Control of Water Chemistry in Alkaline Lakes: Solubility of

Article Cite This: ACS Earth Space Chem. XXXX, XXX, XXX−XXX

Control of Water Chemistry in Alkaline Lakes: Solubility of Monohydrocalcite and Amorphous Magnesium Carbonate in CaCl2−MgCl2−Na2CO3 Solutions Keisuke Fukushi*,† and Haruna Matsumiya‡ †

Institute of Nature and Environmental Technology, Kanazawa University, Kakuma, Kanazawa, Ishikawa 920-1192 Japan School of Natural Systems, College of Science and Engineering, Kanazawa University, Kakuma, Kanazawa, Ishikawa 920-1192 Japan

‡

Downloaded via NORTHERN KENTUCKY UNIV on June 23, 2018 at 10:40:38 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

S Supporting Information *

ABSTRACT: Alkaline lakes play an important role in the global carbon cycle. Nevertheless, little is known about the chemical processes related to sequestration of CO2 in these lakes. Our earlier study demonstrated that the formation of monohydrocalcite (MHC), a hydrous carbonate mineral frequently found in alkaline lakes, requires coexistence of amorphous Mg carbonate (AMC). It is therefore possible that MHC and AMC control the water chemistries of alkaline lakes. To test this hypothesis, we measured the solubilities of MHC and AMC in Na2CO3−MgCl2− CaCl2 solutions and compared them with the water chemistries of alkaline lakes. Results showed that the solubility of MHC is independent of the Mg contents of the system. The solubility product of AMC was almost 2 orders of magnitude higher than that of MHC. The water chemistries obtained from the alkaline saline lakes with pH > 9 around the world closely approximated saturation with respect to both MHC and AMC. MHC, which is a metastable phase, transforms quickly to aragonite or calcite. Under conditions in which MHC controls the solution chemistries, dissolved inorganic carbon is irreversibly and spontaneously isolated via formation of aragonite or calcite in the lakes. Results of this study suggest that alkaline lakes are natural devices for effective CO2 sequestration. KEYWORDS: alkaline lakes, pH, dissolved inorganic carbon, monohydrocalcite, amorphous magnesium carbonate, metastable phase

■

INTRODUCTION Saline lakes, which comprise almost half of inland waters by volume, play an important role in the global carbon cycle.1−4 The CO2 fluxes into saline lakes from the atmosphere reportedly depend strongly on the pH of the lakes.1−4 Finlay et al.4 have monitored the CO2 fluxes and the water chemistries of lakes in the northern Great Plains of Canada and have demonstrated that pH values lower than 9.0 are associated with substantial CO2 evasion, whereas pHs higher than 9.0 are associated with CO2 capture. Saline lakes are typically rich in dissolved inorganic carbon (DIC) rather than dissolved organic carbon.2,5 Chemical uptake of CO2 into solution is enhanced as the pH rises1,4 because the saturation concentration of DIC in water increases exponentially with pH.1,6 These observations suggest that saline, alkaline lakes with pHs greater than 9 (hereinafter, alkaline lakes) play important roles as CO2 sinks. In alkaline lakes, ongoing formation of carbonate minerals is observed frequently.7,8 Alkaline lakes may play a role as natural devices for effective CO2 sequestration because the carbonate minerals in them are formed from DIC (CO32− and HCO3−) derived from atmospheric CO2. Understanding the mechanisms of CO2 sequestration in alkaline lakes requires a quantitative understanding of how carbonate minerals are formed under the © XXXX American Chemical Society

given solution conditions. Calcite and aragonite are common carbonate minerals formed under Earth’s surface conditions.7,8 However, the water of alkaline lakes is always highly supersaturated with respect to calcite and aragonite (Figure S1a, Supporting Information). This supersaturation suggests that anhydrous calcium carbonate may not control the water chemistry of alkaline lakes. Monohydrocalcite (MHC) is a hydrous calcium carbonate with the composition of CaCO3·H2O that is found frequently in the sediments of alkaline lakes.9−11 Nishiyama et al.12 and Fukushi et al.13 have conducted calcium carbonate synthesis experiments with widely different concentrations of Na2CO3, CaCl2, and MgCl2. They observed the formation of vaterite, calcite, aragonite, MHC, and amorphous materials. They found that MHC formed only when the reacted solution was saturated with respect to amorphous Mg carbonates (AMC) and concluded that the formation of MHC requires coexistence of AMC. MHC and AMC are expected to be more soluble than anhydrous carbonates.13,14 We therefore Received: April 11, 2018 Revised: June 5, 2018 Accepted: June 6, 2018

A

DOI: 10.1021/acsearthspacechem.8b00046 ACS Earth Space Chem. XXXX, XXX, XXX−XXX

Article

ACS Earth and Space Chemistry Table 1. Initial Solution Compositions, Duration Times, and log IAPs of MHC in the Experimental Runsa initial composition (mol/L) run run run run

1-1 1-2 2 3

log IAP

CaCl2

MgCl2

Na2CO3

Mg/Ca

0.046 0.044 0.048 0.047

0.015 0.014 0.024 0.047

0.075 0.075 0.075 0.074

0.33 0.31 0.50 1.00

duration 456 648 648 648

h h h h

MHC

AMC1

AMC2

−7.66 −7.70 −7.61 −7.61

−5.3 −5.3 −5.1 −5.1

−5.54 −5.65 −5.62 −5.53

a

AMC1 and AMC2 were obtained from each experimental run (see the text).

electrolyte, ML, the logarithm of the activity coefficient, γi,aq, for the ith ionic aqueous species is given by17,18

hypothesized that the combination of MHC and AMC controls the water chemistries of alkaline lakes. Testing this hypothesis requires accurate solubility data for MHC and AMC. However, the published solubilities are inconsistent for MHC13−15 and are limited for AMC.13 In the present study, we measured the solubilities of MHC and AMC based on longterm precipitation/transformation experiments of carbonate minerals in NaCO3−CaCl2−MgCl2 solutions for up to 650 h under atmospheric CO2 partial pressures. Comparison of the solubilities obtained from the present study with the water chemistries of alkaline lakes revealed that the water chemistries of the alkaline lakes closely approximated saturation with respect to both MHC and AMC.

