Caprock interaction with CO2: A laboratory study of reactivity of shale with supercritical CO2 and brine

Applied Geochemistry 26 (2011) 1975–1989

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Caprock interaction with CO2: A laboratory study of reactivity of shale with supercritical CO2 and brine Binyam L. Alemu a,⇑, Per Aagaard a, Ingrid Anne Munz b,1, Elin Skurtveit c a

University of Oslo, P.O. Box 1047, Blindern, NO-0312 Oslo, Norway Institute for Energy Technology, P.O. Box 40, NO-2027 Kjeller, Norway c Norwegian Geotechnical Institute, P.O. Box 3930, Ullevål Stadion, NO-0806 Oslo, Norway b

a r t i c l e

i n f o

Article history: Received 12 November 2010 Accepted 21 June 2011 Available online 28 June 2011 Editorial handling by R. Fuge

a b s t r a c t Crushed rock from two caprock samples, a carbonate-rich shale and a clay-rich shale, were reacted with a mixture of brine and supercritical CO2 (CO2–brine) in a laboratory batch reactor, at different temperature and pressure conditions. The samples were cored from a proposed underground CO2 storage site near the town of Longyearbyen in Svalbard. The reacting fluid was a mixture of 1 M NaCl solution and CO2 (110 bar) and the water/rock ratio was 20:1. Carbon dioxide was injected into the reactors after the solution had been bubbled with N2, in order to mimic O2-depleted natural storage conditions. A control reaction was also run on the clay-rich shale sample, where the crushed rock was reacted with brine (CO2-free brine) at the same experimental conditions. A total of 8 batch reaction experiments were run at temperatures ranging from 80 to 250 °C and total pressures of 110 bar (40 bar for the control experiment). The experiments lasted 1–5 weeks. Fluid analysis showed that the aqueous concentration of major elements (i.e. Ca, Mg, Fe, K, Al) and SiO2 increased in all experiments. Release rates of Fe and SiO2 were more pronounced in solutions reacted with CO2–brine as compared to those reacted with CO2-free brine. For samples reacted with the CO2–brine, lower temperature reactions (80 °C) released much more Fe and SiO2 than higher temperature reactions (150–250 °C). Analysis by SEM and XRD of reacted solids also revealed changes in mineralogical compositions. The carbonate-rich shale was more reactive at 250 °C, as revealed by the dissolution of plagioclase and clay minerals (illite and chlorite), dissolution and re-precipitation of carbonates, and the formation of smectite. Carbon dioxide was also permanently sequestered as calcite in the same sample. The clay-rich shale reacted with CO2–brine did not show major mineralogical alteration. However, a significant amount of analcime was formed in the clay-rich shale reacted with CO2-free brine; while no trace of analcime was observed in either of the samples reacted with CO2–brine. Ó 2011 Elsevier Ltd. All rights reserved.

1. Introduction Limiting the rise of CO2 concentration in the atmosphere is viewed by many as one of the biggest environmental challenges of the 21st century. In recent years, subsurface CO2 storage has been considered the best option for reducing CO2 emissions from large point sources (Bachu, 2000; IPCC, 2005). This process involves capturing the CO2 and injecting it into various types of traps, such as deep saline formations, depleted oil and gas fields and coal seams (Holloway, 1997). Deep saline aquifers provide the largest storage capacity for underground storage of CO2. They are often found within good proximity to point sources which has the economic advantage of reduced costs for transportation of CO2 to storage sites (Martínez et al., 2009). ⇑ Corresponding author. Tel.: +47 45798581; fax: +47 22854215. E-mail address: [email protected] (B.L. Alemu). Present address: Research Council of Norway, P.O. Box 2700, St. Hanshaugen, NO0131 Oslo, Norway. 1

0883-2927/$ - see front matter Ó 2011 Elsevier Ltd. All rights reserved. doi:10.1016/j.apgeochem.2011.06.028

Svalbard, being an island, has a closed energy system that relies on coal to produce energy for its electricity supply. The largest settlement and administrative center of Svalbard, Longyearbyen, aspires to make itself a zero emission city by the year 2025. With this vision, Longyearbyen could be the first community in the world that has no man-made CO2 emissions while relying totally on C-based energy. The aim is to achieve this by capturing CO2 from the coal power-plant and storing it in close proximity in a deep sandstone reservoir. Currently, the coal-combusting power plant emits c.a. 85,000 tons of CO2/a. The target reservoir is at a depth of 900 m below the surface and has a temperature close to 30 °C. The reservoir is overlain by a thick shale sedimentary succession capable of preventing CO2 loss. Reservoirs selected for CO2 storage should have the capacity to store CO2 with no significant leakage for time periods of 1–10 ka (Bowden and Rigg, 2005). Once the CO2 is injected into the aquifer, it will be sealed by structural or stratigraphic trapping (physically trapped by an impermeable layer) and hydrodynamic trapping as a residual fluid in the reservoir. A small fraction will dissolve

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(aqueous trapping) and after reaction and considerable time, part of the CO2 will be incorporated into carbonate minerals i.e. by mineral trapping (Bachu et al., 1994; IPCC, 2005). The existence of natural CO2 reservoirs trapped for geologic time (Lu et al., 2009) would seem to suggest that the potential to mineralize all CO2 may be limited. Natural analogues from deep CO2-containing reservoirs have demonstrated that shale, mudstone and claystone can effectively trap CO2 over geological time spans, with no significant compromise in integrity (Haszeldine et al., 2005; Lu et al., 2009; Rubert et al., 2009). Any leakage from natural storage traps seems to be limited to escape routes i.e., fractures and faults, rather than, being caused by chemical modifications of porosity (May, 2005; Shipton et al., 2004). In geological storage environments the overall chemical interaction is dictated by the mineralogical composition of the rocks and the evolving fluid composition that will either inhibit, or facilitate the rates of the individual reactions (Lasaga et al., 1994). Therefore, resistance to chemical reaction may possibly be a sitespecific phenomenon. There have been experimental studies on the reactivity of caprock and reservoir material with brine and supercritical CO2 (CO2–brine) (e.g., Credoz et al., 2009; Gunter et al., 1997; Kaszuba et al., 2003, 2005; Wigand et al., 2008) and simulation studies (e.g., Audigane et al., 2007; Cantucci et al., 2009; Gaus et al., 2005; Gherardi et al., 2007; Wigand et al., 2008; Xu et al., 2007), where some reactivity has been reported. This study focuses on the reactivity of shale with reservoir fluids during geological storage of CO2. Laboratory batch reaction experiments were used to investigate the reactivity of shale where crushed rock samples were reacted with CO2–brine. The samples were cored from a caprock overlying a deep sandstone reservoir, near the town of Longyearbyen. Two caprock samples (carbonate-rich and clay-rich shale) were reacted with a CO2–brine mixture in a batch reactor. Reaction parameters were varied in order to determine the reactive responses to different sets of conditions. A total of eight experiments were run for periods ranging from 1 to 5 weeks, at temperatures of 80–250 °C and 110 bar pressure. A control experiment was also run on the clay-rich shale sample by reacting it with a brine solution at 250 °C and 40 bar (vapor pressure of water).

