Effect of organic and inorganic ligands on calcite and magnesite dissolution rates at 60 °C and 30 atm pCO2

Chemical Geology 265 (2009) 33–43

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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 ev i e r. c o m / l o c a t e / c h e m g e o

Effect of organic and inorganic ligands on calcite and magnesite dissolution rates at 60 °C and 30 atm pCO2 O.S. Pokrovsky a,⁎, S.V. Golubev a, G. Jordan b a b

Géochimie et Biogéochimie Expérimentale, Université de Toulouse; UPS (SVT-OMP), CNRS, LMTG, 14 Avenue Edouard Belin, F-31400 Toulouse, France Dept. für Geo-u. Umweltwissenschaften, Ludwig-Maximilians-Universität München, Theresienstr. 41, 80333 München, Germany

a r t i c l e

i n f o

Article history: Accepted 17 November 2008 Keywords: Calcite Magnesite Dissolution Kinetics Carbon dioxide Phosphate Sulphate Borate Silicate Acetate Oxalate Malonate Succinate Phthalate Citrate EDTA pCO2

a b s t r a c t Calcite and magnesite dissolution rates were measured at 60 °C, 30 atm pCO2, 0.1 M NaCl, and pH from 4.95± 0.05 to 5.60 ± 0.05 as a function of organic (acetate, oxalate, malonate, succinate, phthalate, citrate, EDTA) and inorganic (sulphate, phosphate, borate, silicate) ligand concentration in the range of 10− 5 to 10− 2 M. These conditions can be considered as boundary model environments for sedimentary oil-field basins of underground CO2 storage. Experiments on dissolution of magnesite powders (100–200 µm) and calcite crystal planes were performed in a batch reactor with in-situ pH measurements and under controlled hydrodynamic conditions using the rotating disk technique. At 60 °C in circumneutral solutions in the presence of 0.02 M NaHCO3 and 30 atm pCO2 (pH = 4.95), calcite dissolution is weakly affected by the presence of ligands: the rates increase at the maximum by a factor of 2 and, at 0.01 M ligand concentration in solution, the order is: silicate < citrate < NaCl ∼ borate < malonate < EDTA < sulphate < acetate. The order of ligand effects on calcite dissolution at pH = 5.55 (0.1 M NaHCO3, 30 atm pCO2) is: phosphate < NaCl < citrate < acetate < succinate < malonate < phthalate < EDTA. Magnesite dissolution rates at 60 °C, 30 atm pCO2 and 0.02 M NaHCO3 (pH = 4.95) were weakly affected by the presence of acetate, silicate, borate and NaCl but increase in the presence of sulphate, EDTA, citrate and oxalate. These ligandaffected rates were rationalized using a phenomenological equation which postulates the Langmuirian adsorption of a negatively-charged or neutral ligand on rate-controlling surface sites, presumably > MeOH+ 2 (Me = Ca, Mg). Proposed equations of rate–ligand concentration dependencies can be directly incorporated into reaction transport codes. Results obtained in this study demonstrate that both magnesite and calcite reactivity is not appreciably affected by acetate, oxalate, citrate, succinate, sulphate, and phosphate that are most likely present in deep carbonate aquifers at the physical and chemical conditions pertinent to CO2 geological sequestering sites. The concentration of ligands necessary to increase the rates by a factor of 3 to 10 is on the order of 0.01 M. Such a high concentration is unlikely to be encountered in deep sedimentary basins. Therefore, as a first approximation, reactive transport modelling of dissolution induced by CO2 injection in carbonate rocks does not require to explicitly account for the effect of dissolved organics. © 2008 Elsevier B.V. All rights reserved.

1. Introduction Carbonate mineral reactivity has been an issue of active research efforts in the field of biomineralization (Boquet et al.,1973; Morita,1980; Monger et al., 1991; Pokrovsky and Savenko, 1994; Ferris et al., 1994; Pokrovsky and Savenko, 1995; Fujita et al., 2000; Warren et al., 2001; Ferris et al., 2003; Dittrich and Obst, 2004; Lian et al., 2006; Mitchell and Ferris, 2006; Rodriguez-Navarro et al., 2007), and extremely important in both chemical weathering and CO2 sequestration. Weathering of carbonate rocks at the earth surface consumes 0.29 Gt C/year which is considerably more than the overall consumption due to silicate weathering (0.21 Gt C/year). Although on the million-year scale the balance between atmospheric CO2 uptake on the continents and its ⁎ Corresponding author. Tel.: +33 5 61 33 26 17; fax: +33 5 61 33 25 60. E-mail address: [email protected] (O.S. Pokrovsky). 0009-2541/$ – see front matter © 2008 Elsevier B.V. All rights reserved. doi:10.1016/j.chemgeo.2008.11.011

release due to carbonate precipitation in the ocean is nil, on the shortperiod time the variations of carbonate weathering rate and their dependencies on external factors become important for modelling the carbon cycle. Numerous laboratory experiments demonstrated that the dissolution of main carbonates (calcite, magnesite) is controlled by pH, 2− concentration of HCO− 3 /CO3 ions, and temperature. At the same time, in contrast to the large amount of work devoted to the effect of ligands on simple and multiple oxide dissolution (i.e., Wogelius and Walther, 1991; Stumm, 1992; Welch and Ullman, 1993; Welch and Vandevivere, 1994; Welch and Ullman, 1996; Stillings et al., 1996; Oelkers and Schott, 1998; Oelkers and Gislason, 2001; Van Hees et al., 2002; Pokrovsky et al., 2005b to cite just a few), carbonate minerals received much less attention, and most works on carbonates dealt with precipitation of calcite in the presence of various natural and synthetic organic ligands. Despite a number of studies on calcite dissolution in the presence of organic ligands (Compton and Sanders, 1993; Frye and Thomas, 1993;