log γi ,aq =

−A γ zi2,aqI1/2 1 + åBγ I1/2

+ Γγ + bγ I (1)

where zi,aq is the charge of the ith aqueous ion, I is the ionic strength of the solution, å is the ion-size parameter, Aγ and Bγ are the electrostatic Debye−Hückel parameters, and bγ,ML is the extended-term parameter for ML-th electrolyte. The bv and å parameters are 0.064 kg/mol and 3.7 Å for NaCl media.17,18 The Γγ term converts rational activity coefficients to their molality-based counterparts and is given by

■

Γγ = −log(1 + 0.0180153m*)

MATERIALS AND METHODS Solubility Measurements. Temporal evolution of the solution chemistry and the mineralogy from CaCl2−MgCl2− Na2CO3 solutions were monitored in batch experiments under atmospheric CO2 partial pressures. The mixed solutions of CaCl2 and MgCl2 were prepared in 1-L Teflon bottles. A Na2CO3 stock solution was then added to the bottles. Four runs with similar and different solution compositions were conducted in the present study (Table 1). Hereinafter, we use the notation Na, K, Mg, Ca, Cl, SO4, and DIC to express sodium, potassium, magnesium, calcium, chlorine, sulfate, and dissolved inorganic carbon as components. The Ca and DIC concentrations in the initial solutions were similar among all the runs at around 0.045 and 0.075 mol/L, respectively. The Mg concentration range of 0.014−0.047 mol/L (Table 1) caused the Mg/Ca ratio in the system to vary from 0.3 to 1.0. After the addition of Na2CO3, whitish precipitates immediately formed in the vessels. The vessels were kept in an incubator or a temperature-controlled room at 25 °C under atmospheric conditions. The suspensions were stirred continuously using a magnetic stirrer. Aliquots of the suspensions (10−20 mL) were sampled periodically from the vessels. Before collecting the aliquots, we measured the pH of each suspension with a pH meter (HM-21P; DKK-TOA Corp.) that had been calibrated with three buffer solutions. The collected suspensions were filtered through a 0.2 μm membrane under vacuum. The filtrates were acidified with addition of a small volume of concentrated HNO3 for the measurements of Ca and Mg using an inductive coupled plasma-optical emission spectrometer (ICP-OES, ES-710S; Varian Inc.). The solids collected on the filter paper were freeze-dried and analyzed using an X-ray diffractometer (XRD, Ultima IV, Cu Kα, 40 kV, 30 mA; Rigaku Corp.). Speciation analyses of the reacted solutions were conducted using REACT of the Geochemist’s Workbench16 with the thermodynamic database of “thermo.v8.r6+.dat”. The activity coefficient was calculated using the Helgeson, Kirkham, and Flowers17 version of the extended Debye−Hückel equation. The activity model can account for an ionic strength of up to 5 m for NaCl media.17 For a solution containing a 1:1 symmetric

(2)

where m* = ∑i ,aq mi ,aq represents the sum of the molalities (mi,aq) of all the aqueous species in the system. The maximum m* in the experimental solutions was 0.4 m. Therefore, the Γγ term in eq 1 is negligible. Input parameters for speciation analyses were the pH and concentrations of Na, Ca, Mg, and Cl. The Na and Cl concentrations were derived from the initial concentrations of CaCl2, MgCl2, and Na2CO3 (Table 1). The Ca and Mg concentrations were derived from ICP-OES measurements. Our earlier study has shown that the DIC concentrations calculated from the charge balance during speciation analysis are more accurate than those calculated from alkalinity measurements when the contribution of HCO3− and CO32− to total electric charge balance of the solution is greater than 5%.12 In this study, we estimated the DIC concentrations from charge balance calculations because the contributions of HCO3− and CO32− to total electric charge balance always exceeded 5%. Dissolution reactions of calcium or magnesium (Me2+) carbonates with the stoichiometry of MeCO3·nH2O can be written as MeCO3 ·nH 2O = Me 2 + + CO32 − + nH 2O

(3)

The corresponding mass action expression is K sp =

aMe2+aCO32−a Hn 2O aMeCO3·nH 2O

(4)

where ai represents the activity of the ith species. The activities of water and pure solids can be assumed to be unity for the ionic strengths of the solutions in the experiments (I < 0.3). The ionic activity products (IAP) with respect to the calcium carbonates (CaCO3·nH2O) and magnesium carbonates (MgCO3·nH2O) were calculated using the activities of each species from the speciation analyses. Assembling of Water Chemistries in Natural Alkaline Lakes. The water chemistries of alkaline lakes with pH > 9 were assembled from the published literature. Numerous reports in the literature have described the water chemistry of alkaline lakes. Five criteria were used for selection in the B


Article

C

−5.31 not analyzed 250 0.29 9.7

7.2

12

0.099

190

27

31

0.12

−5.47 −7.65 −1.1%

−5.07 not analyzed −5.57 −7.11 8.7% 0.03 660 290 1200 0.19 3.6 110 3400 10.64 2.9

reference

Yang and Williams (2003)21 Yang and Williams (2003)21 Yang (2006)23 −4.06 not analyzed −5.33 −7.12 2200 10.54 2.2

110

6.3

0.19

1200

160

500

0.04

−2.0%

−5.16 aragonite 1600

China (InnerMongolia) Sumujilin-S China (InnerMongolia) Sumujilin-N China (InnerMongolia) Yabulai 11 China (InnerMongolia) Nuoertu

Duikou

China (North Hebei) ErquanjingII China (North Hebei) Kulun China (North Hebei) Qinghai China (Tibet) Huhejilin or China (InnerHuhejaran Mongolia)

17.8

10

1.6

120

6.4

0.31

820

120

380

0.29

−5.37 −6.93 −0.4%

−5.57 not analyzed −4.81 aragonite, calcite −5.45 −5.45 −7.53 −7.07 170 1900 9.23 0.26 10.46 1.9

