- Research article
- Open Access
On the potential for CO2mineral storage in continental flood basalts – PHREEQC batch- and 1D diffusion–reaction simulations
© Pham et al.; licensee BioMed Central Ltd. 2012
- Received: 1 June 2011
- Accepted: 14 June 2012
- Published: 14 June 2012
Continental flood basalts (CFB) are considered as potential CO2 storage sites because of their high reactivity and abundant divalent metal ions that can potentially trap carbon for geological timescales. Moreover, laterally extensive CFB are found in many place in the world within reasonable distances from major CO2 point emission sources.
Based on the mineral and glass composition of the Columbia River Basalt (CRB) we estimated the potential of CFB to store CO2 in secondary carbonates. We simulated the system using kinetic dependent dissolution of primary basalt-minerals (pyroxene, feldspar and glass) and the local equilibrium assumption for secondary phases (weathering products). The simulations were divided into closed-system batch simulations at a constant CO2 pressure of 100 bar with sensitivity studies of temperature and reactive surface area, an evaluation of the reactivity of H2O in scCO2, and finally 1D reactive diffusion simulations giving reactivity at CO2 pressures varying from 0 to 100 bar.
Although the uncertainty in reactive surface area and corresponding reaction rates are large, we have estimated the potential for CO2 mineral storage and identified factors that control the maximum extent of carbonation. The simulations showed that formation of carbonates from basalt at 40 C may be limited to the formation of siderite and possibly FeMg carbonates. Calcium was largely consumed by zeolite and oxide instead of forming carbonates. At higher temperatures (60 – 100 C), magnesite is suggested to form together with siderite and ankerite. The maximum potential of CO2 stored as solid carbonates, if CO2 is supplied to the reactions unlimited, is shown to depend on the availability of pore space as the hydration and carbonation reactions increase the solid volume and clog the pore space. For systems such as in the scCO2 phase with limited amount of water, the total carbonation potential is limited by the amount of water present for hydration of basalt.
- Basaltic Glass
- Reactive Surface Area
Underground sequestration of carbon dioxide is a potentially viable greenhouse gas mitigation option as it reduces the release rate of CO2 to the atmosphere . CO2 can be trapped subsurface by four storage mechanisms: (1) structural and stratigraphic trapping; (2) residual CO2 trapping; (3) solubility trapping; and (4) mineral trapping . Mineral trapping has been considered as the safest mechanism in long-term storage of CO2.
Mineral storage of CO2 in basaltic rocks is favored over siliciclastic reservoirs both by the higher abundance of divalent metal ions in basalt and the faster reactivity of basaltic glass or crystalline basalt . Moreover, basalts such as the Columbia River flood basalts (CRB) are abundant and in many places close to CO2 point source emissions . During the last decade several flood basalts around the world have been mapped for the possibility of CO2 storage, and possible candidates such as CRB in USA and the Deccan traps in India have been identified [4–6].
To be a candidate for CO2 storage, the flood basalt must have a proper sealing and sufficient injectivity, the latter limited by the available connected pore space. In flood basalts, the connected pore space is typically found at zones containing abundant vesicles or in breccias between basalt flows. Because central zones of flood basalts commonly are dense and impermeable without vesicles, and flows are laterally continuous over large areas and commonly stacked vertically for hundreds of meters, flow units can act as seals . The non-porous inner parts of flows may however be penetrated by networks of vertical fractures. These fractures can be open and conductive, or closed by mineralization and non-conductive.