2. Geological setting The study location is a proposed CO2 storage site located in the Central Tertiary Basin (CTB) close to the town of Longyearbyen on Svalbard (Fig. 1). The Central Tertiary Basin consists of a broad NNW-SSE-trending syncline bounded in the east by the Lomfjorden-Agardbukta Fault Zone and in the West by the deformation belt of the West Spitsbergen Orogeny (Harland et al., 1997). The basin formed as a foreland depression in front of an early Tertiary fold and thrust area in the West Spitsbergen Orogenic Belt. The sedimentary succession includes Carboniferous to Paleogene strata inside the basin that overlies a Caledonian basement currently exposed on the west coast and in northern Spitsbergen (Harland et al., 1997). Most of the stratigraphic units in the Central Tertiary Basin are continuous throughout the Barents Shelf. This is because Svalbard represents a more uplifted part of the submerged Barents Shelf (Nøttvedt et al., 1993). The target reservoir is in the Kapp Toscana Group: Late TriassicEarly Jurassic mainly continental sandstones and dense marine deposits. It is overlain by the Adventdalen Group. The Adventdalen, of Jurassic-Early Cretaceous age, is continuous across the Barents Sea Shelf to the Bjarmeland Platform, around the Loppa High and into the Hammerfest and Nordkapp basins (Dallmann, 1999). It is dominated by dark marine mudstones, but also includes deltaic

and shelf sandstones as well as thin carbonate beds. Part of the Adventdalen Group, The Janusfjellet Subgroup, overlies and serves as a caprock for the target reservoir. This Subgroup is a marine shelf to prodeltaic succession dominated by shales containing layers of subordinate siltstone and sandstone (Dypvik et al., 1991). The entire Subgroup was deposited during the Bathonian to Hauterivian time-span in shelf environments represented by a ca. 400 m thick sedimentary succession (Dypvik, 1984). The entire island has been uplifted by 3 km in the late Cenozoic (Nøttvedt et al., 1993). As a result, present day sediments close to the reservoir were buried to 3000 m relative to current depth of 900 m (Vågnes and Amundsen, 1993). This has resulted in a very low porosity and rather brittle shale layers. At the drilling site (DH2), where the two samples were collected, the caprock (Janusfjellet Subgroup) is composed of alternating layers of homogenous marine shale, and silty/sandy shale layers (Fig. 1). Samples for the experiment were collected from two layers of homogenous clay-rich shale and carbonate-rich shale, both with very low porosities of 2–3%. 3. Experimental approach to rock–brine–CO2 interaction In order to investigate the effect of mineralogical compositions on the reactivity of shale, two samples with different mineralogical composition were reacted in a batch reactor in a CO2–brine environment (Table 1). The experiments were run at higher temperature (80–250 °C) than the current reservoir temperature (30 °C) in order to speed up the rate of reactions. This is because the rate of silicate mineral reactions increases with increased temperature (Chen and Brantley, 1997; Gislason and Oelkers, 2003; Aagaard and Helgeson, 1982). For example, similar reactivity studies were conducted at 80– 150 °C and 150 bar (Credoz et al., 2009), 200 °C and 200 bar (Kaszuba et al., 2003, 2005) and 105 °C and 90 bar (Gunter et al., 1997). However caution must be practiced when interpreting high-temperature reactions. In addition, the reactive surface area of the caprock material was increased by crushing it to <500 lm grain size in order to increase the reactivity. A brine solution of 1 M NaCl was used as a substitute for formation water in order to encourage the dissolution process, as the presence of more ions in the initial brine solution, especially divalent cations, have a pH-buffering effect and decrease the acid-induced reactions. The water rock ratio was also increased by as much as 2500 times over geological storage conditions. Further enhancement of dissolution was guaranteed by continuous stirring. 4. Methods 4.1. Material Two core samples, a carbonate-rich shale with high calcite cementation and a clay-rich shale were chosen from Drill hole No. 2 at Flyplassveien (N78°140 1000 –E15°30 3400 ) on Longyearbyen. Both samples were taken from The Adventdalen Group and Janusfjellet Subgroup that have been identified as potential caprocks for the CO2 storage reservoir. The carbonate-rich shale was taken from a depth of 808 m which is close to the sandstone reservoir. It is composed mainly of carbonates, plagioclase, clay and quartz. The clay-rich shale was collected from different strata in the same well at 789 m below the surface. The clay-rich shale consists mainly of clays, quartz, plagioclase (albite dominated), and a very small amount of carbonates. The clay fraction in both samples is dominated by illite and chlorite. Details of the mineralogical arrangement within the clay matrix are shown in the thin section image of the cores (Fig. 2). The carbonate-rich shale has alternating carbonate-rich layers of lm scale having a heterogeneous matrix compared to the clay-rich sample which appears to be dominated

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Fig. 1. (a) Geological Map of the proposed Longyearbyen CO2 storage site (Ogata et al., 2011). Top right corner: geological map of Svalbard (yellow = Central Tertiary Basin) modified from Dallmann (1999). (b) Well correlation between the four wells drilled at the proposed CO2 storage site (Ogata et al., 2011). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)

Table 1 List of experiments and experimental parameters. Sample

Experiment

Carbonate-rich shale

Experiment Experiment Experiment Experiment Experiment Experiment Experiment Experiment

Clay-rich shale

a b

1Aa 1Ba 1Ca 2Aa 2Ba 2C1b 2C2b 2C3b

Starting fluid

Solids (g)

Water/rock ratio

Temperature (°C)

Total pressure (bars)

Duration (days)

1M 1M 1M 1M 1M 1M 1M 1M

12.5 1.50 1.50 12.50 12.50 1.50 1.50 1.50

20 20 20 20 20 20 20 20

250 250 250 250 250 200 150 80

110 110 110 110 39.5 110 110 110

7 14 35 35 35 21 21 21

NaCl + CO2 NaCl + CO2 NaCl + CO2 NaCl + CO2 NaCl NaCl + CO2 NaCl + CO2 NaCl + CO2

Batch reaction experiment. Parallel batch reaction experiment.

by a homogeneous clay matrix with random deposition of other minerals such as quartz, pyrite and albite/plagioclase. Chemical composition determined by X-ray fluorescence (XRF) and X-ray diffraction (XRD) mineralogical analysis of the samples is given in Tables 2 and 3, respectively. 4.2. Experimental set-up and procedures Two set-ups were used for the experiments: a Parr™ titanium reactor (Fig. 3A) and a HastelloyÒ C-276 reactor (Fig. 3B). The crushed rock, 100% passing through a 500 lm sieve, was rinsed with distilled water and dried in an oven at 80 °C for 48 h prior to the experiment. The crushed rock has a grain size which was 10% (500–250 lm), 10% (125–250 lm) and 10% (125–63 lm) and the remaining 70%, less than 63 lm, for all samples reacted. Before the start of all experiments, a mixture of rock and brine solution (1 M NaCl, Merck analytical grade) with a water/rock ratio of 20:1was poured into the reactor and then the mixture was bubbled

with (N2) to remove O2 from the solution, in order to imitate the conditions found in geological reservoirs. The system was then pressurized with CO2 from a commercially supplied bottle (99.9%) and allowed to equilibrate with the mixture for 24 h before heating it to the desired temperatures. In the experiments with CO2–brine, a pressure of 110 bars was achieved by adjusting the initial CO2 pressure before heating (batch reactor) and by using syringe pumps to maintain constant pressure at the experimental temperature (parallel reactor). The system was closed for the duration of the experiment; therefore, no sampling was carried out until the end of the experiment. In all experiments particles in the solution were kept in suspension by turbulent stirring at a speed of 100 rpm. There was a ±1 °C variation in temperature while the pressure stayed within ±3 bar and, once stabilized, stayed relatively stable throughout the experiment. A total of 8 batch reactions were run on the two shale samples at different sets of conditions. A summary of the experimental parameters is given in Table 1.

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Fig. 2. Backscattered Scanning Electron Microscopy (BSEM) image of thin-section depicting mineral distribution pattern in the clay matrix. (A) carbonate-rich clay, (B) magnified portion of A as indicated by the white box and (C) Clay-rich shale. N.B. The thin section indicated in A was 1 cm from the material used in the reactivity study and appears to have a higher ankerite to calcite ratio.

