Ultramafic rocks, such as the Semail Ophiolite in the Sultanate of Oman, are considered to be a potential storage site for CO2. This type of rock is rich in divalent cations that can react with dissolved CO2 and form carbonate minerals, which remain stable over geological periods of time. Dissolution of the ophiolite mobilizes heavy metals, which can threaten the safety of surface and groundwater supplies but secondary phases, such as iron oxides, clays and carbonate minerals, can take up significant quantities of trace elements both in their structure and adsorbed on their surfaces.

Hyperalkaline spring waters issuing from the Semail Ophiolites can have pH as high as 12. This water absorbs CO2 from air, forming carbonate mineral precipitates either as thin crusts on the surface of placid water pools or bottom precipitates in turbulent waters. We investigated the composition of the spring water and the precipitates to determine the extent of trace element uptake. We collected water and travertine samples from two alkaline springs of the Semail Ophiolite. Twenty seven elements were detected in the spring waters. The bulk of the precipitate was CaCO3 in aragonite, as needles, and rhombohedral calcite crystals. Traces of dypingite (Mg5(CO3)4(OH)2·5H2O) and antigorite ((Mg,Fe)3Si2O5(OH)4) were also detected. The bulk precipitate contained rare earth elements and toxic metals, such as As, Ba, Cd, Sr and Pb, which indicated scavenging by the carbonate minerals. Boron and mercury were detected in the spring water but not in the carbonate phases. The results provide confidence that many of the toxic metals released by ophiolite dissolution in an engineered CO2 injection project would be taken up by secondary phases, minimizing risk to water quality.

Capturing and storing CO2 in rock as carbonate minerals, formed during injection of CO2 into geological formations, has great potential for reducing atmospheric concentrations of anthropogenic CO2. Some carbonate minerals form readily in nature but by increasing our understanding, the processes could be enhanced and used on an industrial scale. Improved understanding of carbonate mineral formation would also contribute towards more robust models for predicting the global carbon balance (Mackenzie and Andersson, 2013). Field-scale projects, where CO2-charged waters are injected into the subsurface, have already begun, e.g. into basaltic formations in Iceland and northwestern USA (Alfredsson et al., 2008; Oelkers et al., 2008; Gislason et al., 2010; McGrail et al., 2011; Alfredsson et al., 2013; Gislason and Oelkers, 2014). Carbon dioxide-enriched waters are corrosive (at 1 atm CO2, pH = 3.6) and when injected into geological formations, they attack the host rock, increasing weathering rates. Mafic and ultramafic rock releases ions such as Ca, Mg and Fe, which can combine with the dissolved CO2 and form stable carbonate phases. The overall reaction is energetically favourable (Lackner et al., 1995) but depends on fluid composition, ionic strength, reduction potential, pH, temperature and total pressure (Stumm, 1992; Krauskopf and Bird, 1994; Gislason et al., 2010; Olsson et al., 2012). Mafic and ultramafic rocks are well suited to CO2 storage because of their large divalent cation:Si ratio, and thus greater capacity for carbon capture, and lower tendency to form clay minerals, which use up cations, thus decreasing carbonation potential. Reaction in these rock formations neutralizes the acidic waters, thus promoting carbonate mineral precipitation. However, injecting CO2-charged waters, and the ensuing dissolution of the host rock, can mobilize toxic heavy metals, such as As, Ba, Cd, Cr, F, Hg and Pb (Turekian and Wedepohl, 1961), which could pose a risk for surface and groundwater supplies (Aiuppa et al., 2000; Flaathen et al., 2009; Galeczka et al., 2013; Olsson et al., 2013; Galeczka et al., 2014; Olsson et al., 2014).

The Semail Ophiolite of the Sultanate of Oman is a potential site for CO2-charged water injection for carbon sequestration (Kelemen and Matter, 2008; Matter and Kelemen, 2009; Paukert et al., 2012). It is one of the largest and best preserved ophiolites in the world and covers ~30,000 km2, with an average thickness of ~5 km (Neal and Stanger, 1984b; Nicolas et al., 2000; Chavagnac et al., 2013a). About 30% of this volume is the ultramafic rock, peridotite, which contains mostly olivine and pyroxene and <45% silica (Nicolas et al., 2000; Kelemen and Matter, 2008). Groundwater dissolves the peridotite by reactions such as:
formula
(1)
releasing cations and increasing pH. At low partial CO2 pressure, intermediate products associated with the incomplete dissolution of peridotite are typically serpentine (Mg,Fe)3Si2O5 (OH)4) and talc (Mg3Si4O10(OH)2). Spring waters discharging from ultramafic rocks are extremely alkaline, with a pH as high as 12 and CO2 partial pressure <10−8 bars (Barnes et al., 1978; Neal and Stanger, 1984a; Paukert et al., 2012; Chavagnac et al., 2013b). As the hyperalkaline water emerges at the Earth's surface, it equilibrates with atmospheric CO2, carbonate minerals become supersaturated at the air–water interface and precipitate (Fig. 1) by the overall reaction:
(Ca,Fe,Mg)2++CO2+H2O(Ca,Fe,Mg)CO3+2H+
(2)
Fig. 1.

Precipitation of carbonate minerals at the locations where Sample OM01 (a) and Sample OM08 (b) were collected. In calm hyperalkaline waters, the precipitation rate is controlled by the diffusion of CO2 from the atmosphere through the air–water interface. Carbonate minerals form as a thin crust that looks like ice (a). In fast-moving spring waters (white waters), where the air–water surface film is broken, the nucleation and precipitation is homogenous within the water body (b).

Fig. 1.