0.17 9.6

160

0.14 9.3

120

3.9 91

35 4.7

0.30 0.21

170 830

24 130

19 460

0.29 0.05

0.5% 0.5%

−7.45 not analyzed −5.81 −7.62 3.1

0.077

93

3.9

49

0.34

−0.4%

−7.30 not analyzed −5.78 −7.53 8.3

0.18

93

7.7

20

0.34

0.8%

−6.23 not analyzed −5.62 −7.45 −0.8% 0.19 16 7.5 97 0.23 13 100 0.14 9.4

carbonate mineralogy CaCO3-nH2O MgCO3-nH2O HM C.B. Na K Mg Ca Cl SO4 DIC pCO2 mmol/kg mmol/kg mmol/kg mmol/kg mmol/kg mmol/kg mmol/kg matm I pH temp °C area

RESULTS AND DISCUSSION Solution Chemistry and Mineralogy. XRD patterns of the reacted solids showed similar mineralogical changes among the different experimental conditions (Figure 1 and Figure S2, Supporting Information). Figure 2 presents changes of the intensities of the XRD peak of MHC (20.5° at 2θ) and aragonite (26.2° at 2θ) with time. At the beginning of the reactions, small peaks attributable to MHC appeared. Those peak intensities increased with time (stage 1) and then reached a plateau, which they maintained for some time (stage 2). The intensities of the MHC peaks then decreased with time, while peaks attributable to aragonite appeared and increased with time (stage 3). Finally, the peaks from MHC disappeared completely (stage 4). The times for each sequence were not reproducible, even when the compositions of the initial solutions were similar. The initial conditions for runs 1-1 and 1-2 were almost identical. Stage 3 started after 400 h in run 1-2, although it did not occur in run 1-1 (Figure 2a,b). It has been postulated that the Mg in MHC protects MHC from being transformed.12,39 However, no systematic relationship was found between the transformation rate and the Mg/Ca ratio under the present experimental conditions. The pHs were high (>10) at the beginning of the runs and decreased with time (Figure 3a). The decreases of pH with time from the beginning of the runs were attributable mainly to equilibration with atmospheric CO2 as well as the formations of carbonate minerals. The pHs reached almost constant values after ca. 300 h in all the runs. The plateau pH decreased with increasing Mg/Ca in the initial solutions. The pH of run 3 (highest Mg/Ca) after 300 h was lowest among the runs at around 9.0. Changes of Ca concentrations were clearly related to the changes of mineralogy (Figure 3b). The Ca concentrations decreased during stage 1 and became constant during stage 2. At stage 3, the Ca concentrations were similar to the concentrations during stage 2, but the concentrations suddenly decreased and then became constant again during stage 4. Although the pH and the Ca concentrations decreased monotonically, the Mg concentrations initially either increased or did not decrease at the beginning of the runs (Figure 3c).

lake

■

log IAP

Table 2. Solution Chemistries; Estimated Partial Pressures of CO2 (pCO2); Charge Balance (C.B.) of Solution; log IAPs with Respect to CaCO3·nH2O, MgCO3·nH2O, and Hydromagnesite (HM); and Carbonate Mineralogy of Sediments from Alkaline Lakes around the World

present analyses. (1) The concentrations of major components including Na, Mg, Ca, Cl, SO4, and DIC (or alkalinity) are reported in the literature. (2) The water in the lake is Na−Cl type, in which case it was possible to apply the same activity models used in the laboratory experiments of this study. (3) The electric change balances of the solutions after the speciation analyses are in balance to within 10%. When several analytical data were available from a single lake, we used the data with the smallest electric charge balance. (4) The salinities were higher than 1000 mg/kg, which is threshold level to separate freshwater. (5) The ionic strengths were less than 3, in which case the activity of H2O (>0.95) can be approximated by unity. The mass action expression for the dissolution of AMC contains the activity of H2O. However, the n in MgCO3· nH2O was considered to be variable and unknown.13 It should be noted that the Γγ term in eq 1 is negligible if I < 3. We obtained data from 24 lakes from widely different areas of China,19−23 Mongolia,24,25 Russia,26 Turkey,27,28 Italy,29,30 Australia,11,31 and the United States32−38 (Table 2). Their pHs and ionic strengths were 9.02−10.6 and 0.045−2.9, respectively. Some reports did not state the lake water temperature during sample collection, and in those cases we assumed the temperature to be 25 °C for speciation analysis of the water.

Wen and He (1999)19 Wen and He (1999)19 Wen and He (1999)19 Li et al. (2009)20 Yang and Williams (2003);21 Arp et al. (1998)22 Arp et al. (1998)22

ACS Earth and Space Chemistry


Southern Australia

D

United States (Nebraska) United States (Nevada) United States (Nevada) United States (Nevada)

Wilkinson

Walker

Pyramid

Big Soda

United States (California)

Mono

Western Australia Walyungup Western Australia

Richmond

East Basin

temp °C pH

I

9.2

15.1

12.8

26.2

0.31

1.54

0.41

9.45 0.17

9.39 0.08

9.7

10.04 1.3

9.8

130

58

350

1000

1500

92

9.1

0.15

24

720

280

340

640

220

150

67

320

4.2

7.9

193

45

1.8

0.46

15

10

14

2.3

5.6

7.2

6.0

1.3

1.7

23

7

50

5.2

4.4

64

0.59

18

16

31

0.27

0.21

0.12

0.08

0.10

0.32

0.64

0.81

0.15

0.10

0.23

0.21

1.2

0.74

1.2

63

59

180

470

640

100

24

840

230

160

480

130

110

43

240

21

1.7

58

220

140

4.9

1.8

4.6

7

25

130

7.8

28

16

59

38

19

66

250

460

9.8

11

33

63

150

25

67

16

21

27

−6.66

−7.25 −7.61 −7.43

−7.57

−7.19

0.3%

0.042 −1.0% −9.6%

−3.2%

−3.4%

0.27

0.42

0.26

0.19

0.77

0.25

0.37

0.49

2.1

0.61

1.1

−5.80

−7.66 −7.41

−7.55 −7.45 −7.16 −7.35

5.9% −2.6%

−5.7% −4.1% −9.4% −1.8%

−5.70

−5.74

−5.64

−6.13

−5.96

−5.98

−5.38

−5.95

−5.55

−7.08

−2.8%

−5.87

−7.01

−0.4%

0.45

−5.17

−5.67

−7.02

0.2%

0.67

carbonate mineralogy

−6.76 MHC, calcite, aragonite

−7.19 not analyzed

−6.47 not analyzed

−8.80

−8.51

−7.24

−8.31

−5.45

−8.94

reference

Namsaraev et al. (2010)26

Isupov et al. (2012)25

Baatar et al. (2017)24



Cloern et al. (1983)35 Iwama et al. (1997)36 Cooper and Koch (1984);37 Benson et al. (1991)38