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Thomas et al., 1993; Compton and Brown, 1995; Fredd and Fogler, 1998a, b; Hoch et al., 2000; Wu and Grant, 2002; Perry et al., 2004, 2005; Spanos et al., 2006), the effect of variable ligand concentration on the dissolution rate has been rarely rigorously quantified. There are only a few data on magnesite dissolution in the presence of ligands at elevated temperatures and high ligand concentrations (Hamdona et al., 1995; Demir et al., 2003; Laçin et al., 2005; Bayrak et al., 2006). Jordan et al. (2007) studied ligand-controlled magnesite dissolution at 100 °C and low pCO2 in the presence of a single concentration (0.01 M) of organic and inorganic ligands. In order to extend the range of ligand concentrations to broader environmental conditions, in this study, we focused on calcite and magnesite dissolution in the presence of variable ligand concentrations (10− 5 to 10− 2 M) at otherwise constant solution parameters. The second motivation of this work is related to quantitative and predictive modelling of CO2 sequestration in deep aquifers. Such a modelling requires precise knowledge of carbonate mineral reactivity at conditions pertinent to CO2 storage, i.e., elevated temperatures, pCO2 1 atm and the presence of ligands. Whereas the effect of temperature, salinity, pH, and pCO2 on Ca- and Mg-carbonate dissolution has been extensively studied (Morse and Arvidson, 2002; Higgins et al., 2002a,b; Pokrovsky et al., 2005a; Gledhill and Morse, 2006), the understanding of the influence of various organic and inorganic ligands, omnipresent in deep sedimentary basins, remains very limited and there is no single measurement performed on calcite or other carbonate minerals at elevated pCO2 (>1 atm) and temperatures, pertinent to environmental conditions of CO2 storage. The development of surface complexation models (SCM, Van Cappellen et al., 1993; Pokrovsky et al., 1999) allows predictive rationalization of the dissolution kinetics of carbonate minerals within the framework of surface coordination theory. Within this approach, the fully hydrated metal centers (>MeOH+ 2 , Me=Ca, Mg) govern the H2Oand ligand-promoted carbonate mineral dissolution at pH> 5–6 (Pokrovsky and Schott, 2001, 2002). Although this model exhibits high predictive capacity at 25 °C (i.e., Pokrovsky et al., 2005a), its application for description of carbonate mineral reactivity at elevated temperatures requires the knowledge of both activation energies and enthalpy of H+, 2− HCO− 3 , CO3 , and ligand adsorption reactions. In the absence of this information, a straightforward technique for evaluating the governing factors of carbonate dissolution in potential CO2 storage sites is the experimental measurement of rates in the presence of model organic and inorganic ligands. Assessing the magnitude of the ligand concentration effect on rates under fixed solution parameters should allow extrapolation of the results to a wider range of environmental conditions and finally the incorporation of special coefficients for ligand presence into the computer codes for reactive transport modelling. This study is aimed at filling the existing gap in carbonate mineral reactivity by presenting a detailed study of dissolution of two contrasting carbonates, calcite and magnesite, in the presence of 7 organic (acetate, oxalate, malonate, succinate, phthalate, citrate, EDTA) and 4 inorganic (sulphate, phosphate, borate, silicate) ligands under controlled solution parameters. 2. Experimental methods Large transparent crystals of calcite from hydrothermal veins in basaltic traps (Central Siberia) and magnesite of hydrothermal origin (Satka, Ural), as described in Pokrovsky et al. (1999, 2005a), were used in this study. Electron microprobe and total chemical analysis showed that the samples contained less than 0.5% chemical impurities and no other phases were detected using X-ray diffraction. For rotating disk experiments, cores were drilled into the calcite single crystal normal to the (104) plane and sliced into disks of ~ 2–4 mm in diameter and 5 mm in thickness. The slices were fixed in a polycarbonate holder using an epoxy resin to provide ~0.5–1.5 cm2 of exposed area to solution and polished to a smooth surface (Pokrovsky et al., 2005a). Experiments with magnesite were performed with powders of the

size fraction between 100 and 200 µm that was reacted for several seconds in 1% HCl, ultrasonically cleaned in alcohol to remove adhering fine particles, rinsed repeatedly with distilled water, and dried overnight at 60 °C. Its specific surface area was 660 ± 60 cm2 g− 1 as determined by multi-point krypton absorption using the B.E.T. method. In the course of experiment, the BET surface area of magnesite powder did not change by more than 20%. The geometric surface area of this powder was calculated as 134 cm2 g− 1. 34 magnesite and 50 calcite dissolution experiments at 60 °C, 0.1 M NaCl, pCO2 = 30 atm, pH= 4.95 ± 0.05 and 5.60 ± 0.05, and [NaHCO3] = 0.02 and 0.1 M (Appendix A) were carried out in a titanium high pressure batch reactor which was continuously stirred with a magnetically driven stirrer (Parr Instrument Company). The reactor was equipped with pH and reference electrodes. Carbon dioxide was supplied to the reactor through a Ti 2 µm porous filter and pCO2 was controlled by a calibrated pressure meter. Solution was manually sampled each 0.5–1 h using a valve equipped with a Ti 2 µm porous filter and a back-pressure regulator. Duration of experiments varied from 5 to 15 h depending on fluid composition and the type of solid. The total amount of solution sampled (i.e., 12–15 samples of ~3 mL volume each) did not exceed 10% of the initial mass of solution and necessary corrections were made (e.g., Pokrovsky et al., 2005a). In this closed-system reactor, rates (R) were generated from measured solution composition as a function of time using  h 2 R = Δ Me

