Chemical Geology 386 (2014) 16–21

Contents lists available at ScienceDirect

Chemical Geology

j ourna l homepage: www.e lsev ie r .com/ locate /chemgeo

The gypsum–anhydrite paradox revisited

M. Ossorio a, A.E.S. Van Driessche a,⁎, P. Pérez b, J.M. García-Ruiz a,⁎a Laboratorio de Estudios Cristalográficos, Instituto Andaluz de Ciencias de la Tierra, Consejo Superior de Investigaciones Científicas, Universidad de Granada, Granada, Spainb Departamento de Física General, Universidad de La Habana, Cuba

⁎ Corresponding authors.E-mail addresses: alvdries@alvdries.ac.be (A.E.S. Van D

(J.M. García-Ruiz).

http://dx.doi.org/10.1016/j.chemgeo.2014.07.0260009-2541/© 2014 Elsevier B.V. All rights reserved.

a b s t r a c t

a r t i c l e i n f o

Article history:Received 23 April 2014Received in revised form 22 July 2014Accepted 23 July 2014Available online 11 August 2014

Editor: J. Fein

Keywords:PrecipitationPhase transitionStabilityAnhydriteBassaniteGypsum

Despite much experimentation precipitation of anhydrite from solution in conditions similar to those occurringin sedimentary environments has not yet been reproduced in the laboratory. To resolve this long-standing con-tradiction we have monitored the precipitation and stability behaviors of calcium sulfate during experimentslasting up to two years. Calcium sulfate was precipitated from solution between 40 and 120 °C at three differentsalinities and the formed solid phase was sampled at different time intervals (from 2 min up to 2 yr). We foundthat below 80 °C gypsum is the sole primary phase and in the range of 80 to 120 °C gypsumand bassanite are theprimary phases. The stability of the latter increasedwith increasing salinity. As expected, we did not observe pri-mary anhydrite precipitation, but over time phase transition occurred and anhydrite eventually appeared at tem-peratures N80 °C. We show that intrinsic thermodynamic and kinetic properties severely constrain theprecipitation of anhydrite (compared to gypsum and bassanite), and consequently, a considerable amount oftime (e.g. N2 yr at 60 °C) is needed for anhydrite to form. Even so, at a geological time-scale, anhydrite can beconsidered as a pseudo-primary phase thus resolving the long-standing paradox of our inability to directly pre-cipitate anhydrite in the laboratory at temperatures below 120 °C and the abundant presence of anhydrite inevaporitic environments. Our results also show that at low water activity, bassanite becomes an importantphase, which could be relevant to explain its presence on the surface of Mars.

© 2014 Elsevier B.V. All rights reserved.

1. Introduction

Three crystalline phases are known in the CaSO4–H2O system, distin-guishable by their degree of hydration: gypsum (CaSO4·2H2O),bassanite (CaSO4·0.5H2O), and anhydrite (CaSO4). On Earth, calciumsulfate is mainly encountered as gypsum or anhydrite related withevaporitic environments (e.g. Buick and Dunlop, 1990; Warren, 2006),although anhydrite is also frequently present in low temperature hydro-thermal zones (e.g. Blount and Dickson, 1969). Also large amounts ofgypsum have been detected on the surface of Mars (Langevin et al.,2005) and, although bassanite is only rarely found on Earth (e.g. Allenand Kramer, 1953; Apokodje, 1984; Peckmann et al., 2003), it has alsobeen claimed that significant amounts are present on Mars (Wrayet al., 2010). Calcium sulfate is also an important industrial materialand its applications vary from ground gypsum used in agriculture tospecialized plaster products in art and statuary, medical products, andfast-setting/high strength construction products. On the other hand,the formation of calcium sulfate scalants is a recurring nuisance in sev-eral industrial applications (e.g. Ahmed et al., 2004; Mi and Elimelech,2010), being anhydrite the predominant scale phase formed at

riessche), jmgruiz@ugr.es

higher temperatures and gypsum the most frequent phase at lowertemperatures.

