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The Experimental Determination of Hydromagnesite Precipitation Rates at 22.575ºC

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The Experimental Determination of Hydromagnesite Precipitation Rates at 22.5�75ºC Mineralogical Magazine, November 2014, Vol. 78(6), pp. 1405–1416 OPEN ACCESS The experimental determination of hydromagnesite precipitation rates at 22.5À75ºC 1,2, 2 1 1,3 U.-N. BERNINGER *, G. JORDAN ,J.SCHOTT AND E. H. OELKERS 1 Ge´oscience Environnement Toulouse, CNRS-UPS-OMP, 14 av. E´ douard Belin, 31400 Toulouse, France 2 Department fu¨r Geo- und Umweltwissenschaften, LMU, Theresienstr. 41, 80333 Mu¨nchen, Germany 3 Department of Earth Sciences, UCL, Gower Street, London WC1E 6BT, UK [Received 18 July 2014; Accepted 17 October 2014; Associate Editor: T. Stawski] ABSTRACT Natural hydromagnesite (Mg5(CO3)4(OH)2·4H2O) dissolution and precipitation experiments were performed in closed-system reactors as a function of temperature from 22.5 to 75ºC and at 8.6 < pH < + 2+ À 10.7. The equilibrium constants for the reaction Mg5(CO3)4(OH)2·4H2O+6H =5Mg +4HCO3 + 6H2O were determined by bracketing the final fluid compositions obtained from the dissolution and precipitation experiments. The resulting constants were found to be 1033.70.9,1030.50.5 and 1026.50.5 at 22.5, 50 and 75ºC, respectively. Whereas dissolution rates were too fast to be determined from the experiments, precipitation rates were slower and quantified. The resulting BET surface area- normalized hydromagnesite precipitation rates increase by a factor of ~2 with pH decreasing from 10.7 to 8.6. Measured rates are approximately two orders of magnitude faster than corresponding forsterite dissolution rates, suggesting that the overall rates of the low-temperature carbonation of olivine are controlled by the relatively sluggish dissolution of the magnesium silicate mineral. KEYWORDS: hydromagnesite, carbon storage, carbonate minerals, precipitation. Introduction temperatures (e.g. Burton and Walter, 1987; THE precipitation rates of carbonate minerals are Zuddas and Mucci, 1998; Teng et al., 2000; of current interest due to their potential role in Lakshtanov and Stipp, 2010; Rodriguez-Blanco et carbon storage (e.g. Seifritz, 1990; Lackner et al., al., 2011; Ruiz-Agudo et al., 2013), the precipita- 1995; Xu et al., 2005; Marini, 2007; Flaathen et tion of the anhydrous Mg-carbonate mineral al., 2011). Carbon storage via carbonate mineral magnesite is apparently inhibited at temperatures precipitation appears to be particularly favoured of <~80ºC (e.g. Saldi et al., 2009, 2012). At these in basaltic and ultramafic rocks due to their high lower temperatures, the precipitation of hydrous reactivity and the prevalence of divalent cations Mg carbonates, such as hydromagnesite and such as Ca2+ and Mg2+ (McGrail et al., 2006; dypingite are favoured (Mavromatis et al., 2012; Alfredsson et al., 2008; Oelkers et al., 2008; Shirokova et al., 2013). The goal of the present Matter et al., 2009, 2011; Schaef et al., 2010, study was an improved understanding of the 2011; Gislason et al., 2010; Gysi and Stefansson, precipitation kinetics of hydrous magnesium 2012; Arado´ttir et al., 2012; Oskierski et al., carbonate minerals at temperatures below 80ºC, 2013; Gislason and Oelkers, 2014). Whereas in part to assess their potential to store carbon at anhydrous calcium carbonates such as calcite and aragonite precipitate readily at ambient This paper is published as part of a special issue in Mineralogical Magazine, Vol. 78(6), 2014 entitled * E-mail: [email protected] ‘Mineral–fluid interactions: scaling, surface reactivity DOI: 10.1180/minmag.2014.078.6.07 and natural systems’. # 2014 The Mineralogical Society U.-N. BERNINGER ET AL. ambient temperatures. Hydromagnesite growth where R is the gas constant (= 8.3144 J/Kmol), T rates were measured in closed-system reactors at is absolute temperature, IAPHmg represents the 22.5À75ºC. The results of these experiments are aqueous ion activity product, and OHmg represents used here to evaluate the potential of hydrous the saturation state of the aqueous phase with magnesium carbonate minerals as a carbon- respect to hydromagnesite. All thermodynamic storage host. calculations in this study were performed using the PHREEQC for Windows (Version 2.18) computer code (Parkhurst and Appelo, 1999) Theoretical background together with its llnl database after adding the The standard state adopted in this study for equilibrium constants for Mg2+ hydrolysis and the thermodynamic calculations is that of unit activity carbonic acid dissociation reported by Brown et for pure minerals and H2O at any temperature and al. (1996) and Millero et al. (2007), respectively. pressure. For aqueous species other than H2O, the standard state is unit activity of the species in a Materials and methods hypothetical 1 molal solution referenced to infinite dilution at any