+

i tot

 = Δt = s

ð1Þ

where t (s) designates the elapsed time, [Me2+]tot (mol) stands for the amount of metal (Ca, Mg) released from the solid, and s is the disk geometric area (for calcite cleavages) or total B.E.T. surface area (for magnesite powder). In most experiments presented in this study, the same disk was used only once. The change of specific surface area of the magnesite powder did not exceed 20–30% and thus the initial B.E.T. surface area was used for the rate calculation. A discussion of B.E.T. versus geometric surface area normalization for carbonates dissolution at high temperature can be found elsewhere (Jordan et al., 2007; Golubev et al., 2009-this issue). Reproducibility of rates from batch dissolution experiments was ≤20%. The uncertainties attached to individual rate measurements in closed reactors were assessed from the uncertainty on the slope of [Me2+]tot–time dependencies (Eq. (1)). Solution pH was measured in the course of experiments via solid contact Li–Sn alloy commercial pH electrode coupled with a home-made Ag/AgCl reference electrode providing a constant potential in Cl−-rich solutions (0.1 M NaCl): Sn–Cu, Li–Sn alloy ∣ H-selective glass ∣ Cl−-bearing test solution ∣AgCl/Ag,  E = E + 2:3RT = F log aH ◯

+

+ log aCl −



ð2Þ

where E is the measured electric potential of the system, E° is a standard potential determined from the calibration curve, aH+ and aCl− are the activities of hydrogen and chloride ions in solution, R is gas constant, T is absolute temperature, and F is Faraday constant. The electrode system has an E° of ≤−2200 mV and was connected to a high input impedance, high resolution pH meter (Econix-Expert(R), Russia). The electrode was calibrated in mixture solutions of HCl (0.0003–0.1 M) and NaCl (0 to 0.1 M) having a constant ionic strength of 0.1 M. In basic solutions, a mixture of 0.1 M NaCl, 0.002–0.01 M NaHCO3, and 0.1 M NaCl with 0.01 M Na2B4O7 was used. The pH of the calibration solutions was calculated using the MINTEQA2 computer code with implemented database (Allison et al., 1991). Fast Nernstian response with a slope of 65.13 mV/pH unit versus the theoretical value of 66.1 mV/pH unit was observed and the electrode potentials exhibited sufficient stability to provide the uncertainty of ±0.01 pH units at 60 °C. Between two subsequent calibrations, the electrode

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to 0.1 M NaCl medium (Fig. 3B and D) pointing towards a high pit nucleation rate which is in accordance with the observed high dissolution rate in the presence of this ligand (see below). The surface roughness was greater for calcite than magnesite under similar reaction conditions. Indeed, the roughness of the calcite disk was increased in both NaCl and ligand-bearing solutions (citrate, EDTA, acetate). As shown in Fig. 3G, EDTA produced large and very deep etch pits and acetate and citrate yielded a high amount of triangular pits (Fig. 3E, F). In contrast to magnesite, the calcite disk experiment with acetate (Fig. 3E) shows a surface with a far more advanced morphological transformation. The morphology is already Fig. 1. Example of a solid-state pH electrode calibration coupled with a Ag/AgCl selective electrode.

shift was less than 1 mV/day over 3 days at 60 °C. An example of calibration at 60 °C is given in Fig. 1 and further details are provided elsewhere (Pokrovsky et al., 2005a, 2009-this issue). All solutions were analyzed for magnesium ([Mg2+]t), calcium ([Ca2+]t), alkalinity ([Alk]), and pH as a function of time. Magnesium and calcium were measured by flame atomic absorption with an uncertainty of ± 1% and a detection limit of 4 × 10− 8 and 7 × 10− 8 M, respectively. Alkalinity was determined following a standard HCl titration procedure with an uncertainty of ± 1% and a detection limit of 5 × 10− 5 M. Microscopic analysis of fresh and reacted surfaces were performed using a Jeol JSM840a scanning electron microscope (SEM) after carbon film deposition on the surface. Homogeneous solution equilibrium, as well as carbonate mineral surface speciation, and chemical affinities of their dissociation reactions were calculated for each solution at 60 °C using the MINTEQA2 code (Allison et al., 1991). Thermodynamic constants for aqueous and solid reactions equilibria were taken from Plummer and Busenberg (1982). Throughout the text we always used “M” for molar concentration scale (mol/l). 3. Results Examples of the amount of dissolved metals released as a function of time for different ligands are presented in Fig. 2 for calcite and magnesite. Over the investigated reaction time (20–500 min) a constant release rate of Ca and Mg is observed. This indicates that the contribution of reaction products (backward reaction) to the overall dissolution rate is insignificant. In other words, the rates are unaffected by the Ca and Mg concentrations present in the solution at constant pCO2. At our experimental conditions (4.95 ≤ pH ≤ 5.65 for pCO2 = 30 atm), all solutions were undersaturated with respect to carbonate solid phases (i.e., Ω < 0.2, where Ω is the ratio of ion activity product in solution to ion activity product at equilibrium with the solid phase, K°sp) and more than 90% of dissolved metals (Ca, Mg) were present in the form of free ions. SEM observations revealed that at circumneutral conditions without ligands, magnesite typically shows rhombic shaped etch pits parallel to [481] and [441] with straight edges (Jordan et al., 2001, 2007). However, in the presence of EDTA (Fig. 3A) and citrate (not shown), microscopic investigations of the reacted magnesite surfaces revealed etch pits with rounded edges which were rotated inward from rhombic directions. According to the convenient nomenclature of symmetrically inequivalent step directions on the (104) surface of calcite isotype minerals (Jordan et al., 2001, 2007), these rounded and rotated pit edges correspond to the acute monolayer steps. AFM investigations of the magnesite (104) surface at low pCO2 but comparable pH and temperature revealed the morphology of the obtuse steps to be affected in the presence of these ligands (Jordan et al., 2007) and, thus, suggest a morphologically active role of pCO2. In the presence of 0.01 M oxalate, a higher density of etch pits was observed on magnesite compared

Fig. 2. Amount of dissolved calcium (A, B) or magnesium (C) as a function of elapsed time for different ligand concentrations during calcite (A, B), and magnesite (C) dissolution in batch reactor. Experimental conditions: 60 °C, 30 atm pCO2, 0.1 M NaCl + (0.02–0.1) M NaHCO3 for magnesite and calcite, respectively.