Reimer et al. (2009);27 Kazmierczak and Kempe (2003)28 aragonite, Dongarra et al. hydromagnesite, (1983);29,8 calcite, dolomite Aiuppa et al. (2007)30 MHC, dolomite, Last (1992)11 magnesite, hydromagnesite, aragonite, calcite, Mgcalcite not analyzed Coshell et al. (1998)31 aragonite, Coshell et al. hydromagnesite, (1998)31 dolomite calcite, ikaite, Connell and gaylussite Dreiss (1995);32 Bischoff et al. (1993)33 not analyzed Gosselin (1997)34

−6.41 aragonite, Mg-calcite, calcite

−3.33 calcite

−6.97 no carbonates from XRD −7.04 no carbonates from XRD −7.86 no carbonates from XRD −12.16 calcite

CaCO3-nH2O MgCO3-nH2O HM −5.65

0.73

−0.3%

C.B.

log IAP

−7.02

Na K Mg Ca Cl SO4 DIC pCO2 mmol/kg mmol/kg mmol/kg mmol/kg mmol/kg mmol/kg mmol/kg matm

9.1 0.045

9.15 0.85

9

Western 18.0−19.5 9.02 0.40 Mongolia Western 9.06 0.11 Mongolia Western 9.07 0.20 Mongolia Northwestern 9.28 0.23 Mongolia Russia 9.8 0.77 (Sourtheastarn Transbaikalia) Turkey (eastern 20 9.7 0.43 Anatolia)

area

Specchio di Italy (Pantelleria Venere Island)

Van

Khilganta

Shaazgai

Tsegeen

Telmen

Oigon

lake

Table 2. continued

ACS Earth and Space Chemistry Article


Article


Figure 1. XRD patterns of the reacted solids as a function of the reaction times from run 1-2 (a) and run 3 (b).

Figure 2. Changes of intensities of XRD peaks from MHC (20.5° at 2θ: red squares) and aragonite (26.2° at 2θ: blue circles) with time from run 1−1 (a), run 1−2 (b), run 2 (c), and run 3 (d). Stage 1 is the time interval when the MHC peak intensity increased with time. Stage 2 is the time interval when the intensity of the MHC peak reached a plateau. Stage 3 is the time interval when the MHC and aragonite peaks coexisted. Stage 4 is the time interval when the MHC peak disappeared completely and aragonite alone was present.

(Figure 3d and Figure S3, Supporting Information). The IAPs decreased during stage 1 while the MHC was growing but remained constant during stages 2 and 3 before decreasing again and then finally reaching a minimum constant value at

After reaching maximum levels, the Mg concentrations decreased and then remained almost constant. Solubility of MHC. The IAPs of the CaCO3·nH2O were closely related to the mineralogy of the calcium carbonates E


Article


Figure 3. pH (a), concentrations of Ca (b), and Mg (c), and log IAPs with respect to CaCO3·nH2O (d) and MgCO3·nH2O (e) versus reaction times. The stages based on the mineralogy of the calcium carbonates are indicated by the different plotting symbols.

MHC and precipitation of aragonite from the solution.41 The slight changes of the IAPs during the formation of aragonite indicate that the dissolution rates of MHC are markedly higher than the precipitation rates of aragonite.41 It can be inferred that the concentrations of Ca are determined by the solubility of MHC when the more-soluble MHC coexisted with the lesssoluble anhydrous calcium carbonates. Solubility of AMC. In our experiments, the molar Mg/Ca ratio of >0.2 in the solids calculated from the mass balance of the Ca and Mg concentrations before and after the reactions (Table S1) suggests that a substantial amount of Mg must have been present in the solid phases associated with the MHC. Crystalline phases other than MHC were not apparent in the XRD patterns of the reacted solids (Figure 1). Fukushi et al.13 have examined the speciation of Mg associated with MHC using XANES, ab initio calculations, and geochemical modeling. They concluded that most Mg was present as a discrete amorphous phase with a stoichiometry of MgCO3· nH2O. Figure 3e and Figure S4 in the Supporting Information show IAP as a function of MgCO3·nH2O. During the early part of the runs before ca. 100 h, the log IAP values were almost constant at approximately −5.2 ± 0.1 (Figure S4, Supporting Information). Mg concentrations remained high or even increased while pH decreased with time during the interval. The decreases of pH were mainly attributable to equilibration with atmospheric CO2. The decreases of pH resulted in decreases of CO32− activity, because the dissociation constant of HCO3− to CO32− is 10−10.3. According to the mass action law in eq 4, a decrease of CO32− activity would lead to an increase of Mg2+ activity. An increase of Mg2+ activity suggests that the solutions were equilibrated with a certain MgCO3· nH2O phase. Fukushi et al. have also observed that the log IAP with respect to MgCO3·nH2O averages −5.2 after a 24-h reaction time.13 The constant IAPs observed in the beginning of the runs were consistent with the results of a previous study13 and correspond to the solubility of the freshly formed