All thermodynamic and kinetic calculations were performed using the geochemical code PHREEQC-2 . This code is capable of simulating complex interactions between dissolved gases, aqueous solutions, and mineral assemblages in batch and 1D advection–diffusion-reaction mode. As the code can only model fully saturated systems, natural systems must be simplified to end-member situations, such as given by constant pressure boundary conditions as may be the case close to underground CO2 plumes, or the assumption of packages (batches) of water reacting along a reaction path with a homogenous sediment or rock body. Based on these limitations we divided the simulations into three systems representing different parts of CO2 storage: (1) basalt alteration in the H2O-rich phase at constant CO2 pressure; (2) basalt alteration in a H2O saturated scCO2 phase, and (3) reactions at the boundary of the CO2 plume where CO2 diffuses into the aquifer from the boundary of the scCO2 plume (Figure 1). In the second case, we assumed that the CO2 phase had swept through the systems and dried out residual water, giving only dissolved water in the scCO2 phase. In this case an upper limit of carbonation potential was estimated as reactions were allowed to occur until (nearly) all water was consumed, passing the upper 2 mol/Kgw theoretical limit for the Truesdell-Jones activity model .
The standard state adopted in this study for the thermodynamic calculations was that of unit activity for pure minerals and H2O at any temperature and pressure. For aqueous species other than H2O, the standard state was unit activity of the species in a hypothetical 1 molal solution referenced to infinite dilution at any temperature and pressure. For gases, the standard state was for unit fugacity of a hypothetical ideal gas at 1 bar of pressure. All simulations used the llnl.dat database based on the thermo.com.V8.R6.230 dataset prepared at the Lawrence Livermore National Laboratory, with additions of thermodynamic data for those phases not present (see description below).
CO2 fugacity coefficients were estimated according to the modified Redlich-Kwong (SRK) equation of state  and the solubility was adjusted for by a poynting correction term (exp(v CO2 (P sat - P)/RT) where v denotes molar volume, P pressure, R the universal gas constant and T absolute temperature) . The density of CO2 at 40 C and 100 bar was approximated from Bachu and Stewart  to be 600 Kg/m3 and the solubility of water in scCO2 at the same conditions was approximated to 0.5 mole% [13, 14]
The simulations were divided into batch simulations of the H2O rich and CO2 rich phases respectively, and 1D diffusion of CO2 in the H2O rich phase to obtain information on the CO2-basalt interactions over a continuous range of CO2 pressures. The latter was solved by PHREEQC using , where C denotes molal (mol/Kgw) concentration, q denotes the sink term, subscripts t and x refer to derivatives in time and x-direction respectively, and an efficient diffusion coefficient D L of 0.45x10-9 m2/s was used for CO2 and all solutes.
where ϕ t=0 is the initial porosity, n and are moles and molar volume of mineral i respectively, and V total is the total volume of the system.
The basalt was defined to consist of a mixture of glass and crystalline basalt with mineral and glass fractions chosen based on reported data from CRB [6, 24, 25]. To represent the crystalline basalt, plagioclase (Ca0.5Na0.5Al1.5Si2.5O8) and the pyroxenes augite (Ca0.7Fe0.6 Mg0.7Si2O6) and pigeonite (Ca1.14Fe0.64 Mg0.22Si2O6) were chosen. The hydrolysis equilibrium constants of these phases were estimated using the PHREEQC program assuming ideal solid solutions of the end-members enstatite, ferrosilite and wollastonite for the pyroxenes, and albite and anorthite for the plagioclase. Equilibrium constants for the solid solutions for temperatures up to 100 C were estimated with PHREEQC and from these data coefficients a to e for the PHREEQC built-in analytical expression (log K = a + bT + c/T + dlog 10 (T) + e/T 2 ) were estimated using non-linear regression in MATLAB.
The glass composition (Ca0.015Fe0.095 Mg0.065Na0.025 K0.01Al0.105 S0.003Si0.5O1.35) was taken from  and modified by adding a small fraction of sulfur which is a common minor constituent of the CR basaltic glass .