Due to the importance of this mineral system in both natural and in-dustrial environments, the precipitation of calcium sulfate has been ex-tensively studied (e.g. Ostroff, 1964; Zen, 1965; Hardie, 1967; D'Ans,1968; Billo, 1986; Freyer and Voigt, 2003), and many of these works,both solubility measurements and thermodynamic calculations, indi-cate that anhydrite should directly precipitate from pure CaSO4 aqueoussolutions above temperatures ranging from 42 °C (D'Ans, 1968) to 58 °C(Hardie, 1967). Consequently, at lower temperatures, gypsum shouldform instead of anhydrite. Hardie (1967) pointed out that for chloridebrines that are saturated with respect to sodium chloride and gypsumthe minimum temperature for direct anhydrite precipitation shouldbe even further lowered, to about 18 °C. However, the stability regionof each hydrate in the CaSO4–H2O system is still unclear because markedinconsistencies exist between the solubility measurements andthermodynamic predictions of stability, and the data obtained inthe laboratory from precipitation studies of calcium sulfate phases(e.g. Zen, 1965; Billo, 1986; Freyer and Voigt, 2003). This uncertaintyin the gypsum–anhydrite transition temperature is also reflected inthe different databases employed in speciation programs such asPHREEQC (Parkhurst and Appelo, 1999).

The multiple parameters, such as salinity, temperature and pres-sure, affecting the anhydrite–gypsum precipitation boundary makeit difficult to determine if anhydrite found in evaporitic deposits

Fig. 1. Plots of experimental and calculated solubility curves of anhydrite, bassanite andgypsum as a function of temperature (0–120 °C) and salinity (0, 0.8, 2.8 and 4.3 MNaCl). Experimentally obtained solubility data (dots) were extracted from the review byFreyer and Voigt (2003) and the references therein. Calculated solubility curves of the cal-cium sulfate phases were obtained with the PHREEQC code employing the LLNL database(solid lines) and the PHREEQC database (dashed lines). Purple lines mark the gypsum–

anhydrite transition temperature obtained with the PHREEQC (dotted lines) and LLNL(full lines) databases.

17M. Ossorio et al. / Chemical Geology 386 (2014) 16–21

precipitated primarily or if it formed secondarily upon transforma-tion of gypsum (e.g. Stewart, 1953; Conley and Bundy, 1959;Holliday, 1970; Shearman, 1983). Experimental studies (e.g. Ostroff,1964; Hardie, 1967; Liu and Nancollas, 1969; Smith and Sweett, 1971;Klepetsanis and Koutsoukos, 1990) have shown that the precipitationof the different calcium sulfate phases is driven by supersaturation,temperature and salinity of the mother solution. To complicate thingseven further it was shown recently that for certain experimental condi-tions nanoparticles of bassanite and/or an amorphous calcium sulfatephase are formed previous to gypsum precipitation (Saha et al., 2012;Van Driessche et al., 2012; Wang et al., 2012). Taking into accountthese new insights and all previous works, it becomes obvious thatstill many, old and new, unanswered questions exist with respect tothe precipitation and stability of calcium sulfate phases. To name justa few, in no particular order: “Can anhydrite be considered as a primaryphase in natural evaporitic environments?”, “Under which conditionsdoes anhydrite form in evaporitic environments?”, “Does anhydrite formas a metastable phase during gypsum precipitation?” and “Why isbassanite abundant on the surface of Mars and not on Earth?”.

Thus, to gain relevant insight to resolve these questions, we de-signed a series of experiments to study the precipitating behavior of cal-cium sulfate phase(s) as a function of temperature (40–120 °C), salinity(0.8, 2.8 and 4.3MNaCl) and timeof reaction (2 min–2 yr). This extend-ed time interval allowed us to carefully test the phase stability andphase transition of primary phase(s) to secondary phase(s) bridgingfast time scale laboratory precipitation experiments and long timescale precipitation occurring in nature. In light of these new experimen-tal data and previously obtained information on the kinetics and ther-modynamics of calcium sulfate precipitation we establish a frameworkto formulate answers to the above postulated questions.

2. Materials and method

2.1. Sample preparation

The CaSO4–NaCl–H2O systemwas experimentally studied by precip-itating calcium sulfate by chemical reaction. Supersaturated solutions ofcalcium sulfate were prepared by mixing directly into glass containerstwo equimolar solutions of calcium chloride dihydrate (CaCl2 2H2O,0.2 M, minimum assay 99%, Sigma) and anhydrous sodium sulfate(Na2SO4, 0.2 M, analytical grade, Sigma). Both reactants were dissolvedin aqueous solutions containing sodium chloride (0.8, 2.8 and 4.3 MNaCl). The selected range of temperatures was 40 to 120 °C. All reac-tants were initially filtered (0.45 μm pore size of filters) and temperedat the corresponding temperature of the experiment. Glass bottleswere sealed in order to avoid evaporation. A series of parallel experi-ments were performed in different glass bottles. Replicate experimentswere stopped at different times, from 2min up to tenmonths, but someexperimentswere sampled for up to two years. The obtained precipitatewasfiltered (0.45 μmpore size offilters)with a vacuum-filtering systemand subsequently slightly washed with water to remove excess salt.Sampleswere introduced in oven (40 °C) for 30min to dry. An addition-al series of experiments was conducted by varying the drying time ofthe products and temperature in order to ensure that there was nophase transition due to drying of the samples. All experimental condi-tions were repeated at least three times in order to check the reproduc-ibility of the obtained results.