Precipitation of carbonate minerals at the locations where Sample OM01 (a) and Sample OM08 (b) were collected. In calm hyperalkaline waters, the precipitation rate is controlled by the diffusion of CO2 from the atmosphere through the air–water interface. Carbonate minerals form as a thin crust that looks like ice (a). In fast-moving spring waters (white waters), where the air–water surface film is broken, the nucleation and precipitation is homogenous within the water body (b).

As protons are released, pH decreases. The minerals observed are calcite and aragonite (both CaCO3), as well as brucite (Mg(OH)2) and antigorite ((Mg,Fe)3Si2O5(OH)4) (Neal and Stanger, 1984a; Chavagnac et al., 2013a). Spring waters in natural systems can be considered as an analogue for CO2 injection into ophiolites in general (Kelemen and Matter, 2008; Beinlich et al., 2010, 2012; Paukert et al., 2012) and by studying the uptake of trace metals into the secondary phases, we can gain information about their mobility as a result of CO2 injection. The aim of this study was to analyse ultramafic spring waters and the natural precipitated products to investigate the extent of trace element uptake into the solids and loss from the solution phase.

The samples were collected on 13 and 14 January 2011, a total of eight liquid and eight solid samples from two alkaline springs in Oman. The springs were located near the village of Falaji (Spring #1) and Qafifah (Spring #2) discharging from partly serpentinized peridotite and gabbroic rocks. Samples from the two springs and the surroundings have been described elsewhere (Nicolas et al., 2000; Kelemen and Matter, 2008; Matter and Kelemen, 2009; Paukert et al., 2012; Chavagnac et al., 2013a,b). pH and alkalinity were measured in the field. The pH was measured using a plastic body, double junction electrode, connected to a CyberScan 310 pH meter, calibrated against certified buffers (pH 4.01, 7.00, 10.01 and 12.46) from Oakton. Electrode and buffers were at the same temperature as the spring water. Alkalinity was determined by Gran titration, using HCl (Gran, 1952; Stumm and Morgan, 1981). Samples for ion chromatography (IC) and inductively coupled plasma optical emission spectrometry (ICP-OES) were filtered in the field through 0.2 μm cellulose acetate membranes. They were collected in plastic vials that had been flushed three times with the sample. Samples for ICP-OES were acidified with concentrated, suprapure HNO3 (1:100) onsite to prevent precipitation of solids during transport. An overview of the fluid samples and locations is found in Table 1.

Solid samples were collected in plastic bags and freeze dried for 24 h in the laboratory to minimize alteration. The solids were characterized by means of standard X-ray powder diffraction (XRD) and scanning electron microscopy (SEM), using instrument conditions and procedures described by Olsson et al. (2014). Residual water in the pores could not be removed by freeze drying and is responsible for precipitation of halite (NaCl) observed in the XRD patterns. The solid samples were flash dissolved in a 1% HNO3 solution which dissolves the carbonate mineral phases while minimizing the dissolution of possible silicate mineral grains originating from the ophiolite. Selected spring water and flash-dissolved samples were analysed at Analytica-SGAB, Luleå, Sweden, using ICP-OES and inductively coupled plasma sector field mass spectroscopy (ICP-SFMS). Mercury was measured with atomic fluorescence spectroscopy (AFS). The elements measured by IC were Cl, F and SO42 and by ICP-OES were Al, As, B, Ba, Br, Ca, Cl, Cr, Fe, K, Li, Na, Mg, Mn, Mo, P, S, Sb, Si, Sr, Ti, V and W. In-house, multi-element standards were used for calibration and these are regularly tested against the Standard Canadian River Waters (SLRS-4) and single element international standards (SPEX). The uncertainties of the IC and ICP-OES data were estimated from analytical measurements. The reproducibility of duplicate samples for both methods is within 5%.

Table 1.

Summary of the sample data. At each location, a sample of travertine and spring water were collected. Samples OM01–05 were collected on 13 January 2011 in Spring #1 and Samples OM06, 07 and 08 were collected on 14 January 2011 in Spring #2.

Sample setDistance from spring (m)LocationWater temperature (°C)pHTotal alkalinity (meq/l)
LatitudeLongitude
Spring #1
OM01022.837858.056830.611.65.9
OM025 ± 322.837958.056829.611.65.9
OM0310 ± 322.837958.056726.811.76.3*
OM0413 ± 322.837958.056725.311.65.8
OM0518 ± 322.838058.056724.011.65.9
Spring #2
OM06022.904558.424520.311.53.2
OM0715 ± 322.904458.424622.310.82.2
OM0830 ± 322.904258.424920.310.91.9
Sample setDistance from spring (m)LocationWater temperature (°C)pHTotal alkalinity (meq/l)
LatitudeLongitude
Spring #1
OM01022.837858.056830.611.65.9
OM025 ± 322.837958.056829.611.65.9
OM0310 ± 322.837958.056726.811.76.3*
OM0413 ± 322.837958.056725.311.65.8
OM0518 ± 322.838058.056724.011.65.9
Spring #2
OM06022.904558.424520.311.53.2
OM0715 ± 322.904458.424622.310.82.2
OM0830 ± 322.904258.424920.310.91.9
*

No filtration

The geochemical speciation program PHREEQC (Parkhurst and Appelo, 1999), with the MINTEQ database (Allison et al., 1991), was used to determine the charge balance of the spring water samples, to estimate the dissolved inorganic carbon (DIC) and the saturation state of selected solid phases. The saturation state of lansfordite (MgCO3·5H2O) was determined in a parallel PHREEQC operation run using the LLNL database (Wolery, 1992). The charge balance of all of the spring water samples was within 2% and ionic strength was <20 mm.