stage 4. It can be surmised that the constant IAP obtained during stage 2 correspond to the solubility of MHC (Figure S3, Supporting Information) because little mineralogical change was observed during stage 2. The log IAPs were similar among the different Mg/Ca ratios (Table 1 and Figure S3, Supporting Information). The solubility of MHC must vary with both the Mg content of the solid and the Mg concentration in solution if significant amounts of Mg are incorporated into the MHC structure.40 The very small differences of the IAPs associated with the 3-fold differences of Mg concentrations in the present study (Table 1) confirmed that the solid solubility of Mg in the MHC structure is limited, as XANES and ab initio calculations have indicated.13 The averaged log Ksp of MHC from the four runs was −7.65 ± 0.04 (1σ). Hull and Turnbull have measured the solubility of MHC using dissolution experiments of a natural specimen collected from Lake Fellmongery, Australia.14 The log Ksp from their study was −7.60, which closely approximates the value obtained in the present study. However, other studies have reported a log Ksp of −7.2,13,15 which differs from that of the present study. The solubility estimated by Fukushi et al. was obtained after 24-h precipitation experiments.13 The present study demonstrated that a reaction time of 24 h was insufficient to reach equilibrium (Figure 3d). The solubilities reported by Hull and Turnbull14 and those obtained in the present study seem more reasonable than the values from other studies.13,15 The Gibbs free energy of formation of MHC based on the solubility estimated from present study is −1361.6 ± 0.3 kJ/mol. The Gibbs free energy estimated from the present study matches the value in the thermodynamic database of “thermo.v8.r6+.dat” (−1361.6 kJ/mol), which is based on the solubility reported by Hull and Turnbull.14 The IAPs changed little at stage 3 during aragonite formation and growth (Figure 3d and Figure S3, Supporting Information). The IAPs decreased after the MHC disappeared completely during stage 4. The transformation of MHC to aragonite comprises dissolution of F


Article


values of crystalline calcium carbonates are −8.48 (calcite),6 −8.34 (aragonite),6 −7.91 (vaterite),43 −7.65 (MHC), and −6.58 (ikaite).43 The reported values for amorphous calcium carbonate (ACC) are −6.046 and −6.3.47 The observed log IAPs of the lakes are close to or slightly above the solubility of MHC. The Ca concentrations in alkaline lakes are most likely controlled by the formation of MHC. The observed IAPs of the lakes are slightly but systematically higher than the estimated solubility of MHC but significantly lower than that of ikaite. That difference might simply represent disequilibria of MHC formation. Differences between the observed IAPs and the solubility of MHC are most likely related to the different degrees of evaporation and/or rates of formation of MHC in the respective lakes.The presence of MHC in lake sediments has been observed in East Basin and Walker Lake.8,38 However, the formation of MHC was not observed in the many alkaline lakes22,24−26,31,28,30,33 examined in the present study (Table 2). Instead, formation of anhydrous calcium carbonates such as calcite and/or aragonite was observed almost ubiquitously.22,25,26,31,28,33 The present laboratory study demonstrated that the solution chemistry is governed by MHC when MHC and more stable anhydrous calcium carbonates coexist. We surmised that the formation of MHC might have been overlooked in most studies of alkaline lakes. MHC is a metastable phase and quickly transforms to calcite or aragonite.40,48 In saline lakes subject to high rates of evaporation, the precipitation of MHC occurs in water bodies in the lakes. The precipitates formed from the water body in the lakes settle to the lake bottom. MHC can be transformed to anhydrous calcium carbonates such as calcite and aragonite during sedimentation, if the inhibitors of the transformation of MHC such as phosphate49 are scarce in the lakes. In this scenario, the solution chemistry is governed by MHC, but the analyzed sediments had been anhydrous calcium carbonates transformed from MHC. Negative correlation between the log activities of Mg2+ and CO32− with slopes of ca. −1 (Figure 4) also suggest that the Mg concentrations in the alkaline lakes are controlled by a phase with a stoichiometry of MgCO3·nH2O. The average of the log IAPs of the alkaline lakes was −5.7 ± 0.3 (1σ), which is consistent with the log Ksp of AMC2 (−5.59) estimated from the present study. The presence of crystalline Mg carbonates such as HM has been observed in the East Basin Lake,11 Specchio di Venere Lake,30 and Walyungup Lake.31 However, the IAPs of these two lakes with respect to HM were more than 3 orders of magnitude higher than the Ksp of HM (Table 2). Figure 1 shows that XRD cannot detect AMC. Therefore, the presence of AMC in nature must often be overlooked. The solubility of AMC is almost 2 orders of magnitude higher than that of MHC. Because water chemistry is usually governed by the more soluble phases than by the less-soluble phases,6 important parameters such as the DIC concentrations and the pH of alkaline lakes are probably controlled by the presence of AMC. The solution chemistries such as the pH and concentrations of DIC, Ca, and Mg of alkaline lakes are most likely limited by MHC and AMC2. The pH and concentrations of DIC, Ca, and Mg in the alkaline lakes can be predicted from the charge and mass balances of the solution by assuming equilibria with respect to MHC, AMC2, and the partial pressure of CO2. According to Henry’s law, the partial pressure of CO2 is proportional to the activity of CO2(aq),6 which can be calculated from speciation analyses of the lake waters (Table 1). The calculated partial pressures of CO2 in the lakes were used

AMC. The log IAPs decreased gradually with time and again then reached a constant value that averaged −5.59 ± 0.05 (1σ) for the four runs. Reportedly, the aged MHC from high Mg/Ca solutions exhibited a small peak of hydromagnesite (HM) at around 15° 2θ in the XRD pattern,12,40,42 although that peak was not detected in the present study (Figure 1). The HM associated with aged MHC has been regarded as an alteration product of AMC.13 The stoichiometry of HM, Mg5(CO3)4(OH)2·4H2O, is different from that of MgCO3·nH2O. The calculated log IAPs with respect to the stoichiometry of HM were −6.0 to −6.8 under the conditions of the present study (Table S1). The reported log Ksp of HM is −10.29,43 which is considerably lower than the log IAPs with respect to the stoichiometry of HM calculated from the present study. The Mg/Ca ratio in the solids from run 3 exceeded 0.3 (Table S1). This result suggests that the contribution of Mg carbonate to the reaction products (Mg/Mg + Ca) exceeded 20%. However, the XRD patterns showed no presence of crystalline Mg carbonates (Figure 1b). We surmised that the constant log IAP of −5.59 after ∼200 h (Figure S4) was also attributable to amorphous Mg carbonate, which is more stable than freshly prepared AMC. Hereinafter, we refer to AMCs with log Ksp values of −5.2 and −5.59 as AMC1 and AMC2, respectively. Even after the end of the duration of the runs, no significant decrease of log IAP related to AMC2 was observed which suggests that the AMC associated with MHC must be a durable phase, even in natural environments. Controls of Solution Chemistries in Alkaline Lakes. Saline lakes are found in closed basins with limited outflows in arid regions of the world.44,45 The solutes in the surface water or inlet water to terminal lakes are concentrated in the lakes by evaporation.45 The Ca and Mg concentrations in these lakes are always lower than the Na concentrations by several orders of magnitude (Table 2). The Ca and Mg in solution are expected to be removed by precipitation of Ca and Mg carbonates, as simulated in the present laboratory study.The relation between the log activities of Ca2+ and CO32− from alkaline lakes (Figure 4) is characterized by a negative correlation with