2.2. Solution speciation in the calcium sulfate system

Asmentioned in the Introduction the uncertainty in the gypsum–an-hydrite transition temperature is also reflected in the different data-bases that can be employed with speciation programs. In this work weused the PHREEQC code (Parkhurst and Appelo, 1999) to determinethe solubility of the three calcium sulfate phases implementing two dif-ferent databases, LLNL and PHREEQC (Fig. 1).

In Table 1, the saturation indices of the three calcium sulfate phasesfor all the experimental conditions are shown. These were calculatedwith the PHREEQC code (Parkhurst and Appelo, 1999) using both data-bases ((a) LLNL and (b) PHREEQC). In the case of the database PHREEQCno information is available on the solubility of bassanite and thus onlythe SI of gypsum and anhydrite has been calculated.

2.3. Characterization of precipitates by Powder X-Ray Diffraction

A qualitative characterization of the CaSO4 precipitated from solu-tion was performed by Powder X-Ray Diffraction. Measurements wereperformed on a PANalytical diffractometer, X'Pert Pro MPD, withBragg–Brentano geometry, and PIXcel multichannel detector, using Cu

Table 1Saturation indices of calcium sulfate phases for all the experimental conditions, calculatedwith PHREEQC code (Parkhurst and Appelo, 1999) (a) employing the database LLNL and(b) employing the database PHREEQC (bassanite is not included in this database).

a)

Salinity 0.8 M NaCl 2.8 M NaCl 4.3 M NaCl

T/°C Calcium sulfate phases

Gp Bass Anh Gp Bass Anh Gp Bass Anh

40 0.59 −0.07 0.58 0.23 −0.39 0.28 0.07 −0.52 0.1545 0.59 −0.02 0.63 0.24 −0.33 0.33 0.08 −0.46 0.2150 0.60 0.03 0.68 0.25 −0.28 0.38 0.09 −0.41 0.2660 0.61 0.13 0.78 0.27 −0.17 0.49 0.11 −0.30 0.3880 0.65 0.34 0.99 0.32 0.04 0.71 0.17 −0.08 0.6090 0.68 0.45 1.10 0.35 0.16 0.83 0.20 0.04 0.7299 0.71 0.55 1.20 0.39 0.26 0.93 0.24 0.14 0.82110 0.75 0.67 1.33 0.43 0.39 1.06 0.29 0.27 0.95120 0.80 0.78 1.45 0.48 0.51 1.18 0.33 0.39 1.08

b)

Salinity 0.8 M NaCl 2.8 M NaCl 4.3 M NaCl

T/°C Calcium sulfate phases

Gp Anh Gp Anh Gp Anh

40 0.69 0.57 0.61 0.56 0.68 0.6845 0.68 0.60 0.61 0.59 0.68 0.7150 0.68 0.64 0.60 0.62 0.67 0.7560 0.68 0.72 0.60 0.71 0.67 0.8380 0.70 0.95 0.62 0.94 0.70 1.0690 0.72 1.09 0.64 1.08 0.72 1.2099 0.74 1.23 0.66 1.21 0.74 1.34110 0.77 1.40 0.69 1.38 0.77 1.51120 0.81 1.58 0.72 1.55 0.80 1.68

18 M. Ossorio et al. / Chemical Geology 386 (2014) 16–21

anodeλ, kα1, and kα2. Powder X-RayDiffractionwas performed from5to 60° (2θ). The step size was 0.0263° and the counting time 29 s.Diffractograms were analyzed employing the program XPowder(Martin, 2004), using the database ICDD-PDF2.

3. Results

To test the precipitation behavior and phase stability of the CaSO4–

NaCl–H2O systemcalciumsulfatewas formed from solution by chemicalreaction, considering variations in the concentration of sodium chloridein solution (0.8, 2.8 and 4.3 M NaCl) and in the time during which the

Fig. 2. Experimental results showing the precipitated phases of calcium sulfate as a function ofreaction, but theprecipitate remained to bephase pure gypsum. (*For these experimental conditof small amounts of gypsum ismost likely an artifact of sample preparation due to the fact that itemperature decreases and therefore the solution entered in the stability region of gypsum.)

precipitate was in contact with the mother solution (between 2 minand 2 yr). The calcium sulfate phases detected in the precipitate areschematically represented in Fig. 2.