Table 2 Chemical composition by X-ray fluorescence. Sample

SiO2

Al2O3

Fe2O3

MnO

MgO

CaO

Na2O

K2O

TiO2

P2O5

L.O.I

SUM

Carbonate-rich shale Clay-rich shale

37.2 57.1

10.6 15.9

4.22 10.1

0.4 0.2

2.5 1.8

19.3 0.91

0.89 1.17

2.5 3.2

0.5 0.7

0.1 0.4

20.2 6.41

98.6 97.9

LOI = loss on ignition.

Table 3 Mineralogical composition by X-ray diffraction (semi-quantitative). Minerals

Carbonate-rich shale

Clay-rich shale

Quartz Plagioclase/albite  Ankerite Calcite Pyrite Siderite Illite Chlorite

13 6 7 29 – – 26 19

26 8 – – 1 5 22 38

All values are percentage of total dry weight. Semi-quantification is based on calculating the integrated peak area of respective mineral phases, multiplied by in-house calibrated/published weight factors (Peltonen et al., 2008). All expanding-clays were quantified as smectite whereas 7 Å (chlorite 002) and 10 Å (interpreted as illite) were used for calculating peak area for chlorite and illite.   Plagioclase is the dominant mineral in carbonate-rich shale and albite in the clay mineral-rich shale.

4.3. Analytical methods Liquid and solid samples were collected at the end of the reaction experiment after quenching and releasing the pressure in the

cell. The solution was filtered through a 0.45 lm filter in order to separate the solution from the solids. The liquid samples for metal analysis were acidified to a pH 2 with pure HNO3 in order to prevent the precipitation of metals. Dissolved SiO2 was determined using a Seal-AutoAnalyzer 3™, while basic metals were determined with inductively coupled plasma mass spectrometry (ICPMS) and dissolved anions by ion chromatography (IC). All reported pH and alkalinity measurements were taken 3 h after the samples were equilibrated with atmospheric PCO2 at room temperature (22 °C). Solid samples were washed with distilled water on a 0.45 lm filter and then oven dried at 80 °C for 48 h before they were analyzed. Solid samples were analyzed by XRD, and scanning electron microscopy (SEM). XRD recordings were acquired by Philips X’Pert MPD, using Cu Ka radiation generated at 40 kV and 50 mA. Semiquantification was based on calculating the integrated peak area of the respective mineral phases, multiplied by in-house calibrated weight factors (Peltonen et al., 2008). In order to detect and separate K-feldspar from plagioclase, a slow XRD scan between 26 and 28.5° 2h was run on the bulk samples, but K-feldspar was not detected in all the samples. A slow scan on an oriented clay specimen between 24 and 26° 2h was used to differentiate between kaolinite 002 (3.58 Å) and chlorite 004 (3.52 Å) and there was no kaolinite in

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Fig. 3. Experimental set-up for batch reaction (a) batch reactor, (b) parallel reactor.

any of the samples analyzed. All expanding clays were identified as smectite. The 7 Å peak (chlorite 002) was used for chlorite quantification and 10 Å peak (interpreted as illite) for quantification of illite. No XRD analysis of the <2 lm size fraction was run on reacted samples from experiments 1B, 1C, 2C1, 2C2, 2C3 (Table 1). This is because there was not enough material left to separate the clay fraction as a small amount of material (1.5 g) was reacted. However total clay content in all bulk XRD samples was estimated using the 4.5 Å peak and in-house/published weight factor (Peltonen et al., 2008). Total C (TC), total organic C (TOC) and total inorganic C (TIC) content of solids were determined using acid treatment and a LECOÒ CR-412 Carbon Analyzer for both the starting and the reacted material. 4.4. Geochemical modeling Geochemical modeling was carried out with PHREEQC-2 (version 2.16.02, Parkhurst and Appelo, 1999, using LLNL.dat database). Equilibrium speciation calculations were made for all experiments conducted on the two samples. The calculations were made on the basis of the measured final solution composition of fluids. In addition, a predictive kinetic calculation was also run for the carbonate-rich shale that was reacted with CO2–brine for 1, 2 and 5 weeks at 250 °C and 110 bar. The simulation was run in two steps; first, by equilibrating the minerals with 1 M NaCl at 250 °C, and later, introducing CO2 into the system. The CO2 pressure was introduced by equilibrating the solution from the first step with a CO2 partial pressure of 70 bar, i.e. water vapor pressure at 250 °C (40 bar) minus total pressure (110 bar). Accordingly, the fugacity coefficient of CO2 was estimated using the Soave–Redlich–Kwong (SRK) equation of state (Soave, 1972). The simulation was run for a period of 35 days. The amount of minerals/kg of water was determined on the basis of a water/rock ratio of 20:1 and mineral abundances were expressed as percentages of total solids. Under acidic pH, the contribution of CO2 is mainly limited to the supply of protons and has no inhibiting or accelerating effect in dissolution of silicate minerals (Brady and Carroll, 1994; Golubev et al., 2005; Pokrovsky and Schott, 2004). Pokrovsky et al. (2009) also suggested that the effect of dissolved CO2 on carbonate mineral reactivity to be of secondary importance compared with  pH, dissolved CO2 3 and HCO3 ions. Hence, a general form of rate law as a function of proton activity, derived from the transition state theory (Lasaga, 1984; Aagaard and Helgeson, 1982), was used for the calculation of acidity-driven kinetic dissolution and precipitation.

   Qm Ratem ¼ Am kðTÞm ðaHþ Þn 1  Km

ð1Þ

where Rate depicts the dissolution/precipitation rate (positive values correspond to dissolution, negative values to precipitation), m represents the mineral index, A denotes the reactive surface area per kg of water, k(T) refers to the temperature-dependent rate constant, aH+ stands for the proton activity, n is the order of the reaction with respect to protons, km represents the equilibrium constant for the mineral water reaction and Q designates the corresponding ion activity product. The rate constant at 250 °C was calculated from rate constants at 25 °C using Arrhenius law.

kðTÞ ¼ k25 exp

   Ea 1 1  T 298:15 R

ð2Þ

where Ea represents the activation energy (in J mol1), k25 designates the rate constant at 25 °C (in mol m2 s1), R is the gas constant (8.314 J mol1 K1) and T is the absolute temperature. The kinetic rate constants for the selected minerals were taken mostly from Palandri and Kharaka (2004) and other sources, as indicated, and recalculated using Eq. (2) for 250 °C (Table 7). Similar rate constants were used for both dissolution and precipitation and all minerals were allowed to precipitate. Due to the lack of data on rate constants for analcime and ankerite in the literature, heulandite (Ragnarsdottir, 1993) and dolomite (Palandri and Kharaka, 2004) rate constants were used as a substitute for analcime and ankerite rate constants, respectively. The power value of the rate expression n was assumed to be 0.5 for most of the minerals as the values reported in the literature for the same mineral are not consistent. A spherical grain shape was assumed for all minerals for the calculation of the surface area of each mineral phase. A diameter of 63 lm was assumed for all minerals except clays, where a spherical grain size of 2 lm was used. Grain surface area and molecular volume were used to calculate the specific surface areas used in the rate calculation. 5. Results 5.1. Fluid chemistry 5.1.1. Carbonate-rich shale: reaction with CO2–brine at 250 °C Ions (Ca, SiO2 K, Mg, Fe and Al) previously absent in the starting 1 M NaCl solution, appeared in all three experiments at 7, 14 and 35 days, indicating dissolution of reacting solids (Fig. 4). The concentration of Na in the brine in all three batch reactions did not show any significant variations outside of the analysis margin of error (Table 4). Calcium concentrations in solution reached 11.34 mmol/kg after 1 week of reaction, but it decreased by 20% in the sample reacted for 2 weeks, and then increased to reach

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the 10% margin of error for the Na analysis. The anions in solution did not show significant variation from the starting composition. 5.1.3. Clay-rich shale: reaction with CO2–brine at 80, 150 and 200 °C It is clearly demonstrated in Fig. 6 that the release of Fe and SiO2(aq) was more significant in lower temperature (lower pH) reactions. This resulted in an increase in the concentration of Fe (250%) and SiO2(aq) (118%) in solution at 80 °C compared to the material reacted at 200 °C. The Na concentration increased by 18% from the starting concentration in the two samples reacted at 150 and 200 °C. This could be due to loss of water through leakage. The leakage or dissolution of water into the CO2 phase could go unnoticed as the system was kept at constant pressure by pumping CO2 into the reactors (Fig. 1B). Therefore, the observed relative increase in the concentration of Fe and SiO2(aq) at 80 °C might be lower than it appears, because of the concentration effect. 5.2. Reaction of solid phases

Fig. 4. Brine chemistry as a function of reaction time at 250 °C for the carbonaterich shale reacted with CO2–brine.