Mineralogy included in the model
Initial Weight %
Glass Ca0.015Fe0.095Mg0.065Na0.025K0.01Al0.105 S0.003Si0.5O1.35
Smec high Fe-Mg
Composition of initial formation water
1.0 x 10-3
6.0 x 10-4
1.0 x 10-4
2.0 x 10-5
1.2 x 10-6
2.0 x 10-3
3.0 x 10-4
1.0 x 10-4
2.0 x 10-4
1.0 x 10-6
System 1: Basalt alteration in the H2O-rich phase at constant CO2pressure
i) CRB mineral and glass dissolution and formation of secondary minerals
ii) On the limitation of pore-space for the basalt carbonation
iii) Reduction of pore-space as a function of reactive surface area
System 2: The potential for carbonate growth in a H2O-saurated scCO2phase
The reaction between H2O dissolved in scCO2 and basalt was simulated at 100 bar pressure and 40 C. The initial amount of water was 0.003 Kg and no H2O was allowed to enter the system. This is an ideal end-member case and serves to illustrate the carbonation potential in a volume with limited hydration potential.
System 3: 1D diffusion of CO2into the CRB aquifer
Uncertainty on the reactive surface area
The reactive surface area is considered as a major source of uncertainty (e.g., [20, 34]) and this leads to corresponding high uncertainties in timing and extent of reactions as dissolution rates have a first order dependence on reactive surface areas. Weathering rates in nature are commonly observed to be 1–3 orders of magnitude lower than in laboratory experiments (e.g., [20, 21, 34]), and this may in part be explained by differences in reactive and physical (total) surface area between experimental and natural systems. We assumed in this study a base-case reactive surface area 1 order of magnitude lower than the estimated physical surface area for the basalt. A further two orders of magnitude reduction in the reactive surface area, which is within the range of values expected for natural systems, resulted in little basalt alteration and only minor reduction of porosity (see Figure 5). A better understanding of the surface area of porous basalt and the effect of time (aging) on features such as dislocation densities and reactive surface areas are therefore required to understand the potential for CO2 mineral storage in basaltic rocks.
Uncertainty on the choice of secondary phases used in the model
Growth rate experiments of carbonates such as magnesite and dolomite have shown that the activation energy is high and that growth is negligible at low temperatures (e.g., [35–37]). Dissolution rate studies of siderite suggests that the reaction rate is intermediate between calcite and magnesite [38, 39], and growth rate data suggest that siderite may form down to room temperature . Data on ankerite dissolution and growth is to the knowledge of the authors not known. The crystallographic and physical characteristics of ankerite do resemble those of dolomite and siderite, and the chemistry is related to dolomite with the Mg2+ substituted by various amounts of Fe2+ and Mn2+. If the growth rate is close to the magnesian carbonates such as dolomite and magnesite [41, 42], the amount that may form during low-temperature alteration is likely low. In this case, more iron would be available for siderite growth. If on the other hand the growth rate is closer to siderite, we would expect ankerite or other FeMg solid solution carbonates to grow during low-temperature alteration.
One uncertainty related to the local-equilibrium assumption is on the growth retention time for the secondary carbonates. The local-equilibrium assumption predicts growth of the secondary phases as soon as an infinitesimally small supersaturation is reached . The time it takes to nucleate sufficient mass to initiate a significant growth may however be hundreds to thousands of years for some secondary phases . There are no nucleation rate data for siderite and ankerite and the retention time is hence unknown.
Finally, the total potential for secondary carbonate growth may be affected by the amount of magnesium and iron that enters ferromagnesian calcites. As a significant fraction of the metal cations may substitute for calcium (e.g., ), a iron-magnesium rich calcite may potentially form rather than ankerite and thereby reduce the amount of siderite formed.
Comparisons to experiments, numerical simulations and natural analogues of basalt-CO2interactions
Our simulations suggest that the potential for carbonate growth is limited to siderite or FeMg carbonates at low temperatures as secondary phases such as zeolites outcompeted the carbonates for calcium. We here compare our simulated results with reported data on CO2-basalt interactions from laboratory experiments, natural analogues, and other reported numerical simulations.