Fig. 2 illustrates that the temperature at which the calcium sulfatephases precipitate is strongly influenced by the solution salinity. Atlow salinity (0.8 M NaCl) gypsum is the first phase to precipitate up to99 °C. As expected, we could not find any primary anhydrite, even notat 120 °C. Instead, bassanite is formed previously, which transforms af-terwards in anhydrite. The lower the temperature, the more time isneeded for the reaction products to transform gradually to anhydrite.For example, at 80 °C, we only could found the first signs of anhydritein the reaction products after ten months.

A different picture arises for the experiments conducted at mediumsalinity (2.8 M NaCl). For these conditions gypsum was the exclusivephase found in the precipitation product up to 60 °C. But, more remark-ably is the increased stability region of bassanite, which is now found aslow as 90 °C and with a stability of up to one month. As a consequencethe transition to anhydrite is slowed down compared to the low salinitycase. At 80 °C, no bassanite precipitates and the transition to anhydriteis already observed after two months.

At high NaCl concentrations (4.3 M NaCl), gypsum is the domi-nant phase up to 60 °C during the first ten months of the experimen-tal observation. Starting from 80 °C, bassanite is the primary phase toprecipitate and stays stable up to two months after the onset of theexperiment, after which it transforms into anhydrite. Noteworthy,at 80 °C after 2 min of reaction already a significant amount of pre-cipitate is formed although the calculated bulk bassanite solubilityindicates that the solution is undersaturated with respect tobassanite (SIBass = −0.08).

At 60 °C experiments were monitored for up to two years, but onlygypsum was detected in the corresponding PXRD patterns.

4. Discussion

The first aspect that becomes evident from our experiments isthat, in the lab, anhydrite does not directly precipitates within itspredicted stability field (≤120 °C). This stands in contradictionwith the large amounts of anhydrite that have formed/are formingin evaporitic environments (b~60 °C) on the surface of our planet.To be able to pinpoint the reason(s) behind our inability to produceanhydrite from solution in the lab (inside its stability field), weneed to evaluate the twomain stages of mineral precipitation: nucle-ation and growth.

time, salinity and temperature. Experiments at 60 °C were monitored up to two years ofions some samples showed amixture of bassanite (main phase) andgypsum. Thepresencet takes about 1min to separate the solid phase from the aqueous solution. During this time,

19M. Ossorio et al. / Chemical Geology 386 (2014) 16–21

The formation of a crystal starts with a nucleation event, and thetime needed for this to occur is characteristic for each phase (commonlyreferred to as induction time, tind). Following the classical nucleationtheory conventionalism (Kashchiev, 1999), the induction time can bedefined as the sum of the time for the critical nucleus formation, tn,and growth to detectable size, tgr;

tind ¼ tn þ tgr : ð1Þ

and tind can be related to the experimental conditions as follows;

tind ¼ 1= J � V ð2Þ

being V the volume of the experiment and J the nucleation rate. This canbe expressed as (e.g. De Yoreo and Vekilov, 2003);

J ¼ A � exp −B � γ3eff =σ

2� �

ð3Þ

where A is a pre-exponential factor, γeff is the effective surface free en-ergy between the growing phase and the solution andσ is the supersat-uration value for the phase being formed. The coefficient B containsother factors, such as the molecular volume, the Boltzmann constantand the absolute temperature.

Eq. (3) shows us how strongly the nucleation rate, and thus also theinduction time, will depend on σ and γeff. Thus, for a solution with agiven concentration of Ca2+ and SO4

2− ions the characteristic inductiontime will be mainly determined by the surface free energy (intrinsic toeach calcium sulfate phase) and supersaturation of each phasewhile as-suming the difference in molecular volume between the three phasescan be neglected (see Supp.mat. for details). The reported effective sur-face free energy (i.e. fitting parameter of the nucleation rate dependen-cy on supersaturation), γeff, for anhydrite is significantly larger than thevalues reported for gypsumand bassanite (see Table 2 and Table SM-1.2for details). Thus, when introducing the values of γeff in Eq. (3), for thesame supersaturation, a much longer induction time is obtained foranhydrite. But, in the stability region of anhydrite (N~50 °C), thesolution will always be more supersaturated with respect to anhydritethan to gypsum and bassanite (Fig. 1). Therefore we need to considerthis difference in supersaturation when comparing induction times.For example, for a solution at 80 °C containing 0.1 M CaSO4 and0.8 M NaCl the supersaturation of the different phases will beσBass = 0.78 b σGp = 1.50 b σAnh = 2.28. Introducing these valuestogether with the γeff values into Eqs. (2) and (3), we can estimatethe ratio of induction times between the different phases (Table 2)assuming that the pre-exponential factor, A, is the same for all thethree phases (see Supp. mat.). This example shows that for theseparticular experimental conditions bassanite should be the first phaseto appear. Although the difference in induction time between gypsumand bassanite is very large, when the precipitation product wassampled after 2min only gypsum is detected, confirming that bassanite,if formed, only is of a transient nature, i.e. a precursor phase to gypsum(Van Driessche et al., 2012; Wang and Meldrum, 2012; Wang et al.,2012). The above estimation also indicates that we should expect towait much longer for anhydrite to start precipitating compared togypsum (~1024 times longer). This is without considering that once aprecursor is formed the effective supersaturation of the second phase