Spring-water chemistry

Five water samples were collected from Spring #1 and its runoff and three from Spring #2 and its runoff, at a series of sites downstream from the springs. At Spring #1, the pH was 11.6 and total alkalinity was 5.9 meq/l and these values remained constant downstream (Table 1). A summary of the dissolved components can be found in Tables 2 and 3. With the exception of the DIC, Mg, Ni and Zn, concentrations remained constant with distance downstream for Spring #1 (Fig. 2). The DIC increased as expected because of the influx of CO2 from the atmosphere. Mg concentration was very low, ~6.3 μm, probably because it was controlled by the formation of serpentine, brucite or hydrotalcite (Neal and Stanger, 1984b; Paukert et al., 2012; Chavagnac et al., 2013b). The water in the pool near this spring was covered by a thin crust of solid CaCO3, which inhibited influx of atmospheric CO2. The constant composition of the spring water suggests that relatively little of the dissolved elements formed precipitates. The reasons are: (1) the short period during which precipitates could form during the time the water emerged from the spring outlet to 20 to 30 m downstream, where the last sample was collected; and (2) carbonate minerals could not precipitate because the concentration of dissolved carbonate was hindered by the thin crust of CaCO3 at the water surface (Fig. 1a). The influx of CO2 depends on water turbulence, i.e. dissolution is faster when the gas forms foam and bubbles during mixing, such as at waterfalls, breaking waves or rapid transport (Jähne et al., 1987; Gislason et al., 1996). The water of Spring #1 experienced only slight turbulence, therefore slow CO2 influx.

At Spring #2, pH decreased significantly from 11.5 to 10.9 downstream and alkalinity, from 3.2 to 1.9 meq/l (Table 1). Ca, K, Mg, Si, SO4 and Sr decreased slightly downstream; Ba, Br, Cl and Na increased slightly. In the last sample, F increased by a factor 5 (Fig. 3). The complicated changes in the chemical composition suggest multiple dissolution/precipitation reactions and possible input from a second spring. The DIC increased from 0.22 mm to 0.83 mm over the first 15 m but decreased again to 0.65 mm at 30 m. Precipitation of carbonate minerals consumes DIC. It is likely that magnesium carbonates formed because the Mg concentration in Spring #2 decreased with distance downstream and was more than 2 orders higher than the Mg concentration in Spring #1. Spring #2 water fell over a series of short drops so the turbulence increased the CO2 influx much more than at the calm water–air interface at site OM01 (Fig. 1). CO2 uptake decreased the pH. Rapid influx promoted fast carbonate mineral precipitation, where the rate was limited by nucleation and growth within the water body rather than by CO2 absorption only. Alternatively, the presence of a second spring could discharge water with lower pH and higher DIC into the stream bed, which would promote precipitation.

The water from both Springs #1 and #2 contains several elements which, at high concentration, pose a threat to water quality (Table 2). However, the concentrations of all of these elements are under the WHO threshold limits for clean drinking water (WHO, 2008). Many of the potentially toxic elements in these waters are likely to be scavenged by the precipitating travertine phases (Olsson et al., 2014).

Table 2.

The element concentration of the solid travertine and the spring waters from the two sites. The travertine concentrations were normalized to one mole of CaCO3. Concentration of the elements in the travertine and spring water which potentially pose a threat to the quality of groundwater supplies are included and the WHO standard for drinking water is shown for comparison (WHO, 2008). All concentrations in the spring waters are below the recommend limits of WHO. The precipitating travertine scavenges all toxic elements except Hg.

SpringAsBBaCdCrCuFHgMnMoNiPb
Spring waters (unit: nm)
#1BDLBDL19.2BDLBDL115.89 × 10−30.2323.68BDL7.990.196
#2BDL47000.442BDL0.819BDL1.91 × 10−34.84 × 10−20.88110.19.468.06 × 10−2
WHO threshold1.3 × 1024.6 × 1045.1 × 103279.6 × 1023.1 × 1047.9 × 104307.3 × 1037.3 × 1021.2 × 10348
Solid travertine (unit: mmol/mol of CaCO3)
#17.7 × 10−5BDL8.78 × 10−33.1 × 10−62.1 × 10−43.7 × 10−4NABDL9.51 × 10−3BDL2.61 × 10−36.9 × 10−5
#24.8 × 10−5BDL4.14 × 10−24.0 × 10−62.1 × 10−42.7 × 10−4NABDL1.77 × 10−28.6 × 10−61.33 × 10−39.6 × 10−6
SpringAsBBaCdCrCuFHgMnMoNiPb
Spring waters (unit: nm)
#1BDLBDL19.2BDLBDL115.89 × 10−30.2323.68BDL7.990.196
#2BDL47000.442BDL0.819BDL1.91 × 10−34.84 × 10−20.88110.19.468.06 × 10−2
WHO threshold1.3 × 1024.6 × 1045.1 × 103279.6 × 1023.1 × 1047.9 × 104307.3 × 1037.3 × 1021.2 × 10348
Solid travertine (unit: mmol/mol of CaCO3)
#17.7 × 10−5BDL8.78 × 10−33.1 × 10−62.1 × 10−43.7 × 10−4NABDL9.51 × 10−3BDL2.61 × 10−36.9 × 10−5
#24.8 × 10−5BDL4.14 × 10−24.0 × 10−62.1 × 10−42.7 × 10−4NABDL1.77 × 10−28.6 × 10−61.33 × 10−39.6 × 10−6

BDL, below detection limit. The detection limits are (mm): As, 4 × 10−6; B, 9 × 10−4; Cd, 2 × 10−8; Cr, 2 × 10−7; and Mo, 5 × 10−7.

Table 3.

The element concentration of the solid travertine and the spring waters from sites. The travertine concentrations were normalized to one mole of CaCO3. Small amounts of other elements were detected in the travertine, including Ce, Dy, Er, Eu, Gd, Ho, La, Nd, Pr, Sm, Tb, Tm and Yb.