9.77 mmol/kg in the sample reacted for 5 weeks. Magnesium and Fe in solution remained in micromolal quantities, but almost doubled in samples reacted for 5 weeks compared to samples reacted for 1 and 2 weeks. Aluminum concentration was less than 100 lmol/kg in all three samples, with the lowest measured in the sample reacted for 5 weeks. Potassium and SiO2 increased in concentration almost linearly with increase in reaction time. 5.1.2. Clay-rich shale: reaction with CO2–brine and CO2-free brine at 250 °C As shown in Fig. 5, the release of major ions into solution varied, depending on the reacting fluids. In solutions from reactions with CO2–brine, the concentration of Fe and SiO2(aq) were almost double, whereas concentrations of Ca, Mg, K were lower by 25, 36 and 72%, respectively, compared to the solution from the sample reacted with CO2-free brine. The concentration of Na in the CO2–brine experiment decreased by 5.65% and in the CO2-free brine experiment by 6.52% from the starting concentration, which is within

5.2.1. Carbonate-rich shale: reaction with CO2–brine at 250 °C The content of solid minerals changed extensively at the end of experiments run for 1, 2 and 5 weeks. The main changes involved the dissolution and re-precipitation of carbonates, the dissolution of plagioclase, and the formation and dissolution of clay minerals. Both calcite and ankerite were partially dissolved and re-precipitated as smaller calcite crystals with the newly formed clay minerals (Fig. 8). Ankerite dissolution reached its maximum in the sample reacted for 5 weeks (Fig. 7). Semi-quantitative analysis by XRD on bulk samples showed a decrease in ankerite composition from 7% wt of the total rock at the start of the experiment, to close to 1% wt after reacting for 5 weeks. The same analysis also revealed an increase in calcite weight percentage from 29% at the start of the experiment, to close to 40% of total solids after 5 weeks. This was also confirmed by the increase in total inorganic carbon content (TIC) of the solids, which increased by 0.58% of total weight in the 5 week-batch (Table 6). Comparison of TIC measured from C analysis with TIC calculated from XRD semi-quantification was also in good agreement (Fig. 12). Analysis of <2 lm size oriented clay diffractograms after reaction revealed a significant alteration compared with the starting material. As can be seen in Fig. 9, chlorite and illite, the major components of clay minerals in the original sample, converted partly into expanding clay. Chlorite appears to have dissolved to a greater extent, while a significant amount of illite was also dissolved to form smectite. According to semi-quantification based on XRD, almost 2=3 of the total clay minerals in the solids had transformed

Table 4 Brine chemistry for all experiments (mmol/L). Minerals

Temperature Na⁄ Ca Fe Al K Mg SiO2 Alkalinity (meq/kgw)

Carbonate-rich shale

Clay-rich shale

B + CO2 (1-week)

B + CO2 (2-weeks)

B + CO2 (5-weeks)

B + CO2 (5-weeks)

B (5-weeks)

B + CO2 (3-weeks)

B + CO2 (3-weeks)

B + CO2 (3-weeks)

250 °C 994.35 11.34 0.17 0.10 2.13 0.58 1.83 8.00

250 °C 949.04 9.07 0.17 0.06 2.40 0.42 3.18 5.14

250 °C 975.52 9.97 0.24 0.02 4.73 0.84 6.56 7.32

250 °C 943.48 2.85 0.39 0.05 4.61 0.27 5.86 6.50

250 °C 934.78 3.69 0.21 0.02 6.82 0.64 3.05 0.60

80 °C 984.35 2.46 3.52 0.09 2.36 1.95 3.31 16.25

150 °C 1185.65 2.70 1.35 0.11 0.89 1.97 2.27 18.96

200 °C 1181.74 1.83 1.00 0.01 3.66 1.01 1.53 4.35

Initial solution = 1 M NaCl (Merck analytical grade), average analysis error 2r (7%) Higher analysis error for sodium due to high dilution 2r (10%)

⁄

B + CO2 = reacted with brine and CO2. B = reacted with CO2-free brine.

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Table 5 Saturation indices (SI) for the major minerals involved in the reaction, based on calculation made on the final solution composition at reaction temperature and experimental total pressure of 110 bar ðPCO2 Þ þ PðH2 OÞ. Minerals

Temperature Albite Analcime Ankerite Anorthite Calcite Chalcedony Clinichlore-14A Dolomite Illite Kaolinite K-feldspar Magnesite Quartz Siderite Smectite pHa pH Alkalinitya (meq/kgw) P CO2(mol/kg)

Carbonate-rich shale

Clay-rich shale

B + CO2 (1-week)

B + CO2 (2-weeks)

B + CO2 (5-weeks)

B + CO2 (5-weeks)

B (5-weeks)

B + CO2 (3-weeks)

B + CO2 (3-weeks)

B + CO2 (3-weeks)

250 °C SI 0.43 0.42 5.92 0.45 0.60 0.67 0.05 1.36 0.56 0.19 1.93 0.09 0.52 0.43 1.13 7.70 5.33 8.00 0.64

250 °C SI 0.05 0.17 6.66 0.49 0.21 0.43 1.57 0.53 0.04 0.54 1.83 0.54 0.28 0.15 1.13 7.80 5.16 5.14 0.64

250 °C SI 0.53 0.04 5.88 0.78 0.52 0.12 1.01 1.14 0.06 0.05 0.61 0.04 0.04 0.56 2.32 6.50 5.31 7.32 0.65

250 °C SI 0.79 0.32 6.26 0.61 0.08 0.16 0.35 0.58 0.63 0.56 0.36 0.35 0.01 0.92 2.56 7.60 5.36 6.50 0.65

250 °C SI 0.52 0.66 7.02 1.74 0.19 0.48 13.33 0.26 3.06 4.07 1.5 0.4 0.33 0.14 5.83 7.58 7.41 0.60 –

80 °C SI 2.38 1.43 4.61 4.32 2.22 0.53 19.21 3.09 5.16 6.20 2.30 2.21 0.76 2.80 2.74 8.10 4.21 16.25 1.06

150 °C SI 2.02 1.59 3.10 0.54 0.88 0.12 4.97 0.37 4.72 5.17 1.12 0.58 0.07 0.24 1.93 7.70 4.78 18.96 0.73

200 °C SI 0.67 0.71 5.43 3.28 1.13 0.54 6.54 1.02 0.56 0.69 1.81 0.85 0.37 0.48 0.44 7.90 4.76 4.35 0.69

P SI. pH and CO2 are calculated on the basis of the composition of the final solution at reaction temperatures and their respective CO2 fugacities using PHREEQC-2 (version 2.16.02. Parkhurst and Appelo, 1999, using LLNL.dat database). B + CO2 = reacted with CO2–brine. B = reacted with CO2-free brine. SI = log(IAP/K), where IAP = ion activity product and K = equilibrium constant. a Measured.