Table 2The effective surface free energy and the calculated ratio of induction times for bassaniteand gypsum with respect to anhydrite for a solution containing 0.1 M CaSO4 and 0.8 MNaCl at 80 °C.

CaSO4 phase γeff/J·m−2 tind Bass/tind Anh tind Gp/tind Anh

Bassanite 0.009 2.26·10−51 N.A.Gypsum 0.040 N.A. 2.25·10−24

Anhydrite 0.090 N.A. N.A.

will be lowered by the amount of solute consumed by the precursorphase, and thus increasing even further this induction time.

But, in evaporitic environments, supersaturation is not created in-stantaneously, as in the example discussed above and the experimentsperformed in this work. Thus, we have to improve our approach tomimic more realistically the precipitation behavior in natural settings.During evaporation of a pond filled with (sea)water the solute concen-tration will steadily increase (assuming there are no new influxes of(sea)water). The induction time for precipitation to occur in such apond can then bemodeled by only taking into account the change in su-persaturation as a function of elapsed evaporation time because,throughout the evaporation process, the effective surface free energywill remain constant for each phase. Thus, with elapsing time the ex-pected induction time will decrease due to the steadily increasing su-persaturation and this will continue until eventually precipitation ofone of the phases occurs lowering the solute concentration (and super-saturation). This behavior is graphically represented in Fig. 3, assuminga constant, arbitrary, evaporation rate and considering the differences insolubility and γeff of gypsum and anhydrite. This figure shows that, al-though the solutionwill becomefirst supersaturatedwith respect to an-hydrite (due to the lower solubility of anhydrite at temperaturesN~50 °C), the probability for anhydrite precipitation to occur in thistime period (0–0.4·101 and 102–0.2·103 for a high and low evaporationrate respectively, inset Fig. 3) is practically zero due to the extremelylong induction times (see green and red squares in Fig. 3). This also ex-plains why a solution can be supersaturated and yet no immediate pre-cipitation will occur. Once the solution becomes supersaturated withrespect to gypsum a crossover of induction time between phases occurs(see green and red squares in Fig. 3) and gypsum will have the shorterinduction time, but this still occurs at very high values of tind (N1036)so no precipitation will occur. With continuing evaporation, supersatu-ration will further increase and consequently induction times of bothphases will decrease and judging from Fig. 3 (see dotted red line) gyp-sum should precipitate before anhydrite, due to its much shorter induc-tion times (~4 orders of magnitude).

Thus, the above rationale advocates that anhydrite should not be ex-pected as primary phase to precipitate during evaporation of pureCaSO4 solutions at mild temperatures (50–100 °C). But, natural watersdo of course differ frompure CaSO4 solutions prepared in the laboratory.These waters usually contain additional components (frequently re-ferred to as additives or impurities), which may alter the precipitationbehavior of the calcium sulfate system. These additives could favor theformation of anhydrite in two ways: (i) through the inhibition of gyp-sum precipitation/growth, or (ii) by accelerating the nucleation/growthof anhydrite by lowering its effective surface free energy/activation bar-rier of solute incorporation. Cody and Hull (1980) succeeded in the pre-cipitation of anhydrite at 60 °C from high salinity aqueous solutions inthe presence of additives (polymaleic/polyacrylic acids and phosphateester). Although this irrevocably demonstrates that additives can effec-tively promote the formation of anhydrite the authors also pointed outthat the used concentrations of additives and solutes are not realistic fornatural environments. Even so, it cannot be ruled out that, in an evapo-ritic context, in the presence of certain additives anhydrite might be theprimary phase to form.