Figure 4. Relationship between the log of CO32− activity and the log of Ca2+ activity (black circles) or log of Mg2+ activity (red triangles) from alkaline lakes. Straight black and red lines are the solubilities of MHC and AMC2 estimated from the present laboratory experiments. The dotted lines are errors associated with the solubilities.

a slope of ca. −1. The negative correlation suggests that the Ca concentrations in alkaline lakes are controlled by a phase with a stoichiometry of CaCO3·nH2O. The log IAPs with respect to CaCO3·nH2O of the lakes ranged from −6.93 to −7.66 (Table 2 and Figure S1a, Supporting Information). The log Ksp G


Article


Figure 5. Comparison between the observations and predictions of pH (a), Ca concentration (b), Mg concentration (c), and DIC concentration (d) in alkaline lakes.

MHC and AMC control the water chemistries of alkaline lakes has important implications for CO2 sequestration in alkaline lakes. The main driving force of carbonate precipitation in alkaline lakes is evaporation of the lake water and subsequent accumulation of solutes by drying.44,45 The MHC and AMC precipitate when the lake waters are supersaturated with respect to these phases. The water levels of saline lakes easily fluctuate in response to the climatic changes and hydrological processes such as permafrost degradation.50,51 The increase of water levels must be associated with dilution of the lake water. Dilutions would result in the dissolution of MHC and AMC because of the decrease of their degree of saturation, which suggests that CO2 sequestration by formation of carbonate minerals and restoration of CO2 by dissolution of carbonate minerals are mutually compensated during fluctuations of the lake levels. In contrast, it can be inferred that the contribution of CO2 sequestration always overwhelms the restoration of CO2. Once formed from solution, the MHC spontaneously transforms to stable calcite and aragonite. During dilution of the lake water, the soluble MHC and AMC must be dissolved to keep the lake water saturated with respect to MHC and AMC unless these minerals are consumed completely. The transformed calcite and aragonite hardly dissolve in lake water saturated with these minerals. Therefore, the DIC in the

in the predictions. Figure 5 shows the relationship between the predicted and observed pH and the concentrations of DIC, Ca, and Mg from the examined alkaline lakes. The predicted pH and DIC can reproduce the observations very well. The predicted Mg concentrations can also reproduce the observations reasonably well, except for the lowest Mg concentration (Lake Shaazgai). The estimated IAP with respect to MgCO3·nH2O in this lake is exceptionally low (log IAP = −6.6). The reason for the discrepancy is currently unclear. One possible reason for the overestimation is that the other Mg-bearing minerals may control the Mg concentration in the lake. The predicted Ca concentrations were in most cases underestimations, probably because the lake waters were usually supersaturated with respect to MHC. Their reasonable agreements is consistent with the assumption that the alkaline lake waters are saturated with respect to MHC and AMC2.

■

CONCLUSIONS This study determined the solubilities of MHC and AMC from systematic, long-term precipitation and transformation experiments with Na2CO3−CaCl2−MgCl2 solutions. The obtained solubilities could closely reproduce the water chemistries, including the pH and the concentrations of DIC, Ca, and Mg of alkaline lakes around the world. The fact that metastable H