Fig. 5. Brine chemistry of the clay-rich shale reacted with CO2–brine and CO2-free brine for 5 weeks at 250 °C.

into expanding clay (smectite) after 5 weeks of reaction. The presence of smectite was also confirmed by qualitative analysis using visual appearance of the clays from SEM images (Fig. 8) and the characteristic SEM-EDS spectrum. The expanding clay has been interpreted as trioctahedral smectite given that its d (060) reflection is close to 1.54 Å. There were no traces of other forms of carbonates, such as magnesite, siderite or dolomite, no K-feldspar, kaolinite or any trace of zeolites that could be detected either by XRD or by qualitative X-ray analysis using the EDS of the reacted powders in any of the experiments.

Fig. 6. Brine chemistry as a function of reaction temperature for the clay-rich shale reacted with CO2–brine.

5.2.2. Clay-rich shale: reaction with CO2–brine and CO2-free brine at 250 °C The only observable difference in mineralogical composition between the sample reacted with CO2–brine and the original material is the disappearance of the siderite peak, which implies dissolution. The same dissolution of siderite was also observed in the material reacted with CO2-free brine. According to semi-

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B.L. Alemu et al. / Applied Geochemistry 26 (2011) 1975–1989

Fig. 7. Bulk XRD plot of the carbonate-rich shale, unreacted, reacted for 1, 2, and 5 weeks with CO2–brine at 250 °C. Minerals identified by distinctive peaks and the results depict dissolution of ankerite and increase in calcite.

Fig. 8. SEM picture of the carbonate-rich shale showing calcite crystal evolution and precipitation of secondary clay minerals (smectite). (A) calcite crystal before reaction, (B) calcite crystal (larger) after 2 weeks of reaction (C) re-precipitation of smaller (finer) calcite with clay minerals after 5 weeks (D) calcite crystal growth with clay minerals after 5 weeks.

quantitative XRD analysis on bulk reacted-solids, analcime composition reached as much as 1/5 of the total solids by the end of 5 weeks in the material reacted with CO2-free brine (Table 6). SEM and Qualitative X-ray analysis using EDS also confirmed this (Fig. 10B). In contrast, no analcime was detected on the material reacted with CO2–brine. As opposed to the carbonate-rich shale, there were no significant changes in the clay minerals after

reaction in either case except for a slight indication of mixed clay formation in the sample reacted with CO2-free brine. Total inorganic C (TIC) analysis run on both reacted materials showed a decrease of 0.42% and 0.45% of total weights of the sample, respectively, for material reacted with CO2–brine and CO2-free brine (Table 8). There was no new trace of other carbonates in either of the reacted samples.

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B.L. Alemu et al. / Applied Geochemistry 26 (2011) 1975–1989 Table 6 Mineralogical composition by X-ray diffraction for all experiments after reaction (semi-quantitative). Minerals

Carbonate-rich shale B + CO2 (1-week)

B + CO2 (2-weeks)

Clay-rich shale B + CO2 (5-weeks)

Temperature 250 °C 250 °C 250 °C Ankerite 5 6 1 Calcite 29 26 40 Plagioclase/albitea 13 8 5 Pyrite – – – Quartz 18 17 11 Siderite – – – Zeolites (analcime) – – – Illite N/A N/A 26 Chlorite N/A N/A 9 Smectite N/A N/A 65 B + CO2 = reacted with brine and CO2 mixture B = reacted with CO2-free brine. All values are percentage of total dry weight

B + CO2 (5-weeks)

B (5-weeks)

B + CO2 (3-weeks)

B + CO2 (3-weeks)

B + CO2 (3-weeks)

250 °C – – 9 1 23 – – 42 58 –

250 °C – – 8 – 23 – 23 49 51 –

80 °C – – 8 1 26 5 – N/A N/A N/A

150 °C – – 8 1 26 5 – N/A N/A N/A

200 °C – – 8 1 26 5 – N/A N/A N/A

All values are percentage of total dry weight. N/A = not analyzed. Semi-quantification is based on calculating the integrated peak area of respective mineral phases, multiplied by in-house calibrated/published weight factors (Peltonen et al., 2008). All expanding clays were quantified as smectite whereas 7 Å (chlorite 002) and 10 Å (interpreted as illite). a Plagioclase is the dominant mineral in carbonate-rich shale and albite in the clay mineral-rich shale.

Fig. 9. XRD plot of oriented clay fraction of the carbonate-rich shale, as prepared, treated with ethylene glycol, and heated to 550 °C. (A) un-reacted material, (B) after 5 weeks reacting with CO2–brine showing the formation of expanding clay (smectite) at 14.3 Å.

5.2.3. Clay-rich shale: reaction with CO2–brine at 80, 150 and 200 °C Comparing XRD analysis of the three reacted solids, no mineralogical changes could be detected in any of the three experiments. There was no significant amount of carbonates formed under any of the experimental conditions nor was there any alteration of starting solids that could be detected by bulk XRD or by qualitative X-ray analysis using EDS. 5.3. Geochemical modeling Saturation indices (SI) of primary and secondary minerals, calculated on the basis of the composition of the final solutions, are given in Table 5. The solutions from carbonate-rich shale reacted with CO2–brine at 250 °C was undersaturated with respect to ankerite at all the 3 experimental stages of 1, 2 and 5 weeks. However, solutions from these three stages were saturated with respect to calcite, dolomite and siderite while only the sample reacted for

5 weeks was saturated with respect to magnesite. Given that ankerite dissolved and re-precipitated as calcite in these experiments, the calculated saturation indices are consistent with the observed mineral evolution on the reacted solids. In addition, these three solutions have higher saturation state of smectite compared to other clay minerals such as illite and chlorite. This is in agreement with the observed results where smectite precipitated after dissolution of chlorite and illite. Furthermore the solution from the sample reacted for 5 weeks was undersaturated with respect to kaolinite well in agreement with XRD results where no trace of kaolinite were detected. The observed higher saturation of albite over other feldspars and analcime is also consistent with the mineralogical composition of the reacted solids. The solution from the clay-rich shale reacted with CO2–brine at 250 °C for 5 weeks was saturated with respect to calcite, dolomite and siderite while being undersaturated with respect to ankerite and magnesite. The same solution is also saturated with respect to chlorite, illite, kaolinite

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Fig. 10. (A) Comparison of XRD (bulk) plot of the clay-rich shale unreacted, reacted with CO2–brine and reacted only with CO2-free brine depicting analcime formation only in the material reacted with CO2-free brine. (B) SEM picture of analcime precipitation with clay minerals in the clay-rich shale reacted with CO2-free brine.

Table 7 Kinetic rate parameters used in geochemical modeling. Mass (mmol/kg w)a

Rate logK25 (mol m2 s1)

Rate logK250 (mol m2 s1)

Ea (J mol1)

n

Surface area (m2 g1)b

Source of kinetic rate data

110

13.41

6.56

90.90

0

3.60e02

8.96

12.00

7.10

65.00

0.50

3.63e02

Anorthite*

2.24

8.00

6.75

16.60

1.41

3.46e02

Calcite

147

0.30

0.79

14.40

0.50

3.51e02

Clinochlore14A Ankeritec

15.8

12.80

8.28

60.00

0.50

1.14e+00

Palandri and Kharaka (2004) Palandri and Kharaka (2004) Palandri and Kharaka (2004) Palandri and Kharaka (2004) Brandt et al. (2003)

15.8

3.50

0.78

36.10

0.20

2.97e02

Siderite Illite Kaolinite

– 0,03 –

7.75 12.40 11.31

4.13 8.93 6.34

48.00 46.00 65.90

0.50 0.10 0.20

2.41e02 4.61e01 1.16e+00

Smectite-highFe–Mg Analcimec K-feldspar

–

10.98

9.20

23.60

0.50

1.04e+00

– –

13.10 10.06

8.73 6.16

58.00 51.70

0.50 0.50

4.20e02 3.73e02

Magnesite

–

6.38

5.29

14.40

0.50

3.17e02

Quartz Albite

*

Palandri and Kharaka (2004) Golubev et al. (2009) Köhler et al. (2003) Palandri and Kharaka (2004) Palandri and Kharaka (2004) Ragnarsdottir (1993) Palandri and Kharaka (2004) Palandri and Kharaka (2004)