The reactivity of CRB and other continental flood basalts are available from the long-term (months to years) laboratory experiments done by Schaef and co-workers [6, 24]. In these experiments basalt samples from USA, India, South Africa, and Canada were reacted with CO2 at about 100 bars and 60 to 100 C. Reacted samples from these experiments showed generation of Ca-rich carbonates interpreted as calcites with minor siderite and magnesite. In experiments on CRB using mixtures of H2S and CO2 at 60 C and 100 bar and run for 181 days, pyrite (FeS2) formed together with Mg-Fe poor calcite and a Ca-poor Fe-carbonate . Our simulations at the same temperatures show rapid formation of siderite (60 C) or siderite and magnesite at higher temperatures (Figure 3). Our simulations do not predict any calcite growth as the calcium activity is lowered by zeolite formation. Calcite would however form in our models if the zeolites were not allowed to form at local equilibrium, and possibly if a magnesian ferroan (solid solution) calcite was used in the model instead of the pure end-member calcite. Therefore, the apparent difference between our model and the experiment may be caused by our use of the local equilibrium assumption, whereas the zeolites in the laboratory experiments did not form at low temperatures due to slow kinetics. Recent experiments on basalt dissolution support the preferential release of Mg and Fe over Ca at acidic conditions , suggesting that the MgFe-carbonates will dominate as secondary carbonates during CO2 storage in basalt.
Our numerical simulations share some similarities to other works such as by Marini  and Gysi , but our model and hence the outcome is different in several aspects. The most comprehensive work done earlier is the numerical simulations done by Marini  on the reactivity of crystalline and glassy CFB following CO2 storage. The initial mineralogy was similar to our study whereas the temperature of 60 C was slightly higher than our base case 40 C. In  the CO2-basalt interactions were stretched to last for more than 280000 years compared to our 10000 years perspective. The main differences between our model and  are on the choice of secondary mineral assemblage, and on the focus of limiting factors such as the availability of water for hydration in the present work. The lack of zeolites and hydrous phases other than kaolinite and goethite in  made Ca available for secondary carbonates and the total potential for carbonate formation was higher than in our work. Marini allowed dolomite and magnesite to form at 60 C, whereas our simulations only produced siderite at the similar conditions. Moreover, the formation of dawsonite in  is still uncertain and possibly limited at high silica activities and with an assemblage of stable NaAl-silicates defined to form . Based on two different approaches, the reactive surface area for basalt was estimated to quite similar values. We estimated a specific surface area of approximately 1.5x10-5 m2/gbasalt (= 0.14 m2/Kg water at 10% porosity) based on the A p /V p values estimated by  and reported in , and reduced this value by one order of magnitude to get the reactive surface area. Marini used a geometric model giving a reactive surface area of 0.41 m2/Kg water. The higher reactive surface area and higher temperature of  resulted in faster reactions and more rapid clogging of the pore space (within a few years). Studies of natural basalt systems at similar or higher temperatures may give some insight into how fast pore space is clogged by basalt hydration or carbonation, and this should be used to improve the estimates of reactive surface areas of basalt for future studies.
Another numerical study on low-temperature (25 C, 30 bar CO2) basaltic glass alteration was presented by Gysi et al. . Again a main difference is on the choice of secondary minerals. Gysi et al.  allowed dolomite, magnesite, and Fe-Mg carbonate to form together with calcite and siderite, whereas we did not allow other Mg-Fe carbonates to form than ankerite. As previously stated, the low-temperature formation of dolomite and magnesite is not likely because of the high apparent activation energy and small kinetic coefficients for the growth of Mg-carbonates [35-37]. Other carbonates such as siderite and potentially FeMg-calcites are more likely to form at these low temperatures. The high reactive surface area used in  is based on a geometric model for glass fragments, and is hence not directly comparable with the surface area estimated for a vesicle pore space of a solid basalt. Although no inverse modeling was done to estimate the reactive surface area of the basalt in , fragmented basaltic rocks such as hyaloclastite breccias are expected to have significantly higher reactive surface areas than porous solid basalts, and they are therefore correspondingly more reactive.