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Fig. 3. Scanning Electron Microscopy (SEM) images of reacted carbonate surfaces (0.1 M NaCl, 60 °C, 30 atm pCO2, pH = 4.95–5.5). Acceleration voltage is 10 kV. (A) Magnesite dissolved in the presence of 0.001 M EDTA, 0.01 M Oxalate (B, C) and in 0.1 M NaCl (D). Calcite in the presence of 0.1 M Acetate (E), 0.01 M citrate (F), and 0.01 M EDTA (G).

transformed from an etch pit morphology into an etch hillock morphology (which can be perceived as the remnants of old pit walls which have coincided with other pits). At this stage of morphological development it is very difficult to evaluate the specific

relations between reactivity and direction without a considerably higher spatial resolution of the surface images. Literature information on the influence of ligands on the etch pit morphology of calcite is very limited. For instance, Britt and Hlady

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37

Mg) species concentration due to formation of non-reactive >MeCO3°, >MeHCO3° species (Pokrovsky and Schott, 1999, 2002; Pokrovsky et al., 2005a). Magnesite dissolution rates at 60 °C, 30 atm pCO2 and 0.02 M NaHCO3 (pH = 4.95) were weakly affected by the presence of acetate, silicate, borate and NaCl but increased in the presence of sulphate (Fig. 7A, B), oxalate (Fig. 7C) and citrate (Fig. 7D). The order of magnesite dissolution rate increase in the presence of 0.01 M total ligand concentration was the following: acetate ≅ silicate ≅ borate ≅ NaCl (no ligand) < sulphate < EDTA < oxalate (Fig. 8). EDTA had the strongest effect on magnesite dissolution as the rates increased at [H2EDTA2−] > 10 µM both at pH= 5.55 and 4.95 (Fig. 9A and B, respectively). 4. Discussion: Modeling of ligand-affected carbonate mineral dissolution The effect of ligands on calcite and magnesite dissolution can be modeled within the framework of the surface coordination approach assuming that the overall dissolution rate is controlled by reactions promoted at Ca and Mg centers by various ligands which compete for available surface sites (Stumm, 1992). The effectiveness of ligands depends on the molecular structure and thermodynamic stability of the surface complexes they form. For example, especially efficient are ligands whose functional groups contain two or more oxygen donors and which can form bi- or multidentate mononuclear surface chelates (Stumm, 1992). In contrast, ligands forming bi- or polynuclear complexes, that can bridge two or more metal centers at the surface lattice, are known to retard dissolution. In order to quantitatively model the effect of ligand on carbonate dissolution rate, the simplest, bimolecular surface reaction between positively charged surface groups (>MeOH+ 2 , Me = Ca, Mg, as for other Ca, Mg-bearing multiple oxides) and negatively charged ligands (Ln−) can be considered: Fig. 4. Calcite dissolution rate as a function of inorganic ligand concentration in 0.1 M NaCl, 0.02 M NaHCO3, 30 atm pCO2 and pH = 4.95–5.55 in the presence of borate, sulfate and silicate (A) and phosphate (B). The solid line represents the fit to the data using Eq. (8).

(1997) reported rounded obtuse steps in a saturated CaCO3 (calcite) solution containing 1 mM EDTA. Keith and Gilman (1960) reported clear deviations from [481] and [441] in ~10% acetic acid, saturated citric acid solution, and saturated EDTA solution. An overview on the influence of maleic and fumaric acids is presented by Compton et al. (1989). The presence of inorganic ligands (sulphate, borate and silicate) in circumneutral solutions with 0.02 M NaHCO3 and 30 atm pCO2 (pH=4.95) did not produce significant variation of calcite dissolution rate: no systematic trend of rates as a function of ligand concentration in solution was observed (Fig. 4A). At these conditions, phosphate acts as an inhibitor (Fig. 4B) and calcite dissolution rates weakly depend on the presence of organic ligands such as acetate, malonate, and citrate (Fig. 5A, B and C, respectively). The EDTA increased the rates by a factor of 1.5–2 when its concentration increased from 10− 6 M to 0.01 M (Fig. 5D). The effect of ligands on calcite dissolution became better pronounced in more basic solutions as it is seen in Fig. 5 with citrate and EDTA being the strongest enhancers of dissolution. We can establish an order of calcite rate increase in the presence of 0.01 M total ligand concentration: silicate < citrate < NaCl (no ligand) ≅ borate < malonate < EDTA < sulphate < acetate at pH 4.95 and phosphate < NaCl (no ligand) < citrate < acetate < succinate < malonate < phthalate < EDTA at pH = 5.55 (Fig. 6A and B, respectively). Note that, at far from equilibrium, rates at pH=5.55 (0.1 M NaHCO3) are six times lower than those at pH=4.95 (0.02 M NaHCO3). This is consistent with calcite rate decrease in the presence of carbonate at pH 9 (Lea et al., 2001) and can be explained by the 2− inhibiting effect of HCO− 3 and CO3 ions on metal carbonate dissolution. This effect consists in decreasing the rate controlling >MeOH+ 2 (Me=Ca,

þ

>MeOH2 + L

n−

= > Me −L

1−n

+ H2 O

ð3Þ

According to this scheme, the rate of ligand-controlled dissolution is proportional to the concentration of the surface center–ligand complex >Me–L1 − n which can be deduced from reaction (3) stability 4 ): constant (KMe−L ( 4

KMe−L =

>Me−L1 − n g   >MeOHþ × ½L n −  2

ð4Þ

The asterisk ⁎ means that this conventional reaction is written in terms of surface {>i} and aqueous [Ln−] species concentrations, not activities. We postulate that in the presence of a ligand, the calcite or magnesite forward dissolution rate is thus the sum of H2O (RH2O) and ligand-controlled dissolution, similar to that of brucite (Pokrovsky et al., 2005b), dolomite (Pokrovsky and Schott, 2001), smectite (Golubev et al., 2006), and diopside (Golubev and Pokrovsky, 2006). The assumption that the rate is proportional to the concentration of the surface metal–ligand complex is notoriously known for simple metal oxides (Stumm, 1992) and will be used in this study as the working hypothesis for metal carbonates. Therefore, the forward dissolution rate is given by: n 1−n g R = RH2 O + k>Me − L f × >Me−L

ð5aÞ

and, because surface metal detachment (rate in the absence of ligand) is proportional to the concentration of hydrated metal (Pokrovsky et al., 1999, 2005b), n o f þ RH2 O = k>MeOHþ × >MeOH2 2

ð5bÞ

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Fig. 5. Calcite dissolution rate as a function of phosphate (A), acetate (B), malonate (C), citrate (D) and EDTA (E) in 0.1 M NaCl, 30 atm pCO2, pH = 4.95 (0.02 M NaHCO3) and pH = 5.63 (0.1 M NaHCO3). The solid line represents the fit to the data using Eq. (8). The error bars for experiment at pH = 5.5 with acetate (A) are within the size of the symbols.