CaAlCoFeKMgNaPSiSrTiVZn
Spring waters (unit: mm)
#12.151.6 × 10−41 × 10−74.7 × 10−50.2551.72 × 10−39.57.2 × 10−50.004070.001155 × 10−71.4 × 10−64.8E-05
#20.03148.6 × 10−63 × 10−7BDL0.08141.523.8BDL0.1631.4 × 10−5BDL3.0 × 10−6BDL
Solid travertine (unit: mmol/mol of CaCO3)
#10.03542.5 × 10−41.78 × 10−20.2860.964260.02180.1460.4951.9 × 10−41.7 × 10−46.11 × 10−3
#20.01572.6 × 10−41.04 × 10−2BDL11.35.80.01280.8502.261.9 × 10−52.8 × 10−58.85 × 10−3
CaAlCoFeKMgNaPSiSrTiVZn
Spring waters (unit: mm)
#12.151.6 × 10−41 × 10−74.7 × 10−50.2551.72 × 10−39.57.2 × 10−50.004070.001155 × 10−71.4 × 10−64.8E-05
#20.03148.6 × 10−63 × 10−7BDL0.08141.523.8BDL0.1631.4 × 10−5BDL3.0 × 10−6BDL
Solid travertine (unit: mmol/mol of CaCO3)
#10.03542.5 × 10−41.78 × 10−20.2860.964260.02180.1460.4951.9 × 10−41.7 × 10−46.11 × 10−3
#20.01572.6 × 10−41.04 × 10−2BDL11.35.80.01280.8502.261.9 × 10−52.8 × 10−58.85 × 10−3

BDL, below detection limit. The detection limits are (mm): Ce, 4 × 10−8; Dy, 3 × 10−8; Er, 3 × 10−8; Eu, 3 × 10−8; Fe, 7 × 10−6; Gd, 3 × 10−8; Ho, 3 × 10−8; La, 4 × 10−8; Lu, 3 × 10−8; Nd, 3 × 10−8; P, 2 × 10−5; Pr, 4 × 10−8; Sm, 3 × 10−8; Tb, 3 × 10−8; Th, 9 × 10−8; Ti, 2 × 10−8; Tm, 2 × 10−8; and Yb, 3 × 10−8.

Fig. 2.

Element concentrations in the stream water measured with distance downstream in Spring #1, including the concentration of (a) the alkali and alkali earth metals, (b) transition metals, (c) other metals of interest, and (d) anions. The dissolved inorganic carbon was estimated using PHREEQC. Most concentrations remain constant with distance downstream except for Mg, Fe, Ni and Zn. Uncertainty is generally within the dimension of the symbols. No uncertainty limits are available for Ti.

Fig. 2.

Element concentrations in the stream water measured with distance downstream in Spring #1, including the concentration of (a) the alkali and alkali earth metals, (b) transition metals, (c) other metals of interest, and (d) anions. The dissolved inorganic carbon was estimated using PHREEQC. Most concentrations remain constant with distance downstream except for Mg, Fe, Ni and Zn. Uncertainty is generally within the dimension of the symbols. No uncertainty limits are available for Ti.

Fig. 3.

Element concentrations measured as a function of distance downstream in Spring #2, including the concentration of (a) the alkali metals, alkali earth metals and Si, and (b) anions. The dissolved inorganic carbon was estimated using PHREEQC. Uncertainties are within the dimension of the symbols.

Fig. 3.

Element concentrations measured as a function of distance downstream in Spring #2, including the concentration of (a) the alkali metals, alkali earth metals and Si, and (b) anions. The dissolved inorganic carbon was estimated using PHREEQC. Uncertainties are within the dimension of the symbols.

The travertine deposits

The travertine deposits in the alkaline springs were analysed using XRD and SEM/EDXS to identify the mineral composition and the morphology of the samples collected. The precipitates of the springs consist mainly of aragonite and calcite. The XRD patterns of the samples from Spring #1 also showed traces of dypingite (Mg5(CO3)4(OH)2·5H2O), antigorite ((Mg,Fe)3Si2O5(OH)4) and halite (NaCl). We were not able to detect amorphous material. With the exception of halite, these precipitates are typical of the hyperalkaline springs of the Oman ophiolites (Neal and Stanger, 1984a; Chavagnac et al., 2013a). The presence of halite can be explained by evaporation of residual water when the samples were freeze dried. The SEM images of the thin-surface precipitate show rhombohedral calcite in a network of aragonite crystals (Figs 1a, 4a), whereas the precipitate from the bottom of Spring #1 consisted of dypingite rosettes with bouquets of aragonite (Fig. 4b). The XRD patterns showed that the travertine samples from Spring #2 were predominantly of aragonite. The SEM images showed small networks of rosettes growing on aragonite needles (Fig. 4c,d). The rosettes were rich in magnesium, indicating either dypingite or brucite (Power et al., 2007; Chavagnac et al., 2013a).

The chemical composition of the carbonate minerals in the travertine samples was quantified by flash dissolving the samples and analysing the chemical components released to solution. The average composition for travertine from both springs was determined and normalized to one mole of dissolved CaCO3, assuming all Ca ions originated from pure CaCO3 (Table 2 and 3). The travertine contains a wide range of elements, including rare earth elements (REE) and the toxic elements: As, Ba, Cd, Cr, Cu, Mn, Mo, Ni and Pb. Some elements were below detection limits in the spring water but were found in the travertine which indicates that these trace elements have, to a significant extent, been taken up by the precipitates. Elements observed in the spring waters were also found in the travertine, except for B and Hg. Thus, most toxic elements released from an engineered CO2 injection project would be sequestered by secondary minerals but some, namely B and Hg, would not.

