What we know and still not know about oceanic salts

Simple chlorides

NaCl (halite) and KCl (sylvite): Both salts are well-known and their physical properties have been determined in a wide range of temperature and pressure [25]. NaCl is the main constituent in rock salt formations. A remarkable property of NaCl concerns the visco-plastic behaviour, which caused the geological formation of salt domes and is also of importance in rock salt mining.

NaCl·2H₂O (hydrohalite): This hydrate represents the only hydrated form of NaCl. According to the phase diagram NaCl-H₂O [26] it crystallizes in form of hexagons below 273.3 K (0.2°C) at NaCl concentrations higher than 26.3 mass%. The crystal structure was determined at 150 K [27]. Sodium ion is coordinated octahedrally by four water molecules and two chloride ions in cis-position.

KCl·H₂O: In Gmelins Handbook of Inorganic Chemistry [24] formation of a monohydrate of KCl is mentioned for temperatures near –10°C citing [28–30]. This hydrate should exist only in a very small temperature interval. However, there is no new work reported about this hydrate and [26] stated that the question of existence or no existence cannot be decided. Thus, the hydrate was not considered in the evaluation of the solubility curve in the temperature region between –10 and 0°C.

MgCl₂·RH₂O with R=2, 4, 6, 8 and 12 crystallize from aqueous solutions between the eutectic at 240 K and 523 K [31–33]. All MgCl₂ hydrates are very hygroscopic, the hexa-hydrate will take up water at a relative air humidity above 30% at ambient temperature [34]. For the hydrates R=1–6 thermodynamic standard data at 298.15 K are listed in [35].

MgCl₂·12H₂O represents a cryo-hydrate, which congruently melts at (257.5±0.5) K [36]. These authors report that the crystals are swimming on a saturated solution, which is remarkable, but was not reassessed later. The crystal structure of [37] was re-determined recently confirming the atom positions, but including the hydrogen atom positions [38].

MgCl₂·10H₂O was very recently prepared under high pressure from amorphous MgCl₂ hydrate at low temperature and crystal structure was determined from the crystal powder by means of synchrotron X-ray and neutron diffraction [39].

MgCl₂·8H₂O is formed from saturated MgCl₂ solution below (273.3±0.3) K [36]. Van’t Hoff and Meyerhoffer observed a metastable octahydrate, which they denoted as the β-form. In their notation the α-form is stable, however, later the formation of a metastable form could not be reproduced [38, 40]. Even D’Ans [41] cites the upper formation temperatures for the α- and β-form at 269.7 K and 263.3 K, respectively.

MgCl₂·6H₂O (bischofite) represents the usual hydrate at ambient conditions. Experimentally determined densities vary between 1.56 and 1.59 g/cm³. The crystals melt incongruently at 389.85 K (116.7°C) [41] p. 71, the metastable congruent melting temperature is already at 390.05 K (117°C). The enthalpy of melting was determined as 41.2 cal/g (=172.4 J/g) [42]. Later determinations yielded 171 J/g [43]. Enthalpies of crystallization of the hexa-hydrate were evaluated on the basis of caloric measurements and temperature derivatives in the solubility diagram [44]. The specific heat capacity, Cp, of the crystals between 20 and 70°C is given with (0.3788±0.0070) cal⋅g^–1⋅K^–1 (=1.585±0.029 J⋅g^–1⋅K^–1) [42]. For the molten hexa-hydrate the same authors gave a value of 0.4911 cal/g (=2.055 J⋅g^–1⋅K^–1) in the interval from 117.6 to 129.7°C. With a background in latent heat storage applications DSC measurements have been performed and Cp as a function of temperature and heat of melting (34.6±0.5 kJ⋅mol^–1) had been estimated [45].

Thermal dehydration was extensively investigated under dynamic [46, 47] and static conditions [48] with the focus on hydrolysis, when the lower hydrates are formed. Dehydration enthalpies were derived from vapour pressure measurements by [49].

MgCl₂·4H₂O: Crystal structure from a product formed by dehydration in solid state shows disorder [50], whereas in crystals grown from molten hexa-hydrate no disorder exists [51]. The latter gives a density of 1.65 g⋅cm^–3 at 200 K.

MgCl₂·2H₂O: Crystal structure was solved for the powdered product from dehydration in solid state [52].

Calcium chloride forms a number of stable and metastable hydrates. Temperature ranges of crystallization from aqueous solution and density data are listed in Table 3. The crystal structure of the mono- and 1/3-hydrate is unknown.

A discussion on the lattice energies of the various hydrates can be found in [61]. Heat capacities and phase transition enthalpies between 270 and 400 K were determined for the hexa-, tetra- and di-hydrate [62]. Cooling down to the low-lying cryohydratic point (–55°C) generates no higher hydrate than the hexa-hydrate [63]. Cryoscopic Raman spectra of fluid inclusions do also not indicate other hydrates of CaCl₂ than the hexahydrate [64, 65].