ð7aÞ

experiment, the concentration of >MeHCO3° and >MeCO− 3 surface sites is constant and independent on ligand concentration since there is weak or no interaction of neutral or negatively charged surface groups with negatively-charged ligand. The surface stability constants for reactions (7a) and (7b) are available only for 25 °C (Pokrovsky et al., 1999). Using these constants, we estimated that, at pH=4.95 (0.02 M NaHCO3) and pH=5.55 (0.1 M NaHCO3), 53 and 77%, respectively, of metal sites are occupied by >MeHCO3° species and that the contribution of >MeCO− 3 species is insignificant. Since the enthalpy of surface reactions does not significantly change between 25 and 60 °C (Pokrovsky et al., 2009-this issue) similar speciation is expected at 60 °C. Therefore, we will consider 2− that, although the competition between Ln− and HCO− 3 /CO3 ions for > + MeOH2 surface sites does take place, the contribution of >MeCO− 3 and >MeHCO3° remains constant over the entire range of ligand concentration. This simplification renders our approach fully phenomenological in the sense that parameters of equations are valid only for given pH, pCO2 and temperature. Indeed, even if we consider the rigorous derivation of {>Me–Lx − n} from the SCM taking into account 2− electrostatic contribution and ligand–HCO− 3 /CO3 competition for the surface sites, the limited number and scattering in rate datapoints do not allow more precise fitting. The simplified phenomenological equation for calcite and magnesite dissolution in the presence of ligands can be obtained by combining Eqs. (4)–(6):

ð7bÞ

R = k>MeOHþ

where {>MeOH+ 2 } designates concentration of the rate-controlling + surface sites, non-occupied by the ligand (>MgOH+ 2 , >CaOH2 ), Eq. (5a) can be rearranged to: n o n f þ 1−n g ð5cÞ R = k>MeOHþ × >MeOH2 + k>Me − Lf × >Me−L 2

In Eqs. (5a)–(c), kf>MeOHþ is the rate constant for water-promoted 2 dissolution reaction measured in the absence of ligand at given pH and pCO2 and k>Me − Lf is the empirical kinetic constant pertinent to each ligand. We further postulate that, at the constant pH of our experiments, ligand sorption on carbonate mineral surface follows a Langmuirian adsorption isotherm as it was experimentally observed for the case of ligand adsorption on >MgOH+ 2 sites of brucite (Pokrovsky et al., 2005b). The mass law conservation for the metal surface sites is n o  þ x−n fMeTOTAL g = >MeOH2 + >Me−L g

ð6Þ

where {>MeTOTAL} is the total rate-controlling site number on carbonate mineral surface. This is a simplified representation of carbonate mineral surface speciation, since, at such high pCO2 used in our experiments, adsorption of carbonate and bicarbonate ions on metal surface sites: þ

−

0

>MeOH2 + HCO3 = > MeHCO3 + H2 O þ

−

−

>MeOH2 + CO3 = > MeHCO3 + H2 O

leads to a decrease of {> MeOH+ 2 }. However, we can assume that, at constant pH and pCO2 (and thus, [HCO− 3 ]) of each individual

f

2

·

1−

 !  KMe − L4 · Ln − KMe − L4 · Ln − + k>Me − Lf · n− n− 1 + KMe − L4 · ½L  1 + KMe − L4 · ½L  f

ð8Þ

where the first term represents RH2O, and k>MeOHþ = k>MeOHþ · 2 2 f>MeTOTAL g:

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39

Fig. 6. Calcite dissolution rates in the presence of various ligands (0.01 M total concentration) at pH 4.95 (A) and 5.55 (B).

Application of Eq. (8) to model the experimental dependence of calcite and magnesite dissolution on ligand concentration requires accurate values of reaction (3) stability constant, KMe − L4 , for the ligands investigated in this study. Such values were generated from the fitting of dissolution rate dependences on aqueous ligand concentrations. The fitting procedure performed by “trial and error” consisted in assuming that the values of the equilibrium constants for ligands adsorption are the same as the corresponding values for Me–ligand association reactions in aqueous solution (Schindler and Stumm, 1987). This assumption was earlier tested for brucite dissolution, for which ligand adsorption data are available, and was found to yield reasonable estimation of ligand adsorption constant (Pokrovsky et al., 2005b). The values of k>Me − Lf f were set first equal to k>MeOHþ assuming no inhibition or catalysis to 2 occur. Then, the values of rate constants k>Me − Lf were allowed to vary in order to fit the experimental dependence of dissolution rate on ligand concentration using Eq. (8). The values of kf>MeOHþ at a given pH were 2 determined in ligand-free experiments. When fitting was not possible with the pre-selected value of K>Me − L4 , the ligand adsorption constant was allowed to vary within one order of magnitude from the initial settings. A ≤ 20% disagreement between measured and modeled (Eq. (8)) rate values at each ligand concentration was taken as a criterion of goodness of fit. The final values of constants used in Eq. (8) are listed in Table 1. The uncertainties attached to these values correspond to the range of best fits obtained by varying the k>Me − Lf and KMe − L4 . The degree to which Eq. (8) can be used to describe the effect of the investigated ligands on carbonate dissolution rates can be assessed in Figs. 4, 5, 7 and 9. The solid curves depicted in these figures were