One example of a natural analogue that shed light on CO2 basalt interactions is the CO2 charged basalt hosted groundwaters at Hekla, Iceland. Solution aqueous species sampled from natural cold springs and rivers here showed a drop in total inorganic carbon (TIC) that was interpreted to result from considerable formation of secondary carbonate phases such as calcite . Reaction path modeling of the system suggests however that the carbonate formation is associated with high pH in accordance with the low TIC in the sampled waters. This system is therefore different from basalt CO2 storage projects where higher CO2 pressures may be maintained over time and the pH is lower. In addition to calcite, dolomite was also suggested as a potential storage host for the low temperature reactions in Hekla . This may however be questionable as long-term laboratory experiments at room temperature have failed to form dolomite even at significant super saturations , explained by the high activation energy for dolomite growth [32, 41]. Another natural analogue that more closely corresponds to industrial CO2 storage is the basalt-hosted petroleum reservoir on Nuussuaq, West Greenland. In this system the bulk carbonate formation appears to have occurred as secondary weathering products. Other alteration products such as zeolites and oxides were replaced by dolomite, magnesite, siderite, and calcite at temperatures of 70–120 C . Therefore, taking into account the basalt weathering products and not only primary basalt minerals appears to be vital in estimating the total potential for secondary carbonate formation and the long-term potential for CO2 storage in basalt systems.
Simulations of closed-system (PCO2 = 100 bar, 40 C) and 1D reaction–diffusion (PCO2 = 0–100 bar, 40 C) alteration of basalt suggest that the potential of secondary carbonate formation is limited to siderite at low temperatures as divalent metal cations are preferentially consumed by zeolites and oxides. Higher temperatures 60 – 100 C appear to be in favor of secondary carbonate formation, allowing the precipitation of carbonates such as magnesite, siderite and possibly dolomite and other FeMg carbonates (ankerite). Given an unlimited source of CO2 (fixed CO2 pressure), the total amount of CO2 stored as solid carbonates is orders of magnitude higher than the 1–2 mol/Kg water solubility in the formation water (Figure 4). The total amount trapped might however be reduced if CO2, H2O or pore space are limiting factors. The formation of secondary hydrous and carbonate phases increases the volume of solids and the porosity is correspondingly reduced (Figure 5). This together with the immobilization of CO2 by solid carbonate formation is in favor of safe long-term storage of CO2 in basaltic aquifers.
We highly appreciated constructive comments and suggestions from the reviewers. This work has been funded by SSC-Ramore (Subsurface storage of carbon dioxide - risk assessment, monitoring and remediation) project and (partially) by SUCCESS centre for CO2 storage under grant 193825/S60 from Research Council of Norway (RCN). SUCCESS is a consortium with partners from industry and science, hosted by Christian Michelsen Research as.
- Holloway S: Underground sequestration of carbon dioxide–a viable greenhouse gas mitigation option. Energy. 2004, 30: 2318-2333.View ArticleGoogle Scholar
- Bachu S, Bonijoly D, Bradshaw J, Burruss R, Holloway S, Christensen NP, Mathiassen OM: CO2 storage capacity estimation: Methodology and gaps. Int J Greenhouse Gas Control. 2007, 1: 430-443. 10.1016/S1750-5836(07)00086-2.View ArticleGoogle Scholar
