Icarus 198 (2008) 277–279

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Note

The heptahydrate of sodium sulfate: Does it have a role in terrestrial andplanetary geochemistry?

Christopher Hall ∗, Andrea Hamilton

School of Engineering and Electronics and Centre for Materials Science and Engineering, The University of Edinburgh, The King’s Buildings, Edinburgh EH9 3JL, UK

a r t i c l e i n f o a b s t r a c t

Article history:Received 27 April 2008Revised 8 July 2008Available online 29 July 2008

Keywords:IcesMars, surfaceJupiter, satellitesMineralogy

Sodium sulfate readily forms a metastable heptahydrate from concentrated aqueous solutions on cooling to around10 ◦C. It crystallises much more easily than the well recognised and less soluble decahydrate (mirabilite), althoughthe existence of the heptahydrate is almost entirely ignored in the geochemical literature on sodium sulfate. Thereis strong evidence that the heptahydrate is stable below a triple point temperature of −9.5 ◦C at low water vapourpressures, conditions which are found in cold dry environments such as the surface of Mars and the icy moons ofJupiter.

© 2008 Elsevier Inc. All rights reserved.

1. Introduction

The sodium sulfate–water system is an important component of many environ-mental, geological and planetary chemistries. Sodium sulfate occurs in lacustrinebrines (Sack and Last, 1994), evaporites and playas (Crowley and Hook, 1996); in at-mospheric aerosols (Rankin et al., 2002); and in deep Antarctic ice cores (Ohno etal., 2006). It is implicated in terrestrial rock weathering (Goudie and Viles, 1997)and is thought to occur in the bedded layers of hydrated salts present on the sur-face of Mars (Mangold et al., 2008; Kargel et al., 2007). It is considered a constituentof the icy brine present on Ganymede (McCord et al., 2001a) and the surface crustof Europa (Dalton, 2007; Prieto-Ballesteros and Kargel, 2005). The presence of hy-drated salts on Mars is relevant to its palaeohydrology and palaeobiology. Thesehydrates have potential use as water resources to support advanced exploration ofthe mid-latitudes (Bish et al., 2003).

Anhydrous sodium sulfate occurs in several polymorphic forms (Rasmussen etal., 1996). Polymorph V, thenardite, is the most commonly occurring. Below 32.4 ◦C,in the presence of water, thenardite is replaced by the well-known decahydrate,Na2SO4·10H2O, mirabilite. Another hydrate of sodium sulfate, the metastable hep-tahydrate, has been known since the early 19th century (Loewel, 1850). Howeverthe existence of the heptahydrate is neglected in nearly all modern geochemicalwork. We have recently reported the first crystallographic data on the heptahydrate(Hamilton and Hall, 2008). The purpose of this Note is to suggest that the heptahy-drate should be re-integrated into the sodium sulfate–water system, particularly inrelation to geochemical and planetary processes.

2. Characterisation and properties of sodium sulfate heptahydrate

Studies between 1850 and 1908, notably by Loewel (1850), de Coppet (1907)and Hartley et al. (1908), showed that when aqueous sodium sulfate solutionsare cooled, it is almost always the heptahydrate which crystallises first from solu-tion rather than mirabilite. That the heptahydrate is thermodynamically metastablewith respect to mirabilite is clearly indicated by its higher solubility. Hartleyet al. (1908) investigated the temperature at which solutions of various concen-

* Corresponding author.E-mail address: christopher.hall@ed.ac.uk (C. Hall).

0019-1035/$ – see front matter © 2008 Elsevier Inc. All rights reserved.doi:10.1016/j.icarus.2008.07.001