f computed with Eq. (8) using values of KMe − L4 , kMe − Lf and K MeOHþ 2 listed in Table 1. A reasonable correspondence between the solid curves and experimental data for a broad range of aqueous concentrations of ligands demonstrates the validity of Eq. (8). Note, however, that the effect of phosphate, acetate and citrate was poorly fitted for some range of ligand concentration (Figs. 4B, 5A, C). It is worth noting that the values of surface adsorption constants for acetate, oxalate, citrate, sulfate generated by fitting the rate data are consistent with those recommended earlier for ligand-promoted dolomite, brucite, wollastonite and smectite dissolution which is also controlled by the hydrolysis of alkali-earth metal surface centers >MeOH+ 2 (Pokrovsky and Schott, 2001; Pokrovsky et al., 2005b; Golubev et al., 2006). Given the small number of experimental rate datapoints for each ligand and their certain dispersity, these constants should be considered as the best preliminary estimates subjected to large uncertainties. The uncertainty estimation for kinetic constants is difficult because k and K (Eq. (8)) covary. Therefore, since our fitting procedure did not follow a least-squares minimization algorithm, the co-variance matrix resulting from a least-squares analysis provides uncertainty estimates (in average, ±20% on both kf>MeOHþ and k>Me − Lf ). 2 The sequence of rate constants listed in Table 1 and the comparison f of the ratio kMe − Lf = k>MeOHþ show that carboxylic acids like acetate, 2 that are known to form monodendate surface complexes on oxides, promote dissolution to a much less extent than those forming surface chelates (oxalate, citrate), especially five-membered chelate rings with surface Ca and Mg ions (EDTA). For example, the strong effect of hydroxyl-bearing ligands such as citrate can be understood in view of the marked affinity of their hydroxyl groups for surface Mg and Ca as it is

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Fig. 7. Magnesite dissolution rate as a function of sulfate, borate and silicate (A), acetate (B), oxalate (C), and citrate (D) concentration in 0.1 M NaCl, 0.02 M NaHCO3, and pH = 4.95. The solid line represents the fit to the data using Eq. (8).

the case for calcite (Geffroy et al.,1999), dolomite (Pokrovsky and Schott, 2001) and brucite (Pokrovsky et al., 2005b). Our previous study on magnesite at 100 °C and low pCO2 pressure demonstrated that, among various organic and inorganic ligands, only EDTA and citrate caused a significant increase of the dissolution rate (Jordan et al., 2007). The inhibiting effect of phosphate observed in this study at circumneutral pH (Figs. 4B and 6) is in accord with existing data on phosphate inhibition of calcite dissolution in seawater (Morse and Arvidson, 2002). It follows from the fitting procedure employed here that the stability constants for surface adsorption reactions correlate with

corresponding values for association reactions in homogeneous aqueous solution as it is the case for other simple oxides (Schindler and Stumm, 1987; Ludwig et al., 1995; Pokrovsky et al., 2005b). Such a correlation is depicted in Fig. 10. High uncertainty for both surface and aqueous complex stability constants when log KMe − L4 <2 explains some of the scattering in this correlation. The effect of ligands on calcite dissolution rate is generally lower than that on magnesite dissolution. Indeed, the ratio between ligand-free and f ligand-affected dissolution rate constant (k>Me − Lf = k>MeOH + ), is lower 2 for calcite compared to magnesite (except for EDTA, Table 1). Our previous

Fig. 8. Magnesite dissolution rates at in the presence of various ligands (0.01 M total concentration) at pH 4.95.

O.S. Pokrovsky et al. / Chemical Geology 265 (2009) 33–43

41

Fig. 10. Plot of stability constant at 25 °C, I = 0.1 M for > MeOH+ 2 –ligand surface complexes vs. stability constants of Me2+–ligand aqueous complexes (Table 1).

Fig. 9. Magnesite dissolution rate as a function of EDTA concentration at pH=4.95 (0.02 M NaHCO3, A) and 5.55 (0.1 M NaHCO3, B). The solid line represents the fit to the data using Eq. (8).

studies of ligand effects on dissolution of Ca and Mg bearing minerals (dolomite, brucite, wollastonite, diopside) demonstrated no significant differences in the effect of ligands on Ca or Mg surface sites in terms of Table 1 Parameters of Eq. (8) for magnesite and calcite dissolution in the presence of ligands at pH = 4.95 ± 0.05 unless indicated. Ligand

kf>MeOH

+

2

kf>MeL

f

4 KMeL

kf>MeL = k>MeOHþ 2

5. Conclusions and geological applications

Calcite Sulfate Borate Silicate Acetate (pH = 4.95) Acetate (pH = 5.55) Phosphate (pH=5.63) Malonate Citrate Citrate (pH = 5.63) H2EDTA2− (pH=4.95) H2EDTA2− (pH=5.63)

3.0·10 4.5·10−8 4.0·10−8 3.4·10−8 7.0·10−9 4.0·10−9 3.5·10−8 3.0·10−8 5.0·10−9 3.0·10−8 3.0·10−9

5.0·10 3.0·10−8 1.0·10−8 5.0·10−8 8.0·10−9 1.0·10−9 3.5·10−8 (2±1)·10−8 3.0 · 10− 8 5.0 · 10− 8 5.5 · 10− 8

80 ± 20 100 ± 20 100 ± 20 100 ± 20 N.D. 500 N.D. 300 300 1000 400

1.67 0.67 0.25 1.47 1.14 0.25 1 0.666 6 1.67 1.83

Magnesite Sulfate Borate Silicate Acetate Oxalate Citrate H2EDTA2− (pH=4.95) H2EDTA2− (pH=5.55)

2.0·10−12 2.2·10−12 2.5·10−12 2.2·10−12 2.5·10−12 3.2·10−12 2.2·10−12 2.8 · 10− 12