Mineral saturation estimated using PHREEQC

Numerous carbonate minerals can form in the alkaline spring waters of Oman. We investigated which phases are supersaturated and evaluated if they would be likely to precipitate while the spring water equilibrated with atmospheric CO2. Knowing which phases are present in the travertine is important because trace element coprecipitation and adsorption are controlled by the composition and structure of the sorbing phase. We used the PHREEQC geochemical speciation model to determine the saturation states. In the model, CO2 was added to the spring water in small steps and the saturation states were calculated until equilibrium with atmospheric CO2 was reached (SI = −3.4). No minerals were allowed to precipitate. For the model, we used data from Sample OM01 (water from Spring #1) and OM06 (from Spring #2), including temperature, chemical composition, pH and alkalinity. The saturation index (SI) maximum at pH 10.6, for the pure carbonate phases, corresponds to the point at which the CO32 concentration peaks. At lower pH, the HCO3 ion dominates.

PHREEQC predicted that the supersaturated phases in Spring #1 would be calcite (hexagonal CaCO3), aragonite (orthorhombic CaCO3) and dolomite (CaMg(CO3)2) and in Spring #2, artinite (Mg2(CO3)(OH)2·3H2O), dolomite, huntite (Mg3Ca(CO3)4), magnesite (MgCO3), hydromagnesite (Mg5(CO3)4(OH)2·4H2O), calcite and aragonite (Fig. 5). Lansfordite (MgCO3·5H2O, not shown) and nesquehonite (MgCO3·3H2O) were predicted to be undersaturated. Inorganic synthesis of dolomite, magnesite and huntite at atmospheric pressure and temperature below 40°C has never been reported (Deelman, 1999; Saldi et al., 2009; Deelman, 2011; dos Anjos et al., 2011) so they are not expected to precipitate in the Oman spring waters. Hydrous magnesium carbonate phases, such as artinite, hydromagnesite and dypingite (Mg5(CO3)4(OH)2·5H2O), are far more likely to form (Hsu, 1967; Botha and Strydom, 2001; Hopkinson et al., 2008; Hänchen et al., 2008). No thermodynamic data are available for predicting the saturation state of dypingite but precipitation of both dypingite and hydromagnesite can be mediated by microorganisms (Ming and Franklin, 1985; Braithwaite and Zedef, 1996; Power et al., 2007; Cheng and Li, 2009; Power et al., 2009). The hydrous magnesium carbonate minerals are relatively stable and form more readily than magnesite at temperatures of <80°C and they have therefore been suggested as alternative reaction products for carbonate storage (Ballirano et al., 2013). The magnesium concentration in the outlet spring waters is very low as a result of serpentinization. In an engineered CO2 injection scenario, serpentinization is expected to be minimal and therefore the relative concentration of Mg would be much higher, i.e. ~50 times higher than the Ca concentration (Boudier and Coleman, 1981). In this scenario, magnesium phases would probably dominate the secondary phases and these could complement the scavenging of contaminants above that observed in this study. More research is needed to confirm this assessment.

Fig. 4.

SEM images of the precipitates from Spring #1 (a,b) and Spring #2 (c,d). (a) The carbonate surface film from Fig. 1a, sample OM01: rhombohedral calcite in a network of aragonite needles. (b) Sample OM03, dypingite rosettes (black arrow) surrounded by acicular aragonite; (c) Sample OM08 from Fig. 1b, crystals of aragonite; (d) close-up of the same area showing a secondary phase (black arrows) growing from or nucleating on the aragonite crystals. These crystals are probably dypingite. Samples OM03 and OM08 were collected from the bottom of the springs.

Fig. 4.

SEM images of the precipitates from Spring #1 (a,b) and Spring #2 (c,d). (a) The carbonate surface film from Fig. 1a, sample OM01: rhombohedral calcite in a network of aragonite needles. (b) Sample OM03, dypingite rosettes (black arrow) surrounded by acicular aragonite; (c) Sample OM08 from Fig. 1b, crystals of aragonite; (d) close-up of the same area showing a secondary phase (black arrows) growing from or nucleating on the aragonite crystals. These crystals are probably dypingite. Samples OM03 and OM08 were collected from the bottom of the springs.

The spring waters from the Semail Ophiolite are highly alkaline, with pH between 11.6 and 10.9. The most prominent cations are Ca, K, Na, Mg and Si; anions are Br, Cl, F and SO4. Within uncertainty, the concentrations remain constant downstream for Spring #1 but complicated changes in the chemical composition for Spring #2 suggest dissolution and precipitation reactions and probably input from a second spring.

Fig. 5.

The saturation state of selected carbonate minerals as a function of pH (addition of CO2) with respect to the water from (a) Spring #1 at Sampling site OM01, and (b) Spring #2 at Sampling site OM06. Secondary minerals were not allowed to precipitate in the model. The reaction progress is from right to left in the figures. Before the addition of CO2, the pH was 11.6 and temperature was 30.6°C for Spring #1, and 11.5 and 20.3°C for Spring #2. The logarithm of the initial partial pressure of CO2 was −9.3 bar for Spring #1 water and −8.3 bar in Spring #2. The logarithm of the final CO2 pressure after the solution reached equilibrium with the Earth's atmosphere was −3.4 bar or 395 ppm CO2.

Fig. 5.

The saturation state of selected carbonate minerals as a function of pH (addition of CO2) with respect to the water from (a) Spring #1 at Sampling site OM01, and (b) Spring #2 at Sampling site OM06. Secondary minerals were not allowed to precipitate in the model. The reaction progress is from right to left in the figures. Before the addition of CO2, the pH was 11.6 and temperature was 30.6°C for Spring #1, and 11.5 and 20.3°C for Spring #2. The logarithm of the initial partial pressure of CO2 was −9.3 bar for Spring #1 water and −8.3 bar in Spring #2. The logarithm of the final CO2 pressure after the solution reached equilibrium with the Earth's atmosphere was −3.4 bar or 395 ppm CO2.