CaCl₂·6H₂O (antarcticite) is the stable hydrate from cryo-hydratic temperature 223.45 K (–49.7°C) to 302.85 K (29.7°C). As a mineral antarticite occurs in cold regions as in Antarctica.

According to the phase diagram CaCl₂-H₂O [60] the hexa-hydrate melts incongruently at 302.85 K (29.7°C) with precipitation of tetra-hydrate. Since the dystectic point at 303.85 K (30.2°C) is nearby the peritectic at 29.7°C congruent melting can be observed, when the hydrate is heated not too slowly.

It is remarkable that the most reliable data [60] of the solubility curve below 273 K (0°C) were derived from [63].

The melting and freezing behaviour was investigated extensively for the development of latent heat storage materials for room temperature applications [62, 66, 67]. Main obstacle is loss of heat storage capacity due to stratification of the tetra-hydrate and super-cooling of the molten hexa-hydrate.

The pressure dependence of the solubility up to very high pressures was investigated by [68, 69].

Enthalpy of crystallization was re-evaluated by [44] yielding 37.7±0.4 kJ/mol.

CaCl₂·4H₂O: Crystallizes as α- β- γ-form from aqueous solutions, where the α-form represents the stable one. The coordination of the calcium ion in these three modifications is quite different (Fig. 1).

Fig. 1:

Coordination of Ca⁺⁺ in the three modifications of CaCl₂·4H₂O; green spheres=Cl atom, red spheres=O atom, cyan=Ca.

CaCl₂·2H₂O, CaCl₂·H₂O CaCl₂·1/3H₂O: Besides the equilibrium temperatures not much further characterization is reported. Di-hydrate was confirmed as mineral in Iraq named Sinjarite [70]. The 1/3-hydrate was identified relatively late by [71] in the course of calorimetric and phase equilibrium studies.

With modern Raman spectrometers the various hydrates can be identified by their Raman band patterns in the wavenumber range 2800–4000 cm^–1 [65] even in small fluid inclusions.

Sodium and potassium sulfates: The known minerals and phases are listed below.

Na₂SO₄ and K₂SO₄ exist as minerals thenardite and arcanite. From phase chemistry’s point of view thenardite belongs to the orthorhombic phase Na₂SO₄(V). When heated above 180°C it transforms reversibly to the form (III) and at 240°C into a hexagonal form (I). Re-formation of (III) from (I) on cooling occurs through another form (II) [72]. The observed transition temperatures vary within 10 K since the transformations are sluggish and the high temperature forms can exist days or weeks at room temperature when stabilized by cations as Ca⁺⁺ [73], Ni²⁺, Sr²⁺, Y³⁺ [74, 75]. Presence of humidity accelerates these transformation processes. Na₂SO₄·10H₂O is known as mineral mirabilite or also named Glauber’s salt. It melts incongruently with precipitation of anhydrous sodium sulfate at 305.533 K (32.383°C) accompanied by an enthalpy change of 78.04 kJ⋅mol^–1 [76]. There were many attempts to use this heat of melting (360 kJ per liter molten hydrate) for heat storage applications, but due to the settling of the anhydrous salt the heat can be recovered only partly via rehydration, if settling is not avoided by technical measures [77–79]. In the crystal structure of the deca-hydrate exists some disorder due to thermally activated librational modes of sulfate ions near room temperature [80]. This was discussed as in agreement with earlier heat capacity measurements, from which a zero-point entropy in the order of R⋅ln2 was derived. The latter was interpreted as hydrogen bond disorder of the water molecules [81]. However, recently it could be shown that at low temperatures (120 K) the disorder disappears [82]. In the light of the new diffraction measurements low temperature heat capacity data and structure are not in agreement. A p-V-T equation of state was established for synthetic deuterated mirabilite for a range of 4–300 K by [83]. Na₂SO₄·7H₂O represents a metastable hydrate, but it can be obtained easily by cooling supersaturated aqueous solutions of sodium sulfate from 40°C to about 293–295 K (20–22°C) [84] or to a little bit lower temperature of about 280 K [82]. The transition temperature to the stable deca-hydrate in an aqueous suspension was determined very accurately as (23.465±0.004)°C [84] and proposed as a secondary temperature calibration point. Interestingly, the hepta-hydrate was found also as a mineral (Pustynite) in Kasachstan in salt solution pools in the dessert [85]. Crystal structure analysis by Oswald [86] indicated disorder in one oxygen position of a water molecule at 150 K. Our re-investigation revealed no disorder at 150 K [82], but the structure had to be described in a unit cell of double size than had been used by Oswald et al.

Na₂SO₄·8H₂O was crystallized recently at high pressure (1.5 GPa) from an 3.4 molal solution at 293 K and the structure was solved by in-situ single crystal diffraction [86].