After the nucleation event the second stage of the precipitation pro-cess, i.e. crystal growth, starts and will lead to the typical macroscopiccrystal morphologies observed for the calcium sulfate phases. Whencomparing the growth dynamics of themain faces of anhydrite and gyp-sum (Van Driessche et al., 2011; Morales et al., 2012a), we can find thatanhydrite presents considerably slower growth kinetics than gypsum(Table 3). For example, under similar temperature (~75 °C) and super-saturation values (SI b2), step velocities on the (010) face of gypsum areabove 300 nm/s while step velocities on the (100) face of anhydrite donot exceed5 nm/s. Additionally,Morales et al. (2012b) have shown thattwo-dimensional nucleation is not an effective growth mechanism onthe (100) face of anhydrite, not even at temperatures N80 °C and SI

Fig. 3. Expected induction times of gypsum and anhydrite precipitating from an evaporating (sea)water pond. The solute concentration of Ca2+ and SO42− steadily increases due to evap-

oration and the solution of a pond will become supersaturated with respect to anhydrite and gypsum (blue and pink curves in the inset, two evaporation rates are modeled). Once a su-persaturated state is reached, the crystalline phase can start to form. The waiting time for this to occur is estimated for each phase with respect to the elapsed evaporation time (pink andblue curves in themain figure). The top graphics show the crossover in induction times between anhydrite and gypsum. This occurs at very low supersaturation (green and red squares inthe inset figure) and consequently very long induction times are expected.

20 M. Ossorio et al. / Chemical Geology 386 (2014) 16–21

N2. Consequently, for these conditions screw dislocations are the mainsource of steps, but this growth mechanism presents structure-induced self-inhibition (Pina et al., 1998; Morales et al., 2012b).Hence, the growth of the anhydrite (100) surface is severely inhibited(Shindo et al., 2010), except when a high density of spiral dislocationsis present on the surface (Morales et al., 2012b).

The intrinsic slow growth dynamics of anhydrite have as a conse-quence that even when stable microscopic anhydrite nuclei are formedin an aqueous solution, itwill still take a considerable amount of time forthem to become observable/detectable. This is corroborated by the factthatmetastable formation of gypsumwithin the anhydrite stabilityfieldis even observed in the presence of anhydrite seeds (Nancollas et al.,1973).

Thus, the formation of anhydrite as a primary phase is both hinderedthermodynamically (high effective surface free energy) and kinetically

Table 3Growth kinetics of anhydrite and gypsum.

T/°C βGp/cm·s−1 βAnh/cm·s−1

~60 1.5·10−1 2.0·10−3

~80 4.3·10−1 3.8·10−2

(slow growth kinetics). This constitutes to reconcile the contradictoryobservations in the lab and in nature: enough time is needed forphase transition from gypsum/bassanite to anhydrite to occur in thesystem (mother solution and precipitate phase(s)). This is reflected inFig. 2, where at 80 °C, anhydrite starts to appear after several months.Given enough time anhydrite should also form at 60 °C (and below) inthe laboratory (although we do not see this happen in the lab due tovery long induction times (after 2 yr only gypsum is found)). Additionally,due to the formation of gypsum as a precursor phase the effective super-saturation of anhydrite is very small and hence even longer inductiontimes are to be expected (Eqs. (2) and (3)). Even so, when consideringthe geological time scale, the precipitation of anhydrite can be accountedfor as pseudo-primary (even without the need of additives).

Fig. 2 also confirms that at high ionic strength the precipitation ofbassanite is promoted (e.g. Cruft and Chao, 1970; Fu et al., 2012), butwe have also observed that the stability of bassanite increases with in-creasing salinity, even to such an extent that at high salinity (4.3 MNaCl) bassanite remains stable for more than two months at 80 °C.This is most likely related to the fact that in highly concentrated NaClaqueous solutions, the activity of water aw decreases significantly.High concentrations of NaCl in aqueous solution strongly bind thewater molecules, and thus, the precipitation of bassanite (having 0.5

21M. Ossorio et al. / Chemical Geology 386 (2014) 16–21

molecule of water in its structure) is favored over the precipitation ofgypsum (with 2 molecules of water in its structure). Calculations doneby Tosca et al. (2008) suggest a number ofMartian localities, widely dis-tributed in time and space, hosted fluids thatwere 10 to 100 timesmoresaline that halite-saturated terrestrial waters. Hence, high salinity couldbe the key to understand the abundant presence of bassanite on Mars(Wray et al., 2010).