Travertine samples from the alkaline springs consist of aragonite often mixed with calcite and traces of dypingite and antigorite. Other phases that were supersaturated in the spring waters, and predicted to precipitate at atmospheric conditions, were artinite and hydromagnesite. The most common alkali, alkali earth and transition metals were detected in the solid travertine samples along with REE and toxic elements such as As, Ba, Cd, Cr, Cu, Mn, Mo, Ni and Pb. This suggests active and effective scavenging by precipitation of the carbonate minerals, thus limited mobility of contaminants. However, boron and mercury were not removed from the water. If they were present at increased concentrations they would pose a health risk.

Our study shows that several carbonate minerals precipitate readily in groundwater from the Semail Ophiolite when CO2 dissolves in it. Aragonite and calcite dominate the precipitates and these phases are well known to incorporate trace elements. Thus we can conclude that toxic metals that are released during dissolution of the ophiolite are taken up by the travertine. From this we can draw analogies with likely scenarios for geological CO2 sequestration. Although the reaction conditions in an engineered CO2 injection system would be different from the conditions in our study, in both cases, Ca, Mg, Fe, CO2 and trace metals would be present and where carbonate minerals such as calcite reach supersaturation and precipitate, we would expect the precipitates to scavenge the trace elements. Thus we can conclude that if toxic trace metals are released by dissolution of mafic or ultramafic rocks during CO2 injection, they would probably be sequestered by the carbonate minerals or other secondary phases (such as Fe oxides and clays) that form as the waters approach equilibrium.

Thus, this ophiolite system serves as a natural analogue, providing insight into what could be expected from engineered injection of CO2 into porous peridotite and offers an example scenario of what might happen if CO2 leaks from an injection site. Such studies of natural analogues are useful for providing information about natural processes that laboratory studies of ideal systems cannot. Studying natural systems gives information about how elements or products are favoured or disfavoured in complicated competitive reactions.

The authors thank Eydís Salome Eiríksdóttir, Helene Almind and Iwona Gałeczka for technical help. For assistance in the field and transport of the samples, Pablo García del Real and Amelia Paukert are acknowledged. The authors are also grateful to Christophe Monnin and two anonymous reviewers for their comments during review. The work was funded by the Nordic Volcanological Institute (NORDVULK), the Institute of Earth Sciences, Reykjavík, Iceland; the NanoGeoScience Group, Nano-Science Center, Department of Chemistry, Copenhagen, Denmark; a travel grant for the IODP/ICDP Workshop on “Geological carbon capture & storage in mafic and ultramafic rocks” sponsored by the ESF Magellan Workshop Series; and by the European Commission Framework 7, through the CarbFix Project, Grant Agreement No: FP7 283148.