- IPCC Special Report on Carbon Dioxide Capture and Storage. In Prepared by Working Group III of the Intergovernmental Panel on Climate Change. Edited by: Metz B, Davidson O, Coninck H, Loos M, Meyer L. 2005, Cambridge University Press, Cambridge, United Kingdom and New York, NY, USA, 442 pp-Google Scholar
- Oelkers EH, Gislason SR, Matter J: Mineral Carbonation of CO2. Elements. 2008, 4: 333-337. 10.2113/gselements.4.5.333.View ArticleGoogle Scholar
- McGrail BP, Schaef HT, Ho AM, Chien Y-J, Dooley JJ, Davidson CL: Potential for carbon dioxide sequestration in flood basalts. J Geophysical Res-Solid Earth. 2006, 111: 1-13.View ArticleGoogle Scholar
- Schaef HT, McGrail BP, Owen AT: Carbonate mineralization of volcanic province basalts. Int J Greenhouse Gas Control. 2010, 4: 249-261. 10.1016/j.ijggc.2009.10.009.View ArticleGoogle Scholar
- Schaef HT, Windish CF, McGrail BP, Martin PF, Rosso KM: Brucite [Mg(OH)2] carbonation in wet supercritical CO2: An in situ high pressure x-ray diffration study. Geochim Cosmochim Acta. 2011, 75: 7458-7471. 10.1016/j.gca.2011.09.029.View ArticleGoogle Scholar
- White MD, McGrail BP, Schaef HT, Hu JZ, Hoyt DW, Felmy AR, Rosso KM, Wurstner SK: Multiphase sequestration geochemistry: Model for mineral carbonation. Energy Procedia. 2011, 4: 5009-5016.View ArticleGoogle Scholar
- Parkhurst DL, Appelo CAJ: User's guide to PHREEQC (version 2) - a computer program for speciation, reaction-path, 1D-transport, and inverse geochemical calculations. 1999, U.S. Geological Survey, Water-Resources Investigation Report, 312-Google Scholar
- Soave G: Equilibrium constants from a modified Redlich-Kwrong equation of state. Chem Eng Sci. 1972, 27: 1197-1203. 10.1016/0009-2509(72)80096-4.View ArticleGoogle Scholar
- Hellevang H, Kvamme B: ACCRETE - Geochemistry solver for CO2-water-rock interactions. Proceedings GHGT 8 conference. 2006, 8p-Google Scholar
- Bachu S, Stewart S: Geological sequestration of anthropogenic carbon dioxide in the western Canada sedimentary basin: Suitability analysis. J Can Pet Technol. 2002, 41: 32-40.View ArticleGoogle Scholar
- Coan CR, King AD: Solubility of Water in Compressed Carbon Dioxide, Nitrous Oxide, and Ethane - Evidence for Hydration of Carbon Dioxide and Nitrous Oxide in Gas Phase. J Am Chem Soc. 1971, 93: 1857-1862. 10.1021/ja00737a004.View ArticleGoogle Scholar
- Sabirzyanov AN, Ilìn AP, Akhunov AR, Gumerov FM: Solubility of water in Supercritical carbon dioxide. High Temp. 2002, 40: 203-206. 10.1023/A:1015294905132.View ArticleGoogle Scholar
- Bahar M, Lui K: Measurement of the diffusion coefficient of CO2 in formation water under reservoir conditions: Implication for CO2 storage. 2008, Perth, AustraliaGoogle Scholar
- Palandri JL, Kharaka YK: A Complilation of Rate Parameters of Water-Mineral Interaction Kinetics for Applicatin to Geochemical Modeling. U.S. Geological survey, Open report. 2004, 1068-Google Scholar
- Gislason SR, Oelkers EH: Mechanism, rates, and consequences of basaltic glass dissolution: II. An experimental study of the dissolution rates of basaltic glass as a function of pH and temperature. Geochim Cosmochim Acta. 2003, 67: 3817-3832. 10.1016/S0016-7037(03)00176-5.View ArticleGoogle Scholar
- Marini L: Chapter 7: Reaction path Modelling of geological CO2 sequatration. 2006, Dev Geochem, 319-394.Google Scholar
- Wolff-Boenisch D, Gislason SR, Oelkers EH: The effect of fluoride on the dissolution rates of natural glasses at pH 4 and 25 degrees C. Geochim Cosmochim Acta. 2004, 68: 4571-4582. 10.1016/j.gca.2004.05.026.View ArticleGoogle Scholar