Article


(7) Chagas, A. A. P.; Webb, G. E.; Burne, R. V.; Southam, G. Modern Lacustrine Microbialites: Towards a Synthesis of Aqueous and Carbonate Geochemistry and Mineralogy. Earth-Sci. Rev. 2016, 162, 338−363. (8) Last, F. M.; Last, W. M. Lacustrine Carbonates of the Northern Great Plains of Canada. Sediment. Geol. 2012, 277−278, 1−31. (9) Fukushi, K.; Munemoto, T.; Sakai, M.; Yagi, S. Monohydrocalcite: A Promising Remediation Material for Hazardous Anions. Sci. Technol. Adv. Mater. 2011, 12 (6), 064702. (10) Last, F. M.; Last, W. M.; Halden, N. M. Carbonate Microbialites and Hardgrounds from Manito Lake, an Alkaline, Hypersaline Lake in the Northern Great Plains of Canada. Sediment. Geol. 2010, 225, 34−49. (11) Last, W. M. Petrology of Modern Carbonate Hardgrounds from East Basin Lake, a Saline Maar Lake, Southern Australia. Sediment. Geol. 1992, 81, 215−229. (12) Nishiyama, R.; Munemoto, T.; Fukushi, K. Formation Condition of Monohydrocalcite from CaCl2-MgCl2-Na2CO3 Solutions. Geochim. Cosmochim. Acta 2013, 100, 217−231. (13) Fukushi, K.; Suzuki, Y.; Kawano, J.; Ohno, T.; Ogawa, M.; Yaji, T.; Takahashi, Y. Speciation of Magnesium in Monohydrocalcite: XANES, Ab Initio and Geochemical Modeling. Geochim. Cosmochim. Acta 2017, 213, 457−474. (14) Hull, H.; Turnbull, A. G. A Thermochemical Study of Monohydrocalcite. Geochim. Cosmochim. Acta 1973, 37 (3), 685−694. (15) Kralj, D.; Brečević, L. Dissolution Kinetics and Solubility of Calcium Carbonate Monohydrate. Colloids Surf., A 1995, 96 (3), 287−293. (16) Bethke, C. M. The Geochemist’s Workbench User’s Guide; University of Illinois: Chicago, 1998. (17) Helgeson, H. C.; Kirkham, D. H.; Flowers, G. C. Theoretical Prediction of the Thermodynamic Behavior of Aqueous Electrolytes at High Pressures and Temperatures: IV. Calculation of Activity Coefficients, Osmotic Coefficients, and Apparent Molal and Standard and Relative Partial Molal Properties to 600°. Am. J. Sci. 1981, 281 (10), 1249−1516. (18) Criscenti, L. J.; Sverjensky, D. A. The Role of Electrolyte Anions (ClO4-, NO3-, and Cl-) in Divalent Metal (M2+) Adsorption on Oxide and Hydroxide Surfaces in Salt Solutions. Am. J. Sci. 1999, 299, 828−899. (19) Wen, Z.; Zhi-Hui, H. Biological and Ecological Features of Inland Saline Waters in North Hebei, China. Int. J. Salt Lake Res. 1999, 8 (3), 267−285. (20) Li, M.; Kang, S.; Zhu, L.; Wang, F.; Wang, J.; Yi, C.; Fang, X.; Xie, M. On the Unusual Holocene Carbonate Sediment in Lake Nam Co, Central Tibet. J. Mt. Sci. 2009, 6 (4), 346−353. (21) Yang, X.; Williams, M. A. J. The Ion Chemistry of Lakes and Late Holocene Desiccation in the Badain Jaran Desert, Inner Mongolia, China. Catena 2003, 51 (1), 45−60. (22) Arp, G.; Hofmann, J.; Reitner, J. Microbial Fabric Formation in Spring Mounds (“microbialites”) of Alkaline Salt Lakes in the Badain Jaran Sand Sea, PR China. Palaios 1998, 13 (6), 581−592. (23) Yang, X. Chemistry and Late Quaternary Evolution of Ground and Surface Waters in the Area of Yabulai Mountains, Western Inner Mongolia, China. Catena 2006, 66 (1−2), 135−144. (24) Baatar, B.; Chuluun, B.; Tang, S.-L.; Bayanjargal, O.; Oyuntsetseg, B. Vertical Distribution of Physical−chemical Features of Water and Bottom Sediments in Four Saline Lakes of the Khangai Mountain Region, Western Mongolia. Environ. Earth Sci. 2017, 76 (3), 130. (25) Isupov, V. P.; Ariunbileg, S.; Razvorotneva, L. I.; Lyakhov, N. Z.; Shvartsev, S. L.; Vladimirov, a. G.; Kolpakova, M. N.; Shatskaya, S. S.; Chupakhina, L. E.; Moroz, E. N.; et al. Geochemical Model of Uranium Accumulation in Shaazgai-Nuur Lake (Northwestern Mongolia). Dokl. Earth Sci. 2013, 448 (1), 143−148. (26) Namsaraev, Z. B.; Gorlenko, V. M.; Buryukhaev, S. P.; Barkhutova, D. D.; Dambaev, V. D.; Dulov, L. E.; Sorokin, V. V.; Namsaraev, B. B. Water Regime and Variations in Hydrochemical

lakes is isolated irreversibly and spontaneously via formation of aragonite or calcite. Results of this study suggest that alkaline lakes play important roles in effective CO2 sequestration, but additional studies must be conducted to elucidate the effects of CO2 sequestration in saline alkaline lakes.

■

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsearthspacechem.8b00046. Saturation states of alkaline lake water; XRD patterns of the reacted samples; changes of log IAPs with respect to CaCO3·nH2O and MgCO3·nH2O; summary of the solution compositions, log activities of activities of Ca2+, Mg2+, and CO32−, log IAP with respect to CaCO3·nH2O, MgCO3· nH2O, and hydromagnesite, Mg/Ca in solid, carbonate mineralogy, and the assigned stage number (PDF)

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: fukushi@staff.kanazawa-u.ac.jp; phone: +81-76-2646520; fax: +81-76-264-6545. ORCID

Keisuke Fukushi: 0000-0003-0398-8950 Author Contributions

The manuscript was written with contributions from all authors. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS Discussions with N. Hasebe are greatly appreciated. Financial support was provided by KAKENHI from the Ministry of Education, Culture, Sports, Science, and Technology (MEXT) (Grant No. JP17H06458) and a Grant in Aid for Scientific Research from the Japan Society for Promotion of Science (Grant No JP16H05643).

■

REFERENCES

(1) Duarte, C. M.; Prairie, Y. T.; Montes, C.; Cole, J. J.; Striegl, R.; Melack, J.; Downing, J. A. CO2 emissions from Saline Lakes: A Global Estimate of a Surprisingly Large Flux. J. Geophys. Res. Biogeosciences 2008, 113 (G4), 1−7. (2) Tranvik, L. J.; Downing, J. A.; Cotner, J. B.; Loiselle, S. A.; Striegl, R. G.; Ballatore, T. J.; Dillon, P.; Finlay, K.; Fortino, K.; Knoll, L. B.; et al. Lakes and Reservoirs as Regulators of Carbon Cycling and Climate. Limnol. Oceanogr. 2009, 54, 2298−2314. (3) Finlay, K.; Leavitt, P. R.; Wissel, B.; Prairie, Y. T. Regulation of Spatial and Temporal Variability of Carbon Flux in Six Hard-Water Lakes of the Northern Great Plains. Limnol. Oceanogr. 2009, 54, 2553−2564. (4) Finlay, K.; Vogt, R. J.; Bogard, M. J.; Wissel, B.; Tutolo, B. M.; Simpson, G. L.; Leavitt, P. R. Decrease in CO2efflux from Northern Hardwater Lakes with Increasing Atmospheric Warming. Nature 2015, 519 (7542), 215−218. (5) Müller, B.; Wang, Y.; Wehrli, B. Cycling of Calcite in Hard Water Lakes of Different Trophic States. Limnol. Oceanogr. 2006, 51 (4), 1678−1688. (6) Drever, J. I. The Geochemistry of Natural Waters: Surface and Groundwater Environments, 3rd ed.; Prentice Hall: New York, 1997. I