5. Concluding remarks

In this workwe have established a framework to explain the contra-dictory observations of ubiquitous precipitation of gypsum in the anhy-drite stability region in the laboratory and the abundant presence ofanhydrite in the same stability region in evaporitic environments. Dueto intrinsic thermodynamic and kinetic properties of anhydrite its pre-cipitation is severely constrained compared to gypsum (and bassanite).We argue that anhydrite can be regarded as a pseudo-primary phase inevaporitic settings if the initially formed solid phases are long enough incontact with the mother solution. From our experimental results, wehave determined that the transition to anhydrite at temperatures≥80 °C occurs in an interval of months, and longer times are neededfor the transition to anhydrite to occur at lower temperatures. Our ex-perimental results also demonstrate the large influence of salinity onbassanite stability, which is linked to a decreasedwater activity, and ex-plains the presence (and stability) of bassanite in low water activityenvironments.

Acknowledgments

This work has been carried out within the framework of the projectCGL2010-16882 of the Spanish MINECO and the Consolíder-Ingenio2010 “Factoría de Cristalización” CSD2006-00015. MO acknowledgesthe JAE-Predoc fellowship. We greatly appreciate the input of MADurán-Olivencia on the theoretical aspects of nucleation.

Appendix A. Supplementary data

Supplementary data to this article can be found online at http://dx.doi.org/10.1016/j.chemgeo.2014.07.026.

References

Ahmed, S.B., Tlili, M., Amor, M.B., Bacha, H.B., Eullech, B., 2004. Calcium sulphate scale pre-vention in a desalination unit using the SMCEC technique. Desalination 167, 311–318.

Allen, R.D., Kramer, H., 1953. Occurrence of bassanite in two desert basins in southeasternCalifornia. Am. Mineral. 38, 1266–1268.

Apokodje, E.G., 1984. The occurrence of bassanite in some Australian arid-zone soils.Chem. Geol. 47, 361–364.

Billo, S.M., 1986. Petrology and kinetics of gypsum–anhydrite transitions. J. Pet. Geol. 10(1), 73–86.

Blount, C.W., Dickson, F.W., 1969. The solubility of anhydrite (CaSO4) in NaCl–H2O from100 to 450 °C and 1 to 1000 bars. Geochim. Cosmochim. Acta 33, 227–245.

Buick, R., Dunlop, J.S.R., 1990. Evaporitic sediments of Early Archaean age from theWarrawoona Group, North Pole, Western Australia. Sedimentology 37, 247–277.

Cody, R.D., Hull, A.B., 1980. Experimental growth of primary anhydrite at low tempera-tures and water salinities. Geology 8, 505–509.

Conley, R.F., Bundy, W.M., 1959. Mechanism of gypsification. Geochim. Cosmochim. Acta15, 57–72.

Cruft, E.F., Chao, P.C., 1970. Nucleation kinetics of the gypsum–anhydrite system. Proceed-ings, Symposium on Salt, 3rd. Northern Ohio Geological Society. 1, pp. 109–118.

D'Ans, J., 1968. Der Ubergangspunkt Gips-Anhydrit. Kali. Steinsalz. 5, 109–111.

De Yoreo, J.J., Vekilov, P.G., 2003. Principles of crystal nucleation and growth. In: Dove,P.M., De Yoreo, J.J., Weiner, S. (Eds.), Mineral. Soc. Am. 54, pp. 57–93 (Washington).

Freyer, D., Voigt, W., 2003. Crystallization and phase stability of CaSO4 and CaSO4-basedsalts. Monatsh. Chem. 134, 693–719.

Fu, H., Guan, B., Jiang, G., Yates, M.Z., Wu, Z., 2012. Effect of supersaturation on competi-tive nucleation of CaSO4 phases in concentrated CaCl2 solutions. Cryst. Growth Des.12, 1388–1394.

Hardie, L.A., 1967. The gypsum–anhydrite equilibrium at one atmosphere pressure. Am.Mineral. 52, 171–200.

Holliday, D.W., 1970. The petrology of secondary gypsum rocks: a review. J. Sediment.Petrol. 40, 734–744.

Kashchiev, D., 1999. Nucleation: Basic Theory With Applications. Butterworth,Heinemann, Oxford.

Klepetsanis, P.G., Koutsoukos, P.G., 1990. Spontaneous precipitation of calcium sulfate atconditions of sustained supersaturation. J. Colloid Interface Sci. 143, 299–308.

Langevin, Y., Poulet, F., Bibring, J.P., Gondet, B., 2005. Sulfates in the north polar region ofMars detected by OMEGA/Mars Express. Science 307, 1584–1586.

Liu, S., Nancollas, G.H., 1969. The kinetics of crystal growth of calcium sulphate dihydrate.J. Cryst. Growth 6, 281–289.

Martin, J.D., 2004. Using XPOWDER: a software package for Powder X-Ray Diffractionanalysis. ,p. 105.