- White AF, Peterson ML, et al: Role of reactive-surface-area characterization in geochemical kinetic models. In. Chemical Modeling of Aqueous Systems II. Edited by: Melchior D. 1990, ACS Symposium Series, ACS, Washington, DCGoogle Scholar
- White AF, Brantley SL: The effect of time on the weathering of silicate minerals: why do weathering rates differ in the laboratory and field?. Chem Geol. 2003, 202: 479-506. 10.1016/j.chemgeo.2003.03.001.View ArticleGoogle Scholar
- Gudbrandsson S, Wolff-Boenisch D, Gislason SR, Oelkers EH: An experimental study of crystalline basalt dissolution from 2 ≤ pH ≤ 11 and temperatures from 5 to 75 C. Geochim Cosmochim Acta. 2011, 75: 5496-5509. 10.1016/j.gca.2011.06.035.View ArticleGoogle Scholar
- Helgeson HC: Evaluation of Irreversible Reactions in Geochemical Processes Involving Minerals and Aqueous Solutions .I. Thermodynamic Relations. Geochim Cosmochim Acta. 1968, 32: 853-877. 10.1016/0016-7037(68)90100-2.View ArticleGoogle Scholar
- Schaef HT, McGrail BP, Owen AT: Basalt-CO2-H2O Interactions and Variability in Carbonate Mineralization Rates. Energy Procedia. 2009, 1: 4899-4906. 10.1016/j.egypro.2009.02.320.View ArticleGoogle Scholar
- Schaef HT, McGrail BP: Dissolution of Columbia River Basalt under mildly acidic conditions as a function of temperature: Experimental results relevant to the geological sequestration of carbon dioxide. Appl Geochem. 2009, 24: 980-987. 10.1016/j.apgeochem.2009.02.025.View ArticleGoogle Scholar
- Blake S, Self S, Sharma K, Sephton S: Sulfur release from the Columbia River Basalts and other flood lava eruptions constrained by a model of sulfide saturation. Earth Planet Sci Lett. 2010, 299: 328-338. 10.1016/j.epsl.2010.09.013.View ArticleGoogle Scholar
- Neuhoff PS, Fridriksson T, Arnorsson S, Bird DK: Porosity evolution and mineral paragenesis during low-grade metamorphism of basaltic lavas at Teigarhorn, eastern Iceland. Am J Sci. 1999, 299: 467-501. 10.2475/ajs.299.6.467.View ArticleGoogle Scholar
- Neuhoff PS, Rogers KL, Stannius LS, Bird DK, Pedersen AK: Regional very low-grade metamorphism of basaltic lavas, Disko-Nuussuaq region, West Greenland. Lithos. 2006, 92: 33-54. 10.1016/j.lithos.2006.03.028.View ArticleGoogle Scholar
- Stefansson A, Gislason SR, Arnorsson S: Dissolution of primary minerals in natural waters - II. Mineral saturation state. Chem Geol. 2001, 172: 251-276.Google Scholar
- Reidel SP, Johnson VG, Spane FA: Natural gas storage in basalt aquifers of the Colombia Basin, Pacific Northwest USA: A guide to site characterization. 2002, Report Pacific Northwest National Laboratory; PNNL-13962, HView ArticleGoogle Scholar
- Pham VTH, Lu P, Aagaard P, Zhu C, Hellevang H: On the potential of CO2-water-rock interactions for CO2 storage using a modified kinetic model. International Journal of Greenhouse Gas Control. 2011, 5: 1002-1015. 10.1016/j.ijggc.2010.12.002.View ArticleGoogle Scholar
- Reidel SP: A lava flow without a source: the cohassett flow and its compositional components, Sentinel Bluffs Member, Columbia River Basalt Group. J Geol. 2005, 113: 1-21. 10.1086/425966.View ArticleGoogle Scholar
- Peuble S, Godard M, Gouze P, Luquot L: CO2 sequestration in olivine rich basaltic aquifers: a reactive percolation experimental study. Geochim Cosmochim Acta. 2009, 73: A1635-Google Scholar
- Velbel MA: Constancy of silicate-mineral weathering-rate ratios between natural and experimental weathering: implications for hydrologic control of differences in absolute rates. Chem Geol. 1993, 105: 89-99. 10.1016/0009-2541(93)90120-8.View ArticleGoogle Scholar
