64 (4) 0.37 (4)
0.08 (4) 0.03 (4) 0.11 (4) 0.03 (4) 0.01 (4) 0.01 (4)
0.02 (4) 0.10 (4) 0.06 (4) 0.12 (4) 0.20 (4) 0.49 (4)
0.22 (4) 0.14 (4) 0.14 (4) 0.06 (4) 0.15 (4) 0.13 (4)
The numbers in parenthesis are the error bars.
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balance. TEM observations show that a “graphite-type” phase was formed in association with carbonates (Fig. 6). Such reduced carbon could be this missing phase since graphite, which is not pyrolized together with carbonates, has usually values of δ13C more negative than carbonates (e.g. Des Marais, 2001). Using ΔCO2–graphite values calculated at this temperature (Bottinga, 1969), a putative molar fraction (called R in the following formula) of this reduced carbon phase can be estimated from: 13
13
13
13
13
δ Coxalic acid ¼ Xδ CCO2 þ Yδ CCO þ Zδ Ccarbonate þ Rδ C“graphite”
ð12Þ
and is provided in Table 7, showing that such an ill-ordered graphite phase could represent up to 22 mol% of the initial carbon injected in these experiments as oxalic acid. Also, the “graphite-type” phase present in the samples could contribute to the large isotopic heterogeneity of carbonates. Various proportions of graphite in the samples, having a very negative isotopic composition, close to the carbonates, will make the isotopic compositions of the carbonates very variable. Only the analysis of both content and isotopic composition of this “graphite-type” phase, which is not easy in such small amounts, will lead to a correct view of the isotopic equilibrium or non-equilibrium in these experiments. This result corroborates the predicted oxygen fugacity in the experiments, estimated using the after-equilibration gas composition from the reaction: CO þ 1=2O2 →CO2
ð13Þ
Using enthalpy and entropy data from Holland and Powell (1990), as well as a polynomial function for CO2 fugacity also from Holland and Powell (1990), assuming ideal behaviour of CO and ideal mixing of gases, average values of oxygen fugacities, fO2 (log units using a 1 bar pressure reference) were −32 ± 0.7 and − 25 ± 0.9 at 400 °C and 500 °C, respectively. These values are also close to that obtained using the empirical formula given in Floden et al. (1998). These fO2 values should be used only as qualitative indicator of the redox conditions of the experiments, not as equilibrium fO2 values because of H2 loss from the platinum capsules preventing equilibrium to be reached. These oxygen fugacities are close to the CCO buffer (C as graphite–CO, recalculated at 1 kbar) (Arculus et al., 1990). Such low values are likely due to the use of oxalic acid as a CO2 source, providing H2 during its decomposition. They can explain why we did not identify magnetite in the run products, which should normally form during serpentinisation reaction, and they are in agreement with the presence of a graphite phase in the solid products. 5. Conclusions Carbonation of powdered olivine, orthopyroxene and chrysotile has been studied at high pressure and high temperatures (1 kbar and 400–500 °C). As analysed quantitatively by post-experiment gas recovery, the fluid obtained by oxalic acid decomposition, complemented in some experiments with water and NaCl, is composed of CO2–H2O in most of the experiments and contains CO and H2 in smaller proportions. Due to possible continous H2 loss from capsules, thermodynamic equilibrium is not achieved in these experiments whose purpose is to provide new kinetic data on silicate carbonation as well as to study methods for measuring advancement of the carbonation process. In such conditions, olivine is the most reactive mineral with 57 mol% carbonation obtained at 400 °C, after 4 h experiment, in CO2 + additional NaCl-rich water. As previously described in literature, the reaction proceeds in two steps: silicate dissolution then followed by precipitation of magnesite. Presence of water and/or NaCl in the fluid is found to favour the reaction. Detailed mineralogical study reveals that serpentinisation of olivine is coupled to the carbonation reaction. Carbon isotopic measurements allowed to make mass balance calculations, which suggest that, under the