trations spontaneously crystallise on cooling at 3 ◦C/h when subjected to me-chanical friction. They found the results highly reproducible but had difficulty incrystallising mirabilite spontaneously unless a solution was cooled rapidly below−10 ◦C. We have recently prepared the heptahydrate from concentrated (typically3.4 m) sodium sulfate solutions prepared at 40 ◦C and cooled below 15 ◦C. Hep-tahydrate crystals in suspension persist without conversion to mirabilite providedthey are not cooled below about 0 ◦C, which often results in the formation ofmirabilite. We obtained the powder diffraction pattern (Hamilton and Hall, 2008)from a stirred suspension of heptahydrate crystals in the solution from whichthey crystallised using energy dispersive synchrotron radiation. The full structureof heptahydrate has also been obtained by a single crystal method (Oswald et al.,2008). Indexing of the diffraction pattern and the full crystal structure confirmthat this hydrate, described since Loewel as the heptahydrate, does have the for-mula Na2SO4·7H2O. The indexing provides a heptahydrate crystallographic density,1568 kg/m3 (compared with the mirabilite crystallographic density of 1470 kg/m3).We have also shown by an X-ray synchrotron diffraction method (Hamilton etal., 2008) that the heptahydrate crystallises within a porous calcium silicate ma-trix impregnated with supersaturated sodium sulfate solution and cooled to 10 ◦C.There is indirect evidence that this occurs also in several limestones (Rijniers, 2004;Espinosa Marzal and Scherer, 2008).

We find that the metastable heptahydrate converts to the stable decahydratewhen it is brought into contact with a seed crystal of the stable phase as shownin Fig. 1, but the transformation is not yet understood. The heptahydrate crystal ispseudomorphically replaced (partially or completely) by smaller mirabilite crystals,and the resulting form has an opaque appearance whether it is in contact withthe solution from which it crystallised or isolated in a medium such as mineraloil.

3. Solubility and phase behaviour

Early heptahydrate solubility data are collated in Gmelin (1961). Solubilities inwater were later reported by Eddy and Menzies (1940), from which we have es-timated the solubility product Ksp, using Pitzer ion and water activities calculatedusing Frezchem (Marion and Kargel, 2008). From data over the temperature range0–23.7 ◦C, we find ln Ksp = −45.72755 + 0.2126744T − 2.176074 × 10−4 T 2 where Tis the absolute temperature. We note that the heptahydrate/thenardite co-existencetemperature (23.465 ± 0.004 ◦C) was accurately determined at the National Bureauof Standards by Washburn and Clem (1938) in developing calibration standards for

278 Note / Icarus 198 (2008) 277–279

Fig. 1. (A) A crystal of sodium sulfate heptahydrate, removed from the solution from which it crystallised, suspended in mineral oil and placed in contact with a crystal ofmirabilite (to the left of the image). (B) The same crystal after 10 min. Scale bar 1 mm.

Fig. 2. The phase diagram Na2SO4–H2O at 1 bar pressure, showing the solubility curve of metastable heptahydrate (solid line) and the eutectic point heptahydrate–ice EH7.The eutectic point for mirabilite–ice EH10 is shown for comparison (dotted line, italic phase field labels). The diagram was constructed using Pitzer activities computed usingFrezchem v8.2 modified to include heptahydrate Ksp data (see text).

thermometers. Gans (1978) measured the water vapour pressure curve for the hep-tahydrate/thenardite system from −19.1 to 15.0 ◦C; and also measured the soliddensity of the heptahydrate, 1530 ± 20 kg/m3, a value with which our new crystal-lographic density is consistent. The enthalpy of crystallisation has been estimatedby Espinosa Marzal and Scherer (2008). There has been little other thermodynamicwork on the heptahydrate.

In Fig. 2 we show the calculated phase behaviour on cooling aqueous sodiumsulfate solutions. The heptahydrate–ice eutectic temperature TEH7 = −3.11 ◦C is only1.95 ◦C below that of the mirabilite–ice system, but the solution concentration ismuch higher (1.013 m rather than 0.295 m). Increasing pressure drives both eutec-tics to lower temperatures: the calculated heptahydrate–ice eutectic is −4.00 ◦C at100 bar and −12.69 ◦C at 1000 bar. These rather accessible eutectic temperaturessuggest that it may be possible for metastable heptahydrate once formed at lowtemperatures to become trapped in an ice matrix. There is some support for this inour observation that a crystal of heptahydrate immersed in solution in a capillarytube that was cooled from 273 to 150 K at 120 K/h (Oswald et al., 2008; Oswald,private communication) survived without conversion to mirabilite. While the con-tents of the tube solidified, it is unclear whether the ice formed was amorphousor crystalline. No systematic study has yet been carried out to assess the effects ofsample volume on this transformation, which is the subject of future work. Experi-

mental work is required now also to identify the hydrates formed in crystallisationfrom complex brines at low temperatures.