This paper is published as part of a special issue in Mineralogical Magazine, Vol. 78(6), 2014 entitled ‘Mineral–fluid interactions: scaling, surface reactivity and natural systems’.
1.
Aiuppa
A.
Allard
P.
D'Alessandro
W.
Michel
A.
Parello
F.
Treuil
M.
Valenza
M.
(
2000
)
Mobility and fluxes of major, minor and trace metals during basalt weathering and groundwater transport at Mt. Etna volcano (Sicily)
.
Geochimica et Cosmochimica Acta
 ,
64
,
1827
1841
.
2.
Alfredsson
H.A.
Hardarson
B.S.
Franzson
H.
Gislason
S.R.
(
2008
)
CO2 sequestration in basaltic rock at the Hellisheidi site in SW Iceland: stratigraphy and chemical composition of the rocks at the injection site
.
Mineralogical Magazine
 ,
72
,
1
5
.
3.
Alfredsson
H.A.
Oelkers
E.H.
Hardarsson
B.S.
Franzson
H.
Gunnlaugsson
E.
Gislason
S.R.
(
2013
)
The geology and water chemistry of the Hellisheidi, SW-Iceland carbon storage site
.
International Journal of Greenhouse Gas Control
 ,
12
,
399
418
.
4.
Allison
J.D.
Brown
D.S.
Novo-Gradac
K.J.
(
1991
)
MINTEQA2/PRODEFA2, a geochemical assessment model for environmental systems: Version 3.0 User's Manual
 .
EPA/600/3-91/021
,
US Environmental Protection Agency
,
Athens, Georgia, USA
.
5.
Ballirano
P.
De Vito
C.
Mignardi
S.
Ferrini
V.
(
2013
)
Phase transitions in the Mg-CO2-H2O system and the thermal decomposition of dypingite, Mg5(CO3)4(OH)2·5H2O: implications for geosequestration of carbon dioxide
.
Chemical Geology
 ,
340
,
59
67
.
6.
Barnes
I.
Oneil
J.R.
Trescases
J.J.
(
1978
)
Present day serpentinization in New Caledonia, Oman and Yugoslavia
.
Geochimica et Cosmochimica Acta
 ,
42
,
144
145
.
7.
Beinlich
A.
Austrheim
H.
Glodny
J.
Erambert
M.
Andersen
T.B.
(
2010
)
CO2 sequestration and extreme Mg depletion in serpentinized peridotite clasts from the Devonian Solund basin, SW Norway
.
Geochimica et Cosmochimica Acta
 ,
74
,
6935
6964
.
8.
Beinlich
A.
Plumper
O.
Hovelmann
J.
Austrheim
H.
Jamtveit
B.
(
2012
)
Massive serpentinite carbonation at Linnajavri, N Norway
.
Terra Nova
 ,
24
,
446
455
.
9.
Botha
A.
Strydom
C.A.
(
2001
)
Preparation of a magnesium hydroxy carbonate from magnesium hydroxide
.
Hydrometallurgy
 ,
62
,
175
183
.
10.
Boudier
F.
Coleman
R.G.
(
1981
)
Cross section through the peridotite in the Samail Ophiolite, southeastern Oman Mountains
.
Journal of Geophysical Research: Solid Earth
 ,
86
,
2573
2592
.
11.
Braithwaite
C.J.R.
Zedef
V.
(
1996
)
Hydromagnesite stromatolites and sediments in an alkaline lake, Salda Golu, Turkey
.
Journal of Sedimentary Research
 ,
66
,
991
1002
.
12.
Chavagnac
V.
Ceuleneer
G.
Monnin
C.
Lansac
B.
Hoareau
G.
Boulart
C.
(
2013a
)
Mineralogical assemblages forming at hyperalkaline warm springs hosted on ultramafic rocks: a case study of Oman and Ligurian ophiolites
.
Geochemistry, Geophysics, Geosystems
 ,
14
,
1
22
.
13.
Chavagnac
V.
Monnin
C.
Ceuleneer
G.
Boulart
C.
Hoareau
G.
(
2013b
)
Characterization of hyperalkaline fluids produced by low-temperature serpentinization of mantle peridotites in the Oman and Ligurian ophiolites
.
Geochemistry, Geophysics, Geosystems
 ,
14
,
2496
2522
.
14.
Cheng
W.
Li
Z.
(
2009
)
Precipitation of nesquehonite from homogeneous supersaturated solutions
.
Crystal Research Technology
 ,
44
,
937
947
.
15.
Deelman
J.C.
(
1999
)
Low-temperature nucleation of magnesite and dolomite
.
Neues Jahrbuch für Mineralogie Monatshefte
 ,
289
302
.
16.
Deelman
J.C.
(
2011
)
Low-temperature formation of dolomite and magnesite
 .
Compact Disc Publications
,
Eindhoven, The Netherlands
.
17.
dos Anjos
A.P.A.
Sifeddine
A.
Sanders
C.J.
Patchineelam
S.R.
(
2011
)
Synthesis of magnesite at low temperature
.
Carbonates and Evaporites
 ,
26
,
213
215
.
18.
Flaathen
T.K.
Gislason
S.R.
Oelkers
E.H.
Sveinbjornsdottir
A.E.
(
2009
)
Chemical evolution of the Mt. Hekla, Iceland, groundwaters: a natural analogue for CO2 sequestration in basaltic rocks
.
Applied Geochemistry
 ,
24
,
463
474
.
19.
Galeczka
I.
Wolff-Boenisch
D.
Gislason
S.R.
(
2013
)
Experimental studies of basalt-H2O-CO2 interaction with a high pressure column flow reactor: the mobility of metals
.
Energy Procedia
 ,
37
,
5823
5833
.
20.
Galeczka
I.
Wolff-Boenisch
D.
Oelkers
E.H.
Gislason
S.R.
(
2014
)
An experimental study of basaltic glass-H2O-CO2 interaction at 22 and 50°C: implications for subsurface storage of CO2
.
Geochimica et Cosmochimica Acta
 ,
126
,
123
145
.
21.
Gislason
S.R.
Oelkers
E.H.
(
2014
)
Carbon storage in basalt
.
Science
 ,
344
,
373
374
.
22.
Gislason
S.R.
Arnorsson
S.
Armannsson
H.
(
1996
)
Chemical weathering of basalt in southwest Iceland: effects of runoff, age of rocks and vegetative/glacial cover
.
American Journal of Science
 ,
296
,
837
907
.
23.
Gislason
S.R.
Wolff-Boenisch
D.
Stefansson
A.
Oelkers
E.H.
Gunnlaugsson
E.
Sigurdardottir
H.
Sigfusson
B.
Broecker
W.S.
Matter
J.M.
Stute
M.
Axelsson
G.
Fridriksson
T.
(
2010
)
Mineral sequestration of carbon dioxide in basalt: a pre-injection overview of the CarbFix project
.
International Journal of Greenhouse Gas Control
 ,
4
,
537
545
.
24.
Gran
G.
(
1952
)
Determination of the equivalence point in potentiometric titrations. Part II