*

Plagioclase introduced as 4:1 mixture of end-members albite and anorthite. Calculated on the basis of percentage mass, water rock ratio of 20:1, and molecular weight of individual minerals. Surface area calculated assuming spherical grains with a diameter of 2 lm for all clay minerals and 63 lm for the rest of the minerals and using their respective molecular volumes. c Substitute rate constants used i.e., Heulandite rate for analcime and dolomite rate for ankerite. a

b

and smectite. Smectite has the higher saturation state in this solution even though no evidence of smectite precipitation has been observed. Likewise albite also has the higher saturation state compared to other feldspars and analcime. The solution from the clayrich shale reacted with CO2-free brine at 250 °C for 5 weeks was only saturated with respect to dolomite and siderite. Chlorite and smectite have higher saturation state in the same solution where illite and kaolinite appeared to be undersaturated. There was a small indication of expanding material observed on the clay minerals analyzed from the same reaction, which could be related to the formation of chlorite/smectite mixtures. Illite appeared to have dissolved much as depicted by the higher concentration of K. The observed undersaturation might, therefore, indicate that illite dissolution would continue if the experiment had been allowed to run for longer periods. This solution was also undersaturated with respect to the feldspars and analcime. Solutions from the clay-rich shale reacted with CO2–brine at 80, 150 and 200 °C for 3 weeks were all undersaturated with respect to the carbonates except for

the solutions from the 150 and 200 °C experiment which were also saturated with respect to siderite. Solutions from these 3 experiments were undersaturated with respect to chlorite and smectite (except sample reacted at 150 °C) and saturated with respect to kaolinite and illite (except sample reacted at 150 °C). Both albite and K-feldspar seemed to be more stable at lower temperatures but anorthite was undersaturated in all the three solutions. Analcime also had higher saturation at lower temperatures though no traces of it were detected on the bulk XRD analysis. The predicted fluid composition evolution, and saturation indices of the carbonate-rich shale from kinetic reaction modeling are given in Fig. 13. The simulation results indicate a reasonable match for all elements analyzed except for Mg, which appears to be far from measured values. The respective saturation indices from the same simulation also partially agree with the observed experimental values. The system appears to be supersaturated with respect to smectite and close to saturation with respect to analcime which is in moderate agreement with the observed experimental results.

B.L. Alemu et al. / Applied Geochemistry 26 (2011) 1975–1989

Fig. 11. Activity diagram for albite–analcime–kaolinite–pyrophyllite system at temperature of 250 °C and 110 bar total pressure ðPCO2 Þ þ PðH2 OÞ. Pyrophyllite is an endmember and proxy for montmorillonite (Aagaard and Helgeson, 1983). Carbonate-rich shale reacted with CO2–brine (d), clay-rich shale reacted with CO2–brine (j) and clay-rich shale reacted only with CO2-free brine (N). Activities are calculated on the basis of the final composition of the solution, using PHREEQC2 (version 2.16.02, Parkhurst and Appelo, 1999, using LLNL.dat database).

The simulation depicts the system to be supersaturated with most of the carbonate which has the same trend as the calculated saturation states except for ankerite which is highly unsaturated in the reacted solutions. In addition, both illite and kaolinite have lower saturation states compared to measured compositions. On the other hand the super saturation of chlorite and smectite is fairly in agreement with experimental values. Generally, the model shows a relatively rapid dissolution and precipitation compared with the experimental findings. This is shown clearly in Fig. 13B as the saturation states were established within the first few days and changed little afterward.

6. Discussion The results of the 8 experiments provided data that contributes to the qualitative understanding of fluid rock interactions under geochemical settings of subsurface CO2 storage. Acid-induced reactions between rock forming minerals and CO2–brine were clearly observed. The carbonate-rich shale showed a relatively high reactivity compared to the clay-rich shale. Dissolution of carbonates was significant in the carbonate-rich shale where both calcite and ankerite dissolved and precipitated as calcite. The dissolution of ankerite was clearly identified both in reacted solids and solution. This dissolution of ankerite and reprecipitation is documented by the fact that the final solutions of all carbonate-rich shale experiments were undersaturated with respect to ankerite and saturated with respect to calcite (Table 5). Furthermore, the reacted solids from the same sample have much less ankerite. In addition to ankerite, calcite also appeared to have dissolved and re-precipitated during the experiment. This was revealed by the dissolution of larger calcite crystals (Fig. 8A) that were abundant on the unreacted material which were later replaced by the second generation of finer calcite crystals after reaction (Fig. 8C). The CO2 consumed by the dissolution of carbonates is released when the carbonate finally precipitates. However, the Ca released from silicate mineral dissolution reacts with dissolved CO2, trapping it permanently. The latter process was observed in

1985

Fig. 12. Plot of total inorganic C (TIC) calculated from XRD semi-quantification against TIC measured (C analysis). The carbonate-rich shale (unreacted) (j), clayrich shale (unreacted) (d), carbonate-rich shale reacted with CO2–brine (h) clayrich shale reacted with CO2–brine (N), clay-rich shale reacted with CO2-free brine (s).

the carbonate-rich shale which showed a significant increase in calcite concentration in the reacted solids (Table 6). Such an increase in calcite is quite significant even after accounting for the amount of ankerite that dissolved and re-precipitated as calcite. This increase in calcite was also confirmed by the observed increase in TIC (Table 8). Dissolution of plagioclase was one of the major suppliers of Ca for the formation of new calcite as depicted in Table 6. Other cations in solution, such as Fe and Mg did not react with CO2 to form carbonates. The absence of magnesite is in agreement with its undersaturation in the final solutions from the reacted carbonate-rich shale. On the other hand, there was no significant amount of siderite precipitate that could be detected by XRD analysis even if the final solutions appeared to be saturated with siderite. Hence, Ca has a much greater potential to permanently trap CO2 compared to other common divalent ions at experimental conditions. This is one of the mechanisms by which caprock material can contribute in the mineralogical trapping of CO2. Aluminum appears to be conserved during the formation of clays as shown by the low concentration measured in the final solution in almost all of the experiments. Iron and Mg released by the dissolution of ankerite and chlorite might have been later consumed by the formation of trioctahedral smectite. Therefore, the formation of smectite might be one of the reactions responsible for the observed low concentration of Fe and Mg in the solution of the carbonate-rich shale reacted for 5 weeks. This reaction also consumes SiO2(aq) from the dissolution of silicate minerals such as plagioclase feldspars and clay minerals. The reverse reaction, i.e. the breakdown of smectite to illite through diagenetic-processes releases Fe and Mg and was implied to be one of the factors causing greater ankeritization of calcite at deeper levels in Wilcox (Eocene) sandstones in SW Texas (Boles, 1978). In the caprock, at the pore scale, the limited solute exchange to and from the caprock might lead to higher activity of silica in the pore waters. This higher silica and H+ activity in solution might favor the stability of smectite over illite. Relatively similar results of formation of smectite were reported during experimental studies of clay formation in an acidic environment, chlorite to smectite (Senkayi et al., 1981) and illite to smectite (Vicente et al., 1977) under lower temperatures 100 °C and 50 °C, respectively. Other similar experimental studies of the reactivity of clayey caprock in a CO2–brine environment under hydrothermal

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Fig. 13. (A) Predicted solution composition evolution in comparison with measured experimental values for the carbonate-rich shale reacted with CO2–brine for 5 weeks. (B) Corresponding predicted mineral saturation indices. Kinetic simulation was made using PHREEQC-2 (version 2.16.02, Parkhurst and Appelo, 1999, using LLNL.dat database) at experimental total pressure of 110 bar ðPCO2 Þ þ PðH2 OÞ, and 250 °C.