Article

ACS Earth and Space Chemistry Characteristics of the Soda Salt Lake Khilganta (Southeastern Transbaikalia). Water Resour. 2010, 37 (4), 513−519. (27) Reimer, A.; Landmann, G.; Kempe, S. Lake Van, Eastern Anatolia, Hydrochemistry and History. Aquat. Geochem. 2009, 15 (1− 2), 195−222. (28) Kazmierczak, J.; Kempe, S. Modern Terrestrial Analogues for the Carbonate Globules in Martian Meteorite ALH84001. Naturwissenschaften 2003, 90 (4), 167−172. (29) Dongarrà, G.; Hauser, S.; Alaimo, R.; Carapezza, M.; Tonani, F. Hot Waters on Pantelleria Island. Geochemical Features and Preliminary Geothermal Investigations. Geothermics 1983, 12 (1), 49−63. (30) Aiuppa, A.; D’Alessandro, W.; Gurrieri, S.; Madonia, P.; Parello, F. Hydrologic and Geochemical Survey of the lake “Specchio Di Venere” (Pantelleria Island, Southern Italy). Environ. Geol. 2007, 53 (4), 903−913. (31) Coshell, L.; Rosen, M. R.; Mcnamara, K. J. Hydromagnesite Replacement of Biomineralized Aragonite in a New Location of Holocene Stromatolites, Lake Walyungup, Western Australia. Sedimentology 1998, 45 (6), 1005−1018. (32) Connell, T. L.; Dreiss, S. J. Chemical Evolution of Shallow Groundwater along the Northeast Shore of Mono Lake, California. Water Resour. Res. 1995, 31 (12), 3171−3182. (33) Bischoff, J. L.; Stine, S.; Rosenbauer, R. J.; Fitzpatrick, J. A.; Stafford, T. W. Ikaite Precipitation by Mixing of Shoreline Springs and Lake Water, Mono Lake, California, USA. Geochim. Cosmochim. Acta 1993, 57 (16), 3855−3865. (34) Gosselin, D. C. Major-Ion Chemistry of Compositionally Diverse Lakes, Western Nebraska, U.S.A.: Implications for Paleoclimatic Interpretations. J. Paleolimnol. 1997, 17 (1), 33−49. (35) Cloern, J. E.; Cole, B. E.; Oremland, R. S. Seasonal Changes in the Chemistry and Biology of a Meromictic Lake (Big Soda Lake, Nevada, U.S.A.). Hydrobiologia 1983, 105 (1), 195−206. (36) Iwama, G. K.; McGeer, J. C.; Wright, P. A.; Wilkie, M. P.; Wood, C. M. Divalent Cations Enhance Ammonia Excretion in Lahontan Cutthroat Trout in Highly Alkaline Water. J. Fish Biol. 1997, 50 (5), 1061−1073. (37) Cooper, J. J.; Koch, D. L. Limnology of a Desertic Terminal Lake, Walker Lake, Nevada, U.S.A. Hydrobiologia 1984, 118 (3), 275−292. (38) Benson, L. V.; Meyers, P. A.; Spencer, R. J. Change in the Size of Walker Lake during the Past 5000 Years. Palaeogeogr., Palaeoclimatol., Palaeoecol. 1991, 81 (3−4), 189−214. (39) Dejehet, F.; Idrissi, S.; Debuys, R. Magnesium and Occluded Water in Calcium Carbonate Monohydrate. J. Chim. Phys. Phys.-Chim. Biol. 1999, 96 (4), 741−753. (40) Rodriguez-Blanco, J. D.; Shaw, S.; Bots, P.; Roncal-Herrero, T.; Benning, L. G. The Role of Mg in the Crystallization of Monohydrocalcite. Geochim. Cosmochim. Acta 2014, 127, 204−220. (41) Munemoto, T.; Fukushi, K. Transformation Kinetics of Monohydrocalcite to Aragonite in Aqueous Solutions. J. Mineral. Petrol. Sci. 2008, 103 (5), 345−349. (42) Kimura, T.; Koga, N. Monohydrocalcite in Comparison with Hydrated Amorphous Calcium Carbonate: Precipitation Condition and Thermal Behavior. Cryst. Growth Des. 2011, 11 (9), 3877−3884. (43) Königsberger, E.; Königsberger, L.-C.; Gamsjäger, H. LowTemperature Thermodynamic Model for the System Na2CO3− MgCO3−CaCO3−H2O. Geochim. Cosmochim. Acta 1999, 63 (19), 3105−3119. (44) Bowler, J. M. Spatial Variability and Hydrologic Evolution of Australian Lake Basins: Analogue for Pleistocene Hydrologic Change and Evaporite Formation. Palaeogeogr., Palaeoclimatol., Palaeoecol. 1986, 54 (1−4), 21−41. (45) Yechieli, Y.; Wood, W. W. Hydrogeologic Processes in Saline Systems: Playas, Sabkhas, and Saline Lakes. Earth-Sci. Rev. 2002, 58 (3−4), 343−365. (46) Ogino, T.; Suzuki, T.; Sawada, K. The Formation and Transformation Mechanism of Calcium Carbonate in Water. Geochim. Cosmochim. Acta 1987, 51 (10), 2757−2767.

(47) Brečević, L.; Nielsen, A. E. Solubility of Amorphous Calcium Carbonate. J. Cryst. Growth 1989, 98 (3), 504−510. (48) Blue, C. R.; Giuffre, A.; Mergelsberg, S.; Han, N.; De Yoreo, J. J.; Dove, P. M. Chemical and Physical Controls on the Transformation of Amorphous Calcium Carbonate into Crystalline CaCO3polymorphs. Geochim. Cosmochim. Acta 2017, 196, 179−196. (49) Yagi, S.; Fukushi, K. Removal of Phosphate from Solution by Adsorption and Precipitation of Calcium Phosphate onto Monohydrocalcite. J. Colloid Interface Sci. 2012, 384 (1), 128−136. (50) Szumińska, D. Changes in Surface Area of the Böön Tsagaan and Orog Lakes (Mongolia, Valley of the Lakes, 1974−2013) Compared to Climate and Permafrost Changes. Sediment. Geol. 2016, 340, 62−73. (51) Walther, M.; Dashtseren, A.; Kamp, U.; Temujin, K.; Meixner, F.; Pan, C. G.; Gansukh, Y. Glaciers, Permafrost and Lake Levels at the Tsengel Khairkhan Massif, Mongolian Altai, During the Late Pleistocene and Holocene. Geosciences 2017, 7 (3), 73.

J


Control of Water Chemistry in Alkaline Lakes: Solubility of

Recommend Documents