4. Stability under martian surface conditions

Gans (1978) used vapour pressure data to show that there is a mirabilite–heptahydrate–thenardite triple point at a temperature TTP = −9.5 ◦C and a watervapour pressure of 1.58 mbar. As the temperature falls below TTP the range of thestability zone of heptahydrate increases. We show the complete recalculated phasediagram in Fig. 3. It appears that heptahydrate should be a stable phase in manycold dry environments and in particular on the martian surface. At these tempera-tures sodium sulfate should show a normal dehydration sequence: decahydrate ↔heptahydrate ↔ anhydride, and over a wide range of relative humidities the hep-tahydrate is apparently the stable phase. The study of the thermal and radiolyticstability of mirabilite by McCord et al. (2001b) under simulated Europa environ-mental conditions suggests that mirabilite is stable with respect to dehydration for103 years. Dehydration of mirabilite took place at about 245 K. It is unknown whatphase remained. Our analysis suggests that heptahydrate is stable under these low

Note / Icarus 198 (2008) 277–279 279

Fig. 3. The low temperature phase diagram of the Na2SO4–H2O system. Below 273.15 K, the water vapour pressure p is normalised by the vapour pressure p0 of supercooledwater calculated from the Goff–Gratch equation; above 273.15 K p0 is calculated from the Wagner–Pruss equation for liquid water. EH7 and EH10 are the eutectic pointsfor heptahydrate–ice and mirabilite–ice; TP is the mirabilite–heptahydrate–thenardite triple point. AH7 is the co-existence point for heptahydrate and thenardite; and AH10that for mirabilite and thenardite. The Ice–EH7 line is calculated from the Goff–Gratch equation for ice. Phase boundary H7–TP is the co-existence line for heptahydrate andthenardite; H10–TP that for mirabilite-heptahydrate. The vapour pressures along the EH10–AH10 and EH7–AH7 solubility curves are computed directly rather than calculatedfrom the Gans regression equations. Dashed lines show phase boundaries of the metastable heptahydrate. The Mars surface p–T conditions based on data of Chipera andVaniman (2007) are indicated.

temperature and low humidity conditions, but may possibly exist in a dynamic cy-cle if the water vapour pressure is raised by ice sublimation.

5. Conclusions

On cooling an aqueous sodium sulfate solution we note that heptahydrate ratherthan mirabilite generally forms. At lower temperatures we calculate that heptahy-drate is stable over a range of temperature and vapour pressure conditions, includ-ing those we believe exist on the cold, dry surface of Mars. The slow dehydration ofmirabilite under the cold, high vacuum conditions on the surface of Europa and dy-namic dehydration–rehydration reaction suggest that heptahydrate may be present.It may form first, in preference to mirabilite, in the ice eutectic system; or form byslow dehydration of mirabilite; or by dissolution–recrystallisation.

Acknowledgments

We thank EPSRC UK and the Royal Society of Edinburgh for funding for A.H.;STFC UK for synchrotron access; we acknowledge collaboration with Iain Oswaldand Colin Pulham on single crystal diffraction.

References

Bish, D.L., Carey, J.W., Vaniman, D.T., Chipera, S.J., 2003. Stability of hydrous mineralson the martian surface. Icarus 164, 96–103.

Chipera, S.J., Vaniman, D.T., 2007. Experimental stability of magnesium sulfate hy-drates that may be present on Mars. Geochim. Cosmochim. Acta 71, 241–250.

de Coppet, L.C., 1907. Recherches sur la surfusion et la sursaturation. Ann. Chim.Phys. 8 (10), 457–527.

Crowley, J.K., Hook, S.J., 1996. Mapping playa evaporite minerals and associated sed-iments in Death Valley, California, with multispectral thermal infrared images.J. Geophys. Res. 101, 643–660.

Dalton, J.B., 2007. Linear mixture modeling of Europa’s non-ice material based oncryogenic laboratory spectroscopy. Geophys. Res. Lett. 34. L21205.

Eddy, R.D., Menzies, A.W.C., 1940. The solubility of certain inorganic compounds inordinary water and in deuterium water. J. Phys. Chem. 44, 207–235.

Espinosa Marzal, R.M., Scherer, G.W., 2008. Crystallization of sodium sulfate salts inlimestone. Environ. Geol., doi:10.1007/s00254-008-1441-7.