Table 8 Total C (TC), total organic C (TOC) and total inorganic C (TIC) measured for unreacted and reacted materials from experiments run at 250 °C. Sample Carbonate-rich Carbonate-rich Clay-rich shale Clay-rich shale Clay-rich shale

shale shale reacted with CO2–brine reacted with CO2–brine reacted with brine only

TOC (%)*

TC (%)*

TIC (%)*

0.1383 0.0838 0.3165 0.4100 0.2188

4.878 5.402 0.7876 0.4617 0.239

4.7397 5.3182 0.4711 0.0517 0.0202

Measured only for samples reacted for 35 days. TIC = total inorganic carbon. TOC = total organic carbon. TC = total carbon. * Percentage calculated against total dry weight of the sample.

laboratory conditions, also reported precipitation of smectite (Credoz et al., 2009; Kaszuba et al., 2005; Wigand et al., 2008). This suggests that clay mineral reactions observed under higher temperature (250 °C) in the experiments, might also occur under geological storage conditions, provided that the solution compositions and reacting minerals favor the stability of smectite. In diagenetic processes the illitization of smectite is a common phenomenon, but during storage of CO2 in geological reservoirs, a reverse process might be one of several types of reactions that have the potential to develop in the caprock environment. This is comparable to weathering processes in low temperature environments, where precipitation of smectite after dissolution of illite and chlorite are known to occur (Carnicelli et al., 1997; Herbillon and Makumbi, 1975). However, in clean reservoirs (sandstone), the presence of enhanced solute exchange (higher porosity and permeability) in addition to a low supply of silica, due to the less-reactive nature of the dominant mineral, i.e. quartz, most often leads to lower silica activity. Such low activity of silica might favor the stability of illite over smectite (Sass et al., 1987; Aagaard and Helgeson, 1983). The low reactivity observed in the clay-rich shale might be related to the relatively low content of carbonates in the sample. The fact that no carbonate formed during the reaction of the

clay-rich shale can be attributed to the much lower concentration of Ca in solution. Albite did not dissolve in all the experiments as depicted by the lack of change in composition in reacted minerals (Table 6). This is expected as albite is less reactive than Ca-rich plagioclase feldspars (Gaus et al., 2005). There was no significant dissolution of illite and chlorite and formation of expanding clays as opposed to the carbonate-rich shale, even though both were reacted at the same conditions. The clay-rich shale has relatively higher Fe (Fe2O3 = 10%) compared to the carbonate-rich shale (4%). High Fe concentration, especially in clay minerals such as chlorite, lead in most cases to higher reactivity (Ross and Kodama, 1976). However, Senkayi et al. (1981) suggested that high Fe concentration in chlorite might not lead to an increase in reactivity unless most of the Fe is located in the interlayer hydroxide sheets. If Fe is located in the octahedral structure of the 2:1 units, the reactivity will be slower. Hence, the observed lower reactivity could be related to the structural arrangement of Fe in chlorite, or to the lower supply of other metals such as Mg in solution, which plays an important role in the synthesis of smectite (Kloprogge et al., 1999; Nadeau, 1998). In the case of the carbonate-rich shale, Mg was supplied by the dissolution of ankerite and re-precipitation as calcite which allowed the formation of smectite. However, in the clay-rich shale, reacted at the same conditions, the supply of Mg was very limited as depicted by the composition of the final solution, which had 1=3 of the final concentration of the carbonate-rich shale (Table 4). Therefore, Mg appears to play the most important role in the formation of smectite. The fact that analcime did not form in samples reacted in the presence of CO2 as reported by Kaszuba et al. (2005), could be due to lower Na concentration in the starting fluid (5 N NaCl compared to 1 M NaCl solution in the present study), thus making the analcime unstable (Fig. 11). Higher H+ activity due to the dissolution of CO2 in solution, and increased SiO2(aq) activity (aSiO2) as a result of silicate mineral dissolution, might move the system out of the stable region for analcime. The solution composition of the clay-rich shale reacted with CO2-free brine favored the formation of analcime, compared with other reactions run at the same

B.L. Alemu et al. / Applied Geochemistry 26 (2011) 1975–1989

conditions. According to the phase diagram for the albite–analcime–kaolinite–pyrophyllite system (Fig. 11), formation waters that have low salinity and are supersaturated with silica will have a lower tendency to form analcime. In general, at similar Na+ concentrations, temperatures and pressures, kaolinite-smectite equilibrium requires a lower pH (high aHþ ) than does smectiteanalcime equilibrium (Hutcheon and Abercrombie, 1990). However, the stability of analcime in formation waters could also be influenced by reactions other than the activity of Na+, H+ and SiO2(aq), so all other factors should be taken into account. Experimental results from temperature variations on the clayrich shale showed that lower temperature, i.e., lower pH, resulted in a greater release of SiO2(aq) and Fe, during a relatively short time (Fig. 6). This is because for the same pressure, more CO2 dissolves in water at lower temperature. Hence, the solution at lower temperature would have lower pH which facilitated the dissolution of silicate minerals there by releasing more silica and Fe into solution. The fast release of these solutes in solution could be used as a monitoring tool to detect leakage of CO2 from deep reservoirs to the upper fresh water aquifers. Nevertheless, knowledge of the initial aquifer chemistry is essential to establish background values. The results of the geochemical modeling were in part successful in replicating some of the experimental results. There were some discrepancies such as the predicted low Mg concentration that was far from experimental results and the undersaturation of the solution with respect to illite which is also far from experimental values (close to equilibrium) after 5 weeks of reaction. Because few parameters were used in the model, fitting the model to experimental results was challenging. However, the results of the geochemical modeling are encouraging, especially considering the fact that thermodynamic properties of standard minerals were used, which do not always have the same properties as the minerals in the samples. For these reasons, this type of simulation study would certainly benefit from a larger database that incorporates various minerals which would help in modeling adequate solid solution models. This is especially true for clay minerals which behave uniquely depending on the presence of certain features even within the same group. Under geological storage conditions, the CO2 plume will be on top of the reservoir due to its lower density compared to the formation water thereby resulting in direct contact with the caprock. Due to limited porosity in shales, diffusion of CO2 into the caprock is often slow even at high concentration gradients and may be limited to thin basal layers of the caprock. For example, depth of penetration of CO2 by diffusion into the caprock was about 8 m, 3 ka after placement, based on numerical modeling (Gaus et al., 2005) and about 12 m, 70–80 Ma after placement based on C isotope studies of natural analogues (Lu et al., 2009). Molecular diffusion of CO2 is a continuous process and could be the rate-controlling step in fluid–rock reactions in a transport-limited environment. Most of the reaction would be taking place in the pore space of the caprock due to the acidity induced by the dissolution of CO2 in the pore water. The extent of this reaction also depends on how fast the reacted fluids (solutes) are able to move out of the caprock matrix. In contrast to previous understanding of significant solute exchange through and out of the caprock during diagenesis (e.g., Boles and Franks, 1979; Lynch et al., 1997; Towe, 1962), Thyberg et al. (2010) demonstrated that solute exchange, such as silica migration from caprock to reservoir, is extremely limited. According to Thyberg et al. (2010), silica often precipitates within the clay matrix. These secondary precipitates within the pore space might lead to a reduction in porosity. As a result, both diffusion rate and reaction processes might be further reduced, limiting the progress of fluid–rock interactions. Natural analogues of mudstone caprocks have demonstrated such a capability of limited reactivity upon exposure to CO2 for millions of years (Lu