reducing conditions of these experiments, a non negligible fraction (up to 22 mol% of the total sequestered carbon) can be found under the form of poorly crystalline graphite, also identified by transmission electron microscopy. Further experiments under such reducing conditions should concentrate on a more precise quantification of this reduced carbon phase, which can also be a potential source for carbon storage, mostly never taken into account in the models, and which is to be related to the abiogenic generation of hydrocarbons at lower temperatures (Berndt et al., 1996). These studies could also help to understand the redox sate of carbon during serpentinisation of oceanic crust. Acknowledgements We are grateful to Omar Boudouma and to François Guyot for their help during SEM and TEM studies of the samples. We particularly thank Françoise Pineau for several discussions on isotopic data. The manuscript was improved by insightful reviews by Olivier Vidal, Bill O'Connor and Pascale Bénézeth. This study was supported by the Centre de Recherches sur le Stockage Géologique du CO2, Institut de Physique du Globe de Paris/TOTAL/ SCHLUMBERGER/ADEME partnership. This is IPGP contribution 2383. References Appora-Gnekindy, I., 1998. Etude expérimentale du fractionnement isotopique du carbone et de l'oxygène dans les systèmes CO2–Carbonates liquides: Application aux contextes carbonatitiques. PhD. Thesis. Université Denis Diderot- Paris 7, IPGP, Paris, 351pp. Arculus, R.J., Holmes, R.D., Powell, R., Richter, K., 1990. Metal–silicate equilibria and core formation. In: Newsom, H.E., Jones, J.H. (Eds.), Origin of the Earth. Oxford University Press, New York, pp. 251–271. Bachu, S., 2000. Sequestration of CO2 in geological media: criteria and approach for site selection in response to climate change. Energy Convers. Manag. 41 (9), 953–970. Bachu, S., Adams, J.J., 2003. Sequestration of CO2 in geological media in response to climate change: capacity of deep saline aquifers to sequester CO2 in solution. Energy Convers. Manag. 44 (20), 3151–3175. Berndt, M.E., Douglas, E.A., Seyfried, W.E., 1996. Reduction of CO2 during serpentinization of olivine at 300 °C and 500 bar. Geology 24 (4), 351–354. Bischoff, J.L., Rosenbauer, R.J., 1996. The alteration of rhyolite in CO2 charged water at 200 and 350 °C: the unreactivity of CO2 at higher temperature. Geochim. Cosmochim. Acta 60 (20), 3859–3867. Bottinga, Y., 1969. Calculated fractionation factors for carbon and hydrogen isotope exchange in the system calcite–carbon dioxide–graphite–methane–hydrogen– water vapor. Geochim. Cosmochim. Acta 33 (1), 49–64. Des Marais, D., 2001. Isotopic evolution of the biogeochemical carbon cycle during the Precambrian. In: Valley, J.W., Cole, D.R. (Eds.), Stable Isotope Geochemistry. Reviews in Mineralogy and Geochemistry, vol. 43. Mineralogical Society of America, pp. 555–578. Dessert, C., Dupre, B., Gaillardet, J., François, L.M., Allegre, C.J., 2003. Basalt weathering laws and the impact of basalt weathering on the global carbon cycle. Chem. Geol. 202 (3–4), 257–273. Floden, A.M., Colson, R.O., Nermoe, M.K.B., Hendrickson, T.R., 1998. Effects of CO on the Activity of Nickel in a Simple Silicate Melt. Lunar Planet. Sci. Conf., XXIX, p. 1520. Gaillardet, J., Dupre, B., Louvat, P., Allegre, C.J., 1999. Global silicate weathering and CO2 consumption levels deduced from the chemistry of large rivers. Chem. Geol. 159 (1–4), 3–30. Giammar, D.E., Bruant, J.J., Robert, G., Peters, C.A., 2005. Forsterite dissolution and magnesite precipitation at conditions relevant for deep saline aquifer storage and sequestration of carbon dioxide. Chem. Geol. 217 (3–4), 257–276. Goff, F., Lackner, K.S., 1998. Carbon dioxide sequestring using ultramafic rocks. Environ. Geosci. 