temperature and pressure. Natural hydromagnesite (Mg5(CO3)4(OH)2·4H2O) Hydromagnesite dissolution can be described was collected as stromatolite precipitates from using: Lake Salda in southwest Turkey (cf. Shirokova et + al., 2011) then hand-milled with an agate mortar Mg5(CO3)4(OH)2·4H2O+6H = 2+ À and pestle. The resulting powder was cleaned with 5Mg + 4HCO3 +6H2O (1) H2O2 to remove organics and oven dried at 50ºC Taking account of this standard state, the law of for 24 h. The particles were not cleaned mass action for reaction 1 is given by: ultrasonically due to the fine nature of these 5 4 natural grains. Scanning Electron Microscope a 2þ Á a À Mg HCO3 (SEM) images, obtained using a JEOL JSM- KHmg ¼ 6 ð2Þ aHþ 6360 LV microscope, of the resulting hydro- magnesite are shown in Fig. 1. The hydro- where KHmg stands for the equilibrium constant of magnesite powder prepared consists of reaction 1 and ai represents the activity of the agglomerated crystals. The size of the individual subscripted aqueous species. The chemical hydromagnesite crystals ranges up to ~5 mm, the affinity (A) of reaction 1 can be expressed as: agglomerates range up to 50 mm. No other 8 9 mineral phases were evident from SEM images >IAPHmg> A ¼RT ln: ; ¼RT ln OHmg ð3Þ nor were any detected using backscattered KHmg electron microscopy. The purity of this FIG. 1. SEM images of the initial hydromagnesite seed material (a,b), after the precipitation experiment at 22.5ºC at pH 10.72 (c,d), and after the experiment performed at 75ºC at pH 9.75 (e,f). 1406 HYDROMAGNESITE PRECIPITATION RATES hydromagnesite was further verified by X-ray The fluid in each reactor was sampled regularly diffraction using an INEL CPS-120 diffractometer to monitor reaction progress. The reactive fluid with CoKa radiation, l = 1.78897 A˚ , and a Mg concentration was measured by flame atomic graphite monochromator. X-ray diffraction was absorption spectroscopy using a Perkin Elmer performed from 1 to 110º2y at 0.09º2y/min and at Zeeman 5000 Atomic Absorption Spectrometer a step size of 0.029º2y. This analysis revealed no with an uncertainty of 2% and a detection limit phase other than hydromagnesite. The surface of 6610À7 mol/kg. Alkalinity was determined by area of the hydromagnesite prepared was standard HCl titration using Schott TA 10plus determined to be 8.5 m2/g Ô 10% by multipoint with an uncertainty of 2% and a detection limit krypton adsorption according to the BET method of 5610À5 eq/l. Fluid pH measurements were (Brunauer et al., 1938) using a Quantachrome Gas performed at 22.5ºC immediately after sampling Sorption system. using a standard glass electrode, previously Hydromagnesite dissolution and precipitation calibrated with 4.01, 6.86 and 9.18 NIST pH experiments were performed in closed-system buffers. 0.5 l high-density polypropylene reactors. Approximately 1 g of hydromagnesite powder Results and 500 g of an aqueous sodium carbonate buffer solution with ionic strength of 0.1 mol/kg were Aqueous Mg concentrations as a function of time placed into each reactor. The reactors were during a representative experiment are illustrated subsequently sealed and placed in temperature- in Fig. 2. An approximately constant Mg concen- controlled shaking water baths. The initial solu- tration was obtained by dissolution within tions were composed of deionized H2OandMerck ~30 min. Following the addition of MgCl2 to reagent-grade Na2CO3 and NaHCO3 at varying the fluid phase, the aqueous Mg concentration ratios (Table 1); the resulting aqueous solutions decreased systematically with time. Note the had a pH ranging from 7.7 to 11.2. The original near-constant aqueous Mg concentration attained hydromagnesite was allowed to dissolve in this at the end of the dissolution leg is distinct from fluid for about 5 days until a steady-state aqueous that attained at the end of the precipitation leg of Mg concentration was attained and validated. After the experiment, reflecting the different fluid pH this time, ~3 g of a 1.81 mol/kg aqueous MgCl2 present during these legs. solution was added to each reactor to supersaturate The rate at which steady-state was attained the fluid with respect to hydromagnesite. during hydromagnesite dissolution was too rapid FIG. 2. Temporal evolution of aqueous Mg concentration for a representative experiment performed at 50ºC. This experiment began by the dissolution of hydromagnesite until a near-stationary state was attained. After 100 h, additional Mg was added to the aqueous fluid and precipitation began. Uncertainties in the concentration measurements are within the size of the symbols. Note that steady-state concentration of aqueous Mg during the dissolution leg is reached in <30 min. The initial and final pH values of the fluid during the dissolution and precipitation leg of the experiment are shown. 1407 TABLE 1. Summary of experimental data. T Hydromagnesite
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