Mi, B., Elimelech, M., 2010. Gypsum scaling and cleaning in forward osmosis: measure-ments and mechanisms. Environ. Sci. Technol. 44, 2022–2028.

Morales, J., Astilleros, J.M., Fernández-Díaz, L., 2012a. Nanoscopic characteristics of anhy-drite (100) growth. Cryst. Growth Des. 12, 414–421.

Morales, J., Astilleros, J.M., Fernández-Díaz, L., 2012b. A nanoscopic approach to the kinet-ics of anhydrite (100) surface growth in the range of temperatures between 60 and120 °C. Am. Mineral. 97, 995–998.

Nancollas, G.H., Reddy, M.M., Tsai, F., 1973. Calcium sulfate dihydrate crystal growth inaqueous solution at elevated temperatures. J. Cryst. Growth 20, 125–134.

Ostroff, A.G., 1964. Conversion of gypsum to anhydrite in aqueous salt solutions. Geochim.Cosmochim. Acta 23, 1363–1372.

Parkhurst, D.L., Appelo, C.A.J., 1999. User's guide to PHREEQC (version 2). A computer pro-gram for speciation, batch-reaction, one-dimensional transport, and inverse geo-chemical calculations. Water Resour. Invest. Rep. U.S. Geol. Surv. 99–4259.

Peckmann, J., Goedert, J.L., Heinrichs, T., Hoefs, J., Reitner, J., 2003. The Late Eocene ‘Whis-key Creek’ methane-seep deposit (western Washington State)—part II: petrology,stable isotopes, and biogeochemistry. Facies 48, 241–254.

Pina, C.M., Becker, U., Risthaus, P., Bosbach, D., Putnis, A., 1998. Molecular-scale mecha-nisms of crystal growth in barite. Nature 395, 483–486.

Saha, A., Lee, J., Pancera, S.M., Bräeu, M.F., Kempter, A., Tripathi, A., Bose, A., 2012. New in-sights into the transformation of calcium sulphate hemihydrate to gypsum using time-resolved cryogenic transmission electron microscopy. Langmuir 28, 11182–11187.

Shearman, D.J., 1983. Syndepositional and late diagenetic alteration of primary gypsum toanhydrite. Sixth International Symposium on Salt Northern Ohio Geol. Soc.: Cleve-land, Ohio, pp. 41–50.

Shindo, H., Igarashi, T., Karino, W., Seo, A., Yamanobe-Hada, M., Haga, M., 2010. Stabilitiesof crystal faces of anhydrite (CaSO4) compared by AFM observation of face formationprocesses in aqueous solutions. J. Cryst. Growth 312, 573–579.

Smith, B.R., Sweett, F., 1971. The crystallization of calcium sulphate dihydrate. J. ColloidInterface Sci. 37, 612–618.

Stewart, F.H., 1953. Early gypsum in the Permian evaporates of North-eastern England.Proc. Geol. Assoc. London 64, 33–39.

Tosca, N.J., Knoll, A.H., McLennan, S.M., 2008. Water activity and the challenge for life onearly Mars. Science 320 (5880), 1204–1207.

Van Driessche, A.E.S., García-Ruiz, J.M., Tsukamoto, K., Patiño-Lopez, L.D., Satoh, H., 2011.Ultraslow growth rates of giant gypsum crystals. Proc. Natl. Acad. Sci. U. S. A. 108,15721–15726.

Van Driessche, A.E.S., Benning, L.G., Rodriguez-Blanco, J.D., Ossorio, M., Bots, P., García-Ruiz, J.M., 2012. The role and implications of bassanite as a stable precursor phaseto gypsum precipitation. Science 336, 69–72.

Wang, Y.W., Meldrum, F.C., 2012. Additives stabilize calcium sulfate hemihydrate(bassanite) in solution. J. Mater. Chem. 22, 22055–22062.

Wang, Y.W., Kim, Y.Y., Christenson, H.K., Meldrum, F.C., 2012. A new precipitation path-way for calcium sulfate dihydrate (gypsum) via amorphous and hemihydrate inter-mediates. Chem. Commun. 48, 504–506.

Warren, J.K., 2006. Evaporites: Sediments, Resources and Hydrocarbons. Springer-Verlag,Berlin.

Wray, J.J., Squyres, S.W., Roach, L.H., Bishop, J.L., Mustard, J.F., Noe Dobrea, E.Z., 2010. Iden-tification of the Ca-sulfate bassanite in Mawrth Vallis, Mars. Icarus 209, 416–421.

Zen, E., 1965. Solubility measurements in the system CaSO4–NaCl–H2O at 35, 50 and70 °C and one atmosphere pressure. J. Petrol. 6, 124–164.