8.7 · 10− 12 2.0 · 10− 12 8.7 · 10− 13 1.5 · 10− 12 (4.2±0.2)·10−12 1.8 · 10− 11 1.15 · 10− 11 6.5 · 10− 12

50 ± 10 80 ± 20 70 ± 10 100 ± 20 200 ± 10 3000 ± 100 60,000 ± 10,000 40,000 ± 5000

4.35 0.91 0.348 0.68 1.68 5.62 5.22 2.3

−8

−8

Since our fitting procedure did not follow a least-squares minimization algorithm, the covariance matrix resulting from a least-squares analysis provides uncertainty estimates f (±20% on both k>MeOHþ and k>Me − Lf unless indicated). N.D. means non determined. 2

kinetic constants or thermodynamic adsorption constants. Therefore, since there is no specific reason for Mg and Ca surface centers on calcite and magnesite to interact with ligands in a different way, the only plausible explanation for higher effect of ligands on magnesite compare to calcite is the effect of transport process on calcite dissolution. Indeed, it is notoriously known that the chemical reaction on the surface is not a limiting step for calcite dissolution at circumneutral pH, and diffusive transport to the surface provides significant contribution to the overall dissolution process (Sjöberg and Rickard, 1983, 1984). Although the effect of hydrodynamics is also reported for magnesite at 60 °C and 2< pH< 5 (Higgins et al., 2002b), it is certainly much greater for calcite exhibiting c.a. 10,000 times higher overall dissolution rate. As such, the effect of specific ligand–surface interaction for calcite may be partially masked by the transport of both ligand and reaction products within the Nernstian layer. Since the diffusion coefficients of various ligands (carboxylates, oxyanions) are of the same order of magnitude, the difference between different ligands is smaller at lower pH and the effect of strong surface (and aqueous) complexing ligands such as citrate or EDTA is weakly pronounced (compare Fig. 6a and b). The weaker effect of citrate compared to acetate on calcite dissolution (Fig. 6) is due to weaker dissociation of citric acid compared to acetic acid, and probably some steric hindrance on the surface at these conditions. Alternatively, the diffusion coefficient of the large citrate molecules in the Nernstian layer may be lower than that of the smaller acetate.

The present work represents the first assessment of the effect of organic and inorganic ligands on dissolution of rock-forming carbonate minerals at conditions pertinent to CO2 storage (30 atm pCO2 and 60 °C). In circumneutral solutions, the effect of the investigated ligands is generally weak and does not bring more variation in the dissolution rates than a factor of about 2–3. Only MgCO3 in the presence of EDTA, citrate and oxalate shows a significant increase of the dissolution rate (i.e., more than a factor of two in the presence of 0.01 M ligand compared to ligand-free system). Since the dissolution rates of calcite and magnesite differ by a factor of 10,000, considering both minerals allowed better accounting for the effect of diffusion versus surface chemical reaction. Variations of pH (between 4.95 and 5.60) helped to characterize the effect of ligand and surface speciation on mineral reactivity. Although one would expect the stronger effects of ligands in neutral to alkaline pH region, where the ligands are fully deprotonated and metal hydrated centers are the rate controlling species, the presence of high pCO2 and HCO− 3 ion concentration buffers the pH at the circumneutral range and leads to a competition with ligands for the rate-controlling > MeOH+ 2 sites.

42

O.S. Pokrovsky et al. / Chemical Geology 265 (2009) 33–43

It follows from results of reactive transport modelling on carbonate rocks interacting with CO2-rich fluid that, upon fast dissolution of carbonate minerals at an initial pH of 3–3.5, pH will increase to 5.0–5.5 and will essentially be buffered by released HCO− 3 ions (see Brosse, 2002). It has been reported that the presence of inorganic (sulphate, borate, silicate, phosphate) or organic (acetate, oxalate, citrate) ligands in saline aquifers at concentrations ≤0.01 M might exert important kinetic control both for calcite (e.g., Berner and Morse, 1974; Morse and Arvidson, 2002) and dolomite (Brady et al., 1996; Pokrovsky and Schott, 2001) reactivity. The results of the present study suggest the negligible effect (maximum a factor of 2 to 3) of most common organic (acetate and oxalate) and inorganic (sulphate and phosphate) ligands on Ca(Mg)CO3 dissolution both in CO2-rich and CO2-poor fluids. Only in the presence of 0.01 M citrate or EDTA, the rates increase by a factor of 3 to 10. However, such high concentrations are very unlikely to be encountered in natural sedimentary settings. Therefore, pH and the concentration of HCO− 3 ions are the primary governing parameters of carbonate reactivity in sites of CO2 sequestration, and, within a first-order approximation, computer simulation of reaction sequences during CO2 injection in sedimentary rocks does not require explicit provision for ligand effect on dissolution rate. Acknowledgments We thank the two anonymous reviewers for their valuable comments that allowed great improvement of our manuscript. This work was supported by the Research and Development contract DGEP/TDO/CA/ACOMS CT No. FR00000214, ANR, and by the exchange program PROCOPE between France (Ministry of International Affairs, PAI) and Germany (Deutscher Akademischer Austauschdienst, DAAD). The authors are grateful to J.-C. Harrichoury, A. Castillo and Th. Aigouy for careful technical assistance in the course of this study. Appendix A. The list of experiments performed in this study Exp no

Ligand concentration, M

2

Rate, mol/cm /s

Magnesite pCO2 = 30 atm, 60 TM-18–EDTA TM-54–EDTA TM-53–EDTA TM-31–EDTA TM-32–EDTA TM-33–EDTA pCO2 = 30 atm, 60 TM-52–EDTA TM-72–EDTA TM-73–EDTA TM-74–EDTA TM-75–EDTA pCO2 = 30 atm, 60 TM-18–Citrate TM-55–Citrate TM-56–Citrate TM-35–Citrate TM-57–Citrate TM-36–Citrate TM-58–Citrate TM-62–Oxalate TM-63–Oxalate TM-64–Oxalate TM-65–Oxalate pCO2 = 30 atm, 60 TM-36–Acetate TM-37–Acetate TM-38–Acetate TM-39–Acetate TM-40b–Sulfate TM-40–Sulfate TM-41–Sulfate TM-42–Sulfate