.
Analyst
 ,
77
,
661
671
.
25.
Hopkinson
L.
Rutt
K.
Cressey
G.
(
2008
)
The transformation of nesquehonite to hydromagnesite in the system CaO-MgO-H2O-CO2: an experimental FT-Raman spectroscopic study
.
Journal of Geology
 ,
116
,
387
400
.
26.
Hsu
K.J.
(
1967
)
Chemistry of Dolomite Formation
. Pp.
169
191
in:
Carbonate Rocks – Physical and Chemical Aspects
  (
Chilingar
G.V.
Bissell
H.J.
Fairbridge
R.W.
, editors).
Developments in Sedimentology
,
Elsevier
,
Amsterdam
.
27.
Hänchen
M.
Prigiobbe
V.
Baciocchi
R.
Mazzotti
M.
(
2008
)
Precipitation in the Mg-carbonate system – effects of temperature and CO2 pressure
.
Chemical Engineering Science
 ,
63
,
1012
1028
.
28.
Jähne
B.
Münnich
K.O.
Bösinger
R.
Dutzi
A.
Huber
W.
Libner
P.
(
1987
)
On the parameters influencing air-water gas exchange
.
Journal of Geophysical Research C: Oceans
 ,
92
,
1937
1949
.
29.
Kelemen
P.B.
Matter
J.
(
2008
)
In situ carbonation of peridotite for CO2 storage
.
Proceedings of the National Academy of Science, USA
 ,
105
,
17295
17300
.
30.
Krauskopf
K.B.
Bird
D.K.
(
1994
)
Introduction to Geochemistry
 .
McGraw-Hill College
,
New York
.
31.
Lackner
K.S.
Wendt
C.H.
Butt
D.P.
Joyce
E.L.
Sharp
D.H.
(
1995
)
Carbon dioxide disposal in carbonate minerals
.
Energy
 ,
20
,
1153
1170
.
32.
Mackenzie
F.T.
Andersson
A.J.
(
2013
)
The marine carbon system and ocean acidification during Phanerozoic time
.
Geochemical Perspectives
 ,
2
,
1
227
.
33.
Matter
J.M.
Kelemen
P.B.
(
2009
)
Permanent storage of carbon dioxide in geological reservoirs by mineral carbonation
.
Nature Geoscience
 ,
2
,
837
841
.
34.
McGrail
B.P.
Spane
F.A.
Sullivan
E.C.
Bacon
D.H.
Hund
G.
(
2011
)
The Wallula basalt sequestration pilot project
.
Energy Procedia
 ,
4
,
5653
5660
.
35.
Ming
D.W.
Franklin
W.T.
(
1985
)
Synthesis and characterization of lansfordite and nesquehonite
.
Soil Science Society of America Journal
 ,
49
,
1303
1308
.
36.
Neal
C.
Stanger
G.
(
1984a
)
Calcium and magnesium-hydroxide precipitation from alkaline groundwaters in Oman, and their significance to the process of serpentinization
.
Mineralogical Magazine
 ,
48
,
237
241
.
37.
Neal
C.
Stanger
G.
(
1984b
)
Past and present serpentinisation of ultramafic rocks; an example from the Semail ophiolite nappe of Northern Oman
. Pp.
249
275
in:
The Chemistry of Weathering
  (
Drever
J.I.
, editor).
D. Reidel
,
Dordrecht, The Netherlands
.
38.
Nicolas
A.
Boudier
E.
Ildefonse
B.
Ball
E.
(
2000
)
Accretion of Oman and United Arab Emirates ophiolite – discussion of a new structural map
.
Marine Geophysical Research
 ,
21
,
147
179
.
39.
Oelkers
E.H.
Gislason
S.R.
Matter
J.
(
2008
)
Mineral carbonation of CO2
.
Elements
 ,
4
,
333
337
.
40.
Olsson
J.
Bovet
N.
Makovicky
E.
Bechgaard
K.
Balogh
Z.
Stipp
S.L.S.
(
2012
)
Olivine reactivity with CO2 and H2O on a microscale: implications for carbon sequestration
.
Geochimica et Cosmochimica Acta
 ,
77
,
86
97
.
41.
Olsson
J.
Stipp
S.L.S.
Dalby
K.N.
Gislason
S.R.
(
2013
)
Rapid release of metal salts and nutrients from the 2011 Grímsvötn, Iceland volcanic ash
.
Geochimica et Cosmochimica Acta
 ,
123
,
134
149
.
42.
Olsson
J.
Stipp
S.L.S.
Makovicky
E.
Gislason
S.R.
(
2014
)
Metal scavenging by calcium carbonate at the Eyjafjallajökull volcano: a carbon capture and storage analogue
.
Chemical Geology
 ,
384
,
135
148
.
43.
Paukert
A.N.
Matter
J.M.
Kelemen
P.B.
Shock
E.L.
Havig
J.R.
(
2012
)
Reaction path modeling of enhanced in situ CO2 mineralization for carbon sequestration in the peridotite of the Samail Ophiolite, Sultanate of Oman
.
Chemical Geology
 ,
330–331
,
86
100
.
44.
Power
I.M.
Wilson
S.A.
Thom
J.M.
Dipple
G.M.
Southam
G.
(
2007
)
Biologically induced mineralization of dypingite by cyanobacteria from an alkaline wetland near Atlin, British Columbia, Canada
.
Geochemical Transactions
 ,
8
,
13
13
.
45.
Power
I.M.
Wilson
S.A.
Thom
J.M.
Dipple
G.M.
Gabites
J.E.
Southam
G.
(
2009
)
The hydromagnesite playas of Atlin, British Columbia, Canada: a biogeochemical model for CO2 sequestration
.
Chemical Geology
 ,
260
,
286
300
.
46.
Saldi
G.D.
Jordan
G.
Schott
J.
Oelkers
E.H.
(
2009
)
Magnesite growth rates as a function of temperature and saturation state
.
Geochimica et Cosmochimica Acta
 ,
73
,
5646
5657
.
47.
Stumm
W.
(
1992
)
Chemistry of the Solid–Water Interface: Processes at the Mineral–Water and Particle–Water Interface in Natural Systems
 .
Wiley
,
New York
.
48.
Stumm
W.
Morgan
J. J.
(
1981
)
Aquatic Chemistry: an Introduction emphasizing Chemical Equilibria in Natural Waters
 .
Wiley
,
New York
.
49.
Turekian
K.K.
Wedepohl
K.H.
(
1961
)
Distribution of the elements in some major units of the earth's crust
.
Geological Society of America Bulletin
 ,
72
,
175
192
.
50.
WHO
(
2008
)
Guidelines for Drinking-Water Quality
 .
World Health Organization
,
Geneva
.
51.
Wolery
T.J.
(
1992
)
EQ3/6, A Software Package for Geochemical Modeling of Aqueous Systems
 .
Lawrence Livermore National Laboratory Report
,
California, USA
,
UCRL-MA-110662 PT I
.
P1Freely available online through the publisher-supported open access option.