Gans, W., 1978. Thermodynamic stability of sodium sulfate heptahydrate. Z. Phys.Chem. 111, 39–46.

Gmelin, L., 1961. Gmelins Handbuch der anorganischen Chemie, Systeme Nr. 21, Na-trium, Weinheim, Bergstrasse.

Goudie, A.S., Viles, H.A., 1997. Salt Weathering Hazards. Wiley, Chichester.

Hamilton, A., Hall, C., 2008. Sodium sulfate heptahydrate: A synchrotron energy-dispersive diffraction study of an elusive metastable hydrated salt. J. Anal. At.Spectrom. 23, 840–844.

Hamilton, A., Hall, C., Pel, L., 2008. Salt damage and the forgotten metastable sodiumsulfate heptahydrate: Direct observation of crystallization in a porous material,J. Phys. D Appl. Phys., submitted for publication.

Hartley, H., Jones, B.M., Hutchinson, G.A., 1908. The spontaneous crystallisation ofsodium sulphate solutions. J. Chem. Soc. Trans. 93, 825–833.

Kargel, J.S., Furfaro, R., Prieto-Ballesteros, O., Rodriguez, J.A.P., Montgomery, D.R.,Gillespie, A.R., Marion, G.M., Wood, S.E., 2007. Martian hydrogeology sustainedby thermally insulating gas and salt hydrates. Geology 35, 975–978.

Loewel, H., 1850. Observations sur la sursaturation des dissolutions salines. Ann.Chim. Phys. 29, 62–127.

McCord, T.B., Hansen, G.B., Hibbitts, C.A., 2001a. Hydrated salt minerals onGanymede’s surface: Evidence of an ocean below. Science 292, 1523–1525.

McCord, T.B., Orlando, T.M., Teeter, G., Hansen, G.B., Sieger, M.T., Petrik, N.G., VanKeulen, L., 2001b. Thermal and radiation stability of the hydrated salt mineralsepsomite, mirabilite and natron under Europa environmental conditions. J. Geo-phys. Res. 106 (E2), 3311–3319.

Mangold, N., Gendrin, A., Gondet, B., LeMouelic, S., Quantin, C., Ansan, V., Bibring,J.-P., Langevin, Y., Masson, P., Neukum, G., 2008. Spectral and geological study ofthe sulfate-rich region of West Candor Chasma, Mars. Icarus 194, 519–543.

Marion, G.M., Kargel, J.S., 2008. Cold Aqueous Planetary Geochemistry withFREZCHEM. Springer Series in Astrobiology and Biogeophysics. Springer-Verlag,Berlin and Heidelberg.

Ohno, H., Igarashi, M., Hondoh, T., 2006. Characteristics of salt inclusions in polarice from Dome Fuji, East Antarctica. Geophys. Res. Lett. 33. L08501.

Oswald, I.D.H., Hall, C., Hamilton, A., Marshall, W.G., Pryor, T.J., Pulham, C.R., 2008.In situ characterization of elusive salt hydrates—The crystal structures of theheptahydrate and octahydrate of sodium sulfate, J. Am. Chem. Soc., submittedfor publication.

Prieto-Ballesteros, O., Kargel, J.S., 2005. Thermal state and complex geology of a het-erogeneous salty crust of Jupiter’s satellite, Europa. Icarus 173, 212–221.

Rankin, A.M., Wolff, E.W., Martin, S., 2002. Frost flowers: Implications for tropo-spheric chemistry and ice core interpretation. J. Geophys. Res. 107 (D23), 4683.

Rasmussen, S.E., Joergensen, J.-E., Lundtoft, B., 1996. Structures and phase transitionsof Na2SO4. J. Appl. Cryst. 29, 42–47.

Rijniers, L.A., 2004. Salt crystallization in porous materials: An NMR study. Ph.D.thesis. Technical University of Eindhoven.

Sack, L.A., Last, W.M., 1994. Lithostratigraphy and recent sedimentation history ofLittle Manitou Lake, Saskatchewan, Canada. J. Paleolimn. 10, 199–212.

Washburn, E.R., Clem, W.J., 1938. The transition temperature of sodium sulfate hep-tahydrate. J. Am. Chem. Soc. 60, 754–757.