°C, 425 rpm; 0.1 M NaCl + 0.02 1.00E−07 1.00E− 06 1.00E− 05 1.00E−04 0.001 0.01 °C, 425 rpm; 0.1 M NaCl + 0.08 1.00E− 07 1.00E−05 1.00E−04 1.00E− 03 0.01 °C, 425 rpm; 0.1 M NaCl + 0.02 1.00E− 06 1.00E− 05 0.0001 0.001 0.001 0.01 0.01 1.00E− 06 1.00E−04 1.00E−03 1.00E−02 °C, 425 rpm; 0.1 M NaCl + 0.02 1.00E−06 0.0001 0.001 0.01 1.00E− 05 0.01 0.03 0.1

M NaHCO3; pH = 4.95 +/− 0.05 1.96E−12 3.53E− 12 5.15E− 12 1.09E− 11 1.06E− 11 1.13E−11 M NaHCO3; pH = 5.55 +/− 0.05 8.50E− 13 3.42E− 12 4.62E− 12 6.25E− 12 7.38E− 12 M NaHCO3; pH = 4.95 +/− 0.05 1.96E−12 5.03E− 12 7.00E− 12 7.55E−12 7.54E− 12 1.35E−11 1.35E− 11 2.50E− 12 3.23E− 12 9.21E− 12 2.61E−11 M NaHCO3; pH = 4.95 +/− 0.05 2.20E− 12 2.37E− 12 1.97E−12 1.84E− 12 1.96E− 12 3.87E− 12 5.46E−12 8.21E− 12

AppendixAA(continued) (continued) Appedix Exp no

Ligand concentration, M

Rate, mol/cm2/s

Magnesite pCO2 = 30 atm, 60 °C, 425 rpm; 0.1 M NaCl + 0.02 M NaHCO3; pH = 4.95 +/− 0.05 TM-43b–Borate 1.00E− 05 1.96E− 12 TM-43–Borate 0.001 3.23E− 12 TM-44–Borate 0.05 1.95E− 12 TM-59–Silicate 0.00001 2.64E− 12 TM-60–Silicate 0.0001 2.91E− 12 TM-69–Silicate 0.001 2.41E− 12 TM-68–Silicate 0.01 1.84E− 12 Exp No

Ligand concentration, M

pH

Rate, mol/cm2/s

Calcite pCO2 = 30 atm, 60 °C, 425 rpm; 0.1 M NaCl + 0.02 M NaHCO3; pH = 4.95 +/− 0.01 EDTA TC-22, 20 TC-90 TC-91 TC-94 TC-92 TC-93 TC-95 CITRATE TC-96 TC-97 TC-98 TC-99 Malonate TC-100 TC-101 TC-102 TC-103 Acetate TC-104 TC-105 TC-106 TC-107 Sulfate TC-108 TC-109 TC-110 TC-111 Borate TC-112 TC-113 TC-114 Silicic acid TC-115 TC-116 TC-117

1.00E− 06 1.00E− 05 1.00E− 04 1.00E− 04 1.00E− 03 0.01 1.00E− 02

4.81 4.81 4.93 4.93 4.92 4.92 4.95

2.66E− 08 3.58E− 08 4.61E− 08 3.95E− 08 3.41E− 08 4.26E− 08 4.77E− 08

1.00E− 05 1.00E− 04 1.00E− 03 1.00E− 02

4.85 4.85 4.85 4.95

3.13E− 08 2.87E− 08 3.33E−08 2.57E− 08

1.00E− 05 1.00E− 04 1.00E− 03 1.00E−02

4.85 4.85 4.80 4.90

3.36E− 08 3.47E−08 3.89E− 08 3.51E− 08

1.00E− 04 1.00E− 03 1.00E− 02 0.1

4.85 4.90 4.80 4.80

1.00E−03 1.00E− 02 0.03 0.1

4.9 4.85 4.95 4.85

3.35E− 08 4.18E− 08 5.41E− 08 3.34E− 08 2.70E− 08 3.97E− 08 4.85E− 08 3.89E− 08 4.42E− 08

1.00E− 03 0.01 0.05

4.92 4.9 4.95

4.67E− 08 3.17E− 08 3.12E− 08

1.00E− 04 1.00E− 03 1.00E− 02

4.98 4.83 4.85

3.70E− 08 3.96E− 08 2.33E− 08

0.0001 0.001 0.01 0.05

5,65 5,62 5,61 5,63

4.09E− 09 5.77E− 09 2.56E− 08 5.53E− 08

1.00E−04 0.001 0.01 0.05

5,63 5,65 5,61 5.62

4.90E− 09 4.70E− 09 6.16E− 09 3.45E− 08

1.00E− 05 0.001 0.01 0.1

5.55 5.50 5.50 5.51

6.27E− 09 7.70E− 09 7.44E−09 7.26E− 09

1.00E− 05 1.00E− 04 1.00E− 03 1.00E− 02 3.00E−02 no ligand 0.01 M KHPhtalate 0.01 M Maleic 0.01 M Succinic

5,63 5,59 5,63 5,65 5,60 5.55 5.50 5.50 5.50

3.82E− 09 4.06E− 09 2.77E− 09 7.72E− 10 2.44E− 09 6.27E− 09 2.03E− 08 1.09E− 08 9.70E− 09

0.1 M NaCl + 0.1 M NaHCO3; 60 °C, 30 atm pCO2 EDTA C-8 C-9 C-10b C-10c Citrate C-11 C-12 C-13 tc-123 Acetate tc-122 tc-124 tc-125 tc-126 Phosphate C-17 C-18 C-19 C-20 C-20b tc-122 tc-119 tc-120 tc-121

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