5 (3), 89–101. Hanor, J.C., 1994. Physical and chemical controls on the composition of waters in sedimentary basins. Mar. Pet. Geol. 11 (1), 31–45. Holland, T.J.B., Powell, R., 1990. An enlarged and updated internally consistent thermodynamic dataset with uncertainties and correlations: the system K2O– Na 2O–CaO–MgO–MnO–FeO–Fe2O3–Al2O3–TiO2–SiO2–C–H2–O2. J. Metamorph. Geol. 8, 89–124. Holloway, S., 1997a. An overview of the underground disposal of carbon dioxide. Energy Convers. Manag. 38, 193–198. Holloway, S., 1997b. Safety of the underground disposal of carbon dioxide. Energy Convers. Manag. 38 (Supplement 1), S241–S245. Holloway, J.R., Burnham, W.C., Millhollen, G.L., 1968. Generation of H2O–CO2 mixture for use in hydrothermal experimentation. J. Geophys. Res. 73 (20), 6598–6600. Javoy, M., Pineau, F., Liyama, I., 1978. Experimental determination of the isotopic fractionation between gaseous CO2 and carbon dissolved in tholeitic magma. Contrib. Mineral. Petr 67 (1), 35–39. Jendrzejewski, N., Trull, T.W., Pineau, F., Javoy, M., 1997. Carbon solubility in Mid-Ocean Ridge basaltic melt at low pressures (250–1950 bar). Chem. Geol. 138 (1–2), 81–92.
F. Dufaud et al. / Chemical Geology 265 (2009) 79–87 Kaszuba, J.P., Janecky, D.R., Snow, M.G., 2003. Carbon dioxide reaction processes in a model brine aquifer at 200 °C and 200 bars: implications for geologic sequestration of carbon. Appl. Geochem. 18, 1065–1080. Matter, J.M., Takahashi, T., Goldberg, D., 2007. Experimental evaluation of in situ CO2– water–rock reactions during CO2 injection in basaltic rocks: implications for geological CO2 sequestration. Geochem. Geophys. Geosyst. 8 (2), 1–19. McGrail, B.P., Schaef, H.T., Anita, M., Chien, Y.J., Dooley, J.J., Davidson, C.L., 2006. Potential for carbon dioxide sequestration in flood basalts. J. Geophys. Res. 111 (No B12), B12201. O'Connor, W.K., Dahlin, D.C., Rush, G.E., Dahlin, C.L., Collins, W.K., 2002. Carbon dioxide sequestration by direct mineral carbonatation: process mineralogy of feed and products. Miner. Metall. Process. 19 (2), 95–101. Oelkers, E.H., Gislason, S.R., 2001. The mechanism levels and consequences of basaltic glass dissolution: I. An experimental study of the dissolution levels of basaltic glass as a function of aqueous Al, Si and oxalic acid concentration at 25 °C and pH = 3 and 11. Geochim. Cosmochim. Acta 65 (21), 3671–3681.
87
Oelkers, E.H., Schott, J., 2005. Geochemical aspects of CO2 sequestration. Chem. Geol. 217, 183–186. Pineau, F., Javoy, M., Bottinga, Y., 1976. 13C/12C ratios of rocks and inclusions in popping rocks of the mid-aztlantic ridge and their bearing on the probleme of isotopic composition of deep-seated carbon. Earth Planet. Sci. 29, 413–421. Richet, P., Bottinga, Y., Javoy, M., 1977. A review of hydrogen, carbon, nitrogen, oxygen, sulphur and chlorine stable isotope fractionation among gaseous molecules. Annu. Rev. Earth Planet. Sci. 5, 65–110. Surdam, R.C., Boese, S.W., Crossey, L.J., 1984. The chemistry of secondary porosity. In: McDonald, D.A., Surdam, R.C. (Eds.), Clastic Diagenesis. AAPG Memoir, vol. 37, pp. 127–151. Wildenborg, T., Lokhorst, A., 2005. Introduction on CO2 geological storage. Classification of storage options. Oil Gas Sci. Technol. 60, 513–515.