- Saldi GD, Jordan G, Schott J, Oelkers EH: Magnesite growth rates as a function of temperature and saturation state. Geochim Cosmochim Acta. 2009, 73: 5646-5657. 10.1016/j.gca.2009.06.035.View ArticleGoogle Scholar
- Arvidson RS, Mackenzie FT: The dolomite problem; control of precipitation kinetics by temperature and saturation state. Am J Sci. 1999, 299: 257-288. 10.2475/ajs.299.4.257.View ArticleGoogle Scholar
- Arvidson RS, Mackenzie FT: Tentative kinetic model for dolomite precipitation rate and its application to dolomite distribution. Aquat Geochem. 1996, 2: 273-298. 10.1007/BF01160046.View ArticleGoogle Scholar
- Golubev SV, Bébnézeth P, Schott J, Dandurand JL, Castillo A: Siderite dissolution kinetics in acidic aqueous solutions from 25 to 100 C and 0 to 50 atm pCO2. Chem Geol. 2009, 265: 13-19. 10.1016/j.chemgeo.2008.12.031.View ArticleGoogle Scholar
- Pokrovsky OS, Golubev SV, Schott J, Castillo A: Calcite, dolomite and magnesite dissolution kinetics in aqueous solutions at acid to circumneutral pH, 25 to 150 C and 1 to 55 atm pCO2: New constraints on CO2 sequestration in sedimentary basins. Chem Geol. 2009, 265: 20-32. 10.1016/j.chemgeo.2009.01.013.View ArticleGoogle Scholar
- Greenberg J, Tomson M: Precipitation and dissolution kinetics and equilibria of aqueous ferrous carbonate vs temperature. Appl Geochem. 1992, 7: 185-190. 10.1016/0883-2927(92)90036-3.View ArticleGoogle Scholar
- Arvidson RS, Mackenzie FT: Tentative Kinetic Model for Dolomite Precipitation Rate and Its Application to Dolomite Distribution. Aquat Geochem. 1997, 2: 273-298. 10.1007/BF00119858.View ArticleGoogle Scholar
- Saldi GD, Jordan G, Schott J, Oelkers EH: Magnesite growth rates as function of temperature and saturation state: An HAFM study. Geochim Cosmochim Acta. 2009, 73: A1149-Google Scholar
- Busenberg E, Plummer LN: Thermodynamics of magnesian calcite solid-solutions at 25 C and 1 atm total pressure. Geochim Cosmochim Acta. 1989, 53: 1189-1208. 10.1016/0016-7037(89)90056-2.View ArticleGoogle Scholar
- Gysi AP, Stefansson A: Numerical modelling of CO2-water-basalt interaction. Mineral Mag. 2008, 72: 55-59. 10.1180/minmag.2008.072.1.55.View ArticleGoogle Scholar
- Hellevang H, Declercq J, Aagaard P: Why is dawsonite absent in CO2 charged reservoirs? Oil & Gas Science and Technology - Re. IFP Energies nouvelles. 2011, 66: 119-135. 10.2516/ogst/2011002.View ArticleGoogle Scholar
- Saar MO, Manga M: The relationship between permeability, porosity, and microstructure in vesicular basalts. 1998, Master Thesis, Univ. of Oregon, Eugene, 91-Google Scholar
- Flaathen TK, Gislason SR, Oelkers EH, Sveinbjornsdottir AE: Chemical evolution of the Mt. Hekla, Iceland, groundwaters: a natural analogue for CO2sequestration in basaltic rocks. Appl Geochem. 2009, 24: 463-474. 10.1016/j.apgeochem.2008.12.031.View ArticleGoogle Scholar
- Land LS: Failure to precipitate dolomite at 25 degrees C from dilute solution despite 1000-fold oversaturation after 32 years. Aquat Geochem. 1998, 4: 361-368. 10.1023/A:1009688315854.View ArticleGoogle Scholar
- Rogers KL, Neuhoff PS, Pedersen AK, Bird DK: CO2metasomatism in a basalt-hosted petroleum reservoir, Nuussuaq, West Greenland. Lithos. 2006, 92: 55-82. 10.1016/j.lithos.2006.04.002.View ArticleGoogle Scholar
This article is published under license to BioMed Central Ltd. This is an Open Access article distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/by/2.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.