1987

et al., 2009). Depending on rock porosity/permeability and mineral composition, the dissolution of minerals might also be the dominant process leading to a partial increase in porosity of the caprock. Such enhanced permeabilities were reported in studies of sealing capacity and fluid-transport on a carbonate-rich shale (marl), where the effective gas permeability increased by repeated diffusion of CO2 across a water wet sample (Wollenweber et al., 2010). The authors suggested that dissolution of carbonate and transportation during flooding as a main reason for the increase in permeability under laboratory experimental conditions (open system). However, in subsurface geological conditions, dissolution/precipitation reactions appear to be occurring within the pore space and they are best represented as closed system (Bjørlykke, 1983; Ehrenberg and Nadeau, 1989; Nadeau, 1998). In addition, formation of smectite in a carbonate-rich environment leads to reduction of permeability as the smectite bridges the pore space there by significantly reducing the permeability of the system (Nadeau, 1998). Furthermore, laboratory measured permeability measurements in clay-containing rocks appear to be significantly affected by drying and re-saturation processes due to the structural rearrangement of clay minerals in the pore space (Nadeau, 1998). Therefore, caution must be practiced in interpreting laboratory measured permeability in clay-rich samples to field scale. In general, both chemical reaction and transport systems in the caprock should be considered when one analyzes the integrity of reservoir seals for geological storage of CO2. 7. Summary and conclusions The experimental study suggests that caprock is an active participant in the reaction process in the geological storage of CO2 and plays a role in the trapping of CO2 into stable mineral forms such as calcite. Reaction of the mixture of shale and CO2–brine mixture was revealed by the changes in aqueous chemistry and dissolution of minerals and mineralogical alterations. Over all, the carbonate-rich shale was more reactive than the clay mineral-rich shale as it was documented by the dissolution and re-precipitation of carbonates, dissolution of plagioclase, illite, chlorite and the formation of smectite. Carbon dioxide was also permanently trapped in the same sample as calcite. The clay-rich shale did not show significant mineralogical alterations except for dissolution of silicate minerals. The availability Mg through dissolution of ankerite, among other factors, could be the most likely reason for the formation of smectite in the carbonate-rich shale. In the experiments, the release of SiO2 and Fe into solution appears to be more sensitive to the increase in acidity created by the increase in the dissolution of CO2 in brine at lower temperatures. Analcime did not form in samples reacted with CO2–brine probably due to the high silica activity and low Na+/H+ activity ratio. The results of the geochemical modeling showed reasonable agreement with the experimental results, even though some discrepancy existed. Nonetheless, availability of a larger thermodynamic database would help in increasing the accuracy of such models. If subsurface storage of CO2 is to be practical, it is essential to understand the basic chemical reactions that take place between caprock material, CO2 and the formation water. As evidenced by the experimental results, reactivity of shale is highly dependent on the mineral composition of the rock since individual mineral reactions are interdependent. In addition, fluid chemistry plays a crucial role as it controls the evolution of dissolution–precipitation processes. In the carbonate-rich layers the dissolution of carbonates might lead to increase in porosity only in cases where transport of solutes is significant. The fact that the thin carbonate dominated layers are arranged in alternating layers with clay-rich matrix suggests a limited dissolution progress due to solute movement restrictions. In the shaly-caprock, the potential alteration of

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matrix permeability/porosity due to interaction between CO2–brine and caprock appears to be extremely low. Therefore, in the caprock at the study site, the enhancement of CO2 movement due to fluid–rock interaction appears to be very limited. The results presented are only in reference to intact shale matrix and caution should be practiced while interpreting the results of this study as the presence of other factors such as micro-fractures and sand bodies have the ability to modify the system. Ultimately, these types of chemical study of water–rock interactions should be incorporated into the regional 3D coupled models that take into account all the relevant factors such as transport and geomechanics in order to predict the integrity of the caprock in a geological CO2 storage environment. Acknowledgements This research is part of the Subsurface Storage of CO2 – Risk Assessment, MOnitoring and REmediation (SSC-RAMORE) project. Funding for the project was provided by: The Research Council of Norway (Grant No. 178008), ConocoPhillips, RWE-DEA, Schlumberger, Shell Technology and Statoil. The authors would like to thank Berit Løken Berg, Mufak Naroz and Ruikai Xie, for the analysis of samples, Julien Declercq, Helge Hellevang, Olivier Regnault and Øyvind Brandvoll for their support and helpful discussions during the experiment. We would also like to thank Professor Alvar Braathen and the Lonyearbyen CO2 Lab for providing materials for the experiment, and Amy Dale for correcting the language in an earlier version of the manuscript. Finally we would like to thank Paul Nadeau and an anonymous reviewer for their constructive and detailed reviews that significantly improved the manuscript. References Aagaard, P., Helgeson, H.C., 1982. Thermodynamic and kinetic constraints on reaction-rates among minerals and aqueous-solutions. 1. Theoretical considerations. Am. J. Sci. 282, 237–285. Aagaard, P., Helgeson, H.C., 1983. Activity composition relations among silicates and aqueous-solutions: II. Chemical and thermodynamic consequences of ideal mixing of atoms on homological sites in montmorillonites, illites, and mixedlayer clays. Clays Clay Miner. 31, 207–217. Audigane, P., Gaus, I., Czernichowski-Lauriol, I., Pruess, K., Xu, T.F., 2007. Twodimensional reactive transport modeling of CO2 injection in a saline Aquifer at the Sleipner site, North Sea. Am. J. Sci. 307, 974–1008. Bachu, S., 2000. Sequestration of CO2 in geological media: criteria and approach for site selection in response to climate change. Energy Convers. Manage. 41, 953– 970. Bachu, S., Gunter, W.D., Perkins, E.H., 1994. Aquifer disposal of CO2: hydrodynamic and mineral trapping. Energy Convers. Manage. 35, 269–279. Bjørlykke, K., 1983. Diagenetic reactions in sandstones. In: Parker, A., Shellwood, B.W. (Eds.), Sediment Diagenesis. Reidel Publishing Company, Dordrecht, pp. 169–213. Boles, J.R., 1978. Active ankerite cementation in the subsurface Eocene of southwest Texas. Contrib. Mineral. Petrol. 68, 13–22. Boles, J., Franks, S., 1979. Clay diagenesis in Wilcox sandstones of southwest Texas: implications of smectite diagenesis on sandstone cementation. J. Sediment. Petrol. 49, 55–70. Bowden, A., Rigg, A., 2005. Assessing reservoir performance risk in CO2 storage projects. In: Rubin, E., Keith, D., Gilboy, C. (Eds.), Greenhouse Gas Control Technologies. Elsevier Ltd, pp. 683–691. Brady, P.V., Carroll, S.A., 1994. Direct effects of CO2 and temperature on silicate weathering – possible implications for climate control. Geochim. Cosmochim. Acta 58, 1853–1856. Brandt, F., Bosbach, D., Krawczyk-Barsch, E., Arnold, T., Bernhard, G., 2003. Chlorite dissolution in the acid pH-range: a combined microscopic and macroscopic approach. Geochim. Cosmochim. Acta 67, 1451–1461. Cantucci, B., Montegrossi, G., Vaselli, O., Tassi, F., Quattrocchi, F., Perkins, E.H., 2009. Geochemical modeling of CO2 storage in deep reservoirs: the Weyburn Project (Canada) case study. Chem. Geol. 265, 181–197. Carnicelli, S., Mirabella, A., Cecchini, G., Sanesi, G., 1997. Weathering of chlorite to a low-charge expandable mineral in a spodosol on the Apennine mountains, Italy. Clays Clay Miner. 45, 28–41. Chen, Y., Brantley, S.L., 1997. Temperature- and pH-dependence of albite dissolution rate at acid pH. Chem. Geol. 135, 275–290. Credoz, A., Bildstein, O., Jullien, M., Raynal, J., Pétronin, J., Lillo, M., Pozo, C., Geniaut, G., 2009. Experimental and modeling study of geochemical reactivity between clayey caprocks and CO2 in geological storage conditions. Energy Procedia 1, 3445–3452.

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