Na₂SO₄·3K₂SO₄: This double salt is known as the mineral glaserite. However, the ratio of 1:3 for the two sulfates as given in the formula represents a limiting case. At enhanced temperatures and high Na/K ratio in the solution glaserite can uptake Na₂SO₄ reaching ratios up to 1:1. The latter salt is sometimes named aphtitalite. Thus, between these two ratios the phase represents a solid solution of glaserite with Na₂SO₄ [87], which can be understood structurally. Cations can occupy three different sites in the glaserite structure with coordination numbers 6 (1b), 10 (2d, two-fold), and 12(1a) [88]. At low temperatures Na⁺ resides only in 1b, at increasing T up to 50% of the two-fold 2d site can be occupied by Na⁺ resulting in a 1:1 ratio of Na⁺/K⁺ in the lattice. According to the chemical analyses of Foote [89] also at 25°C about 4% variation in the sodium sulfate occurs. We had studied the system Na₂SO₄-K₂SO₄-H₂O up to 200°C where the Na/K ratio increased to 1.25 at the extremum [90].

Calcium sulfate phases

Calcium sulfate occurs as minerals anhydrite, CaSO₄ , gypsum, CaSO₄·2H₂O and bassanite, CaSO₄·0.5H₂O. According to the solubility diagram CaSO₄-H₂O the hemi-hydrate is always metastable. Crystal structures of anhydrite and gypsum are well-known. A discussion and investigation of the morphology of gypsum (twinning, intergrowth of crystals etc.) in relation to growth conditions and crystal structure is given by [128]. Further studies on gypsum morphology and twinning were provided by [129, 130]. Estimations of the transition temperature between gypsum and anhydrite vary from 42 to 60°C due to the scatter of experimental solubility data and the obtuse angle, under which the solubility curves are crossing [131].

The structure of the hemi-hydrate itself was under discussion several decades, because of variations in water content and crystallinity. At appropriate conditions (temperature, water activity) the α-form of hemi-hydrate crystallizes from aqueous solutions. When drying gypsum in absence of solution the β-form is obtained. Both forms are characterized by a crystal structure, where the water molecules occupy sites in channels along c-axis. The β-form can be considered as a not completely ordered form of α-CaSO₄·0.5H₂O, although some authors argued for a specific structure of β-hemi-hydrate [132, 133]. Structural disorder is reflected in broadened X-ray diffraction patterns, larger specific surfaces, sub-micro crystal sizes and higher hydration rates to form gypsum with water. The α- and β-form can be distinguished easily by recording DTA curves and looking for the appearance and temperature of an exothermic effect. For the β-form this effect is more pronounced and occurs at much higher temperatures (320–375°C) [134]. The reason for the exothermic effect has obviously to be seen in the collapse of the anhydrous intermediate γ-form (also called “soluble anhydrite”), which has the hemi-hydrate structure preserved, but with empty channels [135]. The steps of dehydration from gypsum through hemihydrate and γ-CaSO₄ were recently confirmed by means of micro-XRD inside of a DTA set-up with fast recording of reflexion patterns [136]. For α-hemihydrate the appearance of the exothermic effect depends also on the crystal morphology [134]. In literature related to building materials the anhydrous forms of CaSO₄ are often denoted as AIII and AII for the γ-form and the stable anhydrite, respectively.

The controversial aspects of the α-CaSO₄⋅0.5H₂O concern the choice of the unit cell, effect of twinning and the water occupation (positions and amount) in the channels of the structure. The preferred symmetry is monoclinic as C2 or I2 [135, 137, 138], but trifold [139] or sixfold [140] twinning can give pseudo-trigonal or pseudo-hexagonal symmetry. Indications for the existence of CaSO₄ -subhydrates with water contents between 0.5<×<0.8 were reported from gravimetric analyses and peak splitting of X-ray patterns in dependence on relative humidity [141, 142]. Weiss and Bräu [140] claimed that such sub-hydrates does not exist and single crystal diffraction patterns are only effected by many-fold twinning. However, from careful single crystal X-ray structure analysis at different humidities correlated with systematic powder X-ray diffraction and thermogravimetric investigations a reversible phase transition between a 0.5 hydrate and a 0.625 hydrate could be observed when crossing a relative humidity between 30% and 75% [143]. Whereas the structure of the hemi-hydrate possess only one type of water channels with 3 molecules water per unit translation along the channel the sub-hydrate has two types of channels, one with 3 and the other with 4 water molecules per unit translation [143]. For the analogous calcium selenate the sub-hydrate with 0.625 H₂O represents the phase crystallizing when a di-hydrate suspension in water is heated in an autoclave to 190°C. It shows the same channel structure as found for the sulfate with 0.625 H₂O [144].

Recently, Robertson and Bish [145] investigated the kinetics of dehydration of gypsum and rehydration of anhydrite under controlled humidity insitu in a diffractometer. At 220 K and 6% relative humidity the extrapolated kinetic curves predict the beginning of gypsum dehydration after about 500 years. Depending on humidity they also discussed besides bassanite with 0.5 H₂O a sub-hydrate with 0.67 H₂O as intermediate phases and its possible role on the surface of Mars. Finally, X-ray patterns sent from the Curiosity rover on Mars revealed 1.5% anhydrite in a sample from Gale crater [146], which underlines the importance of the CaSO₄ phase chemistry for an understanding of the processes on this planet.

CaSO₄⋅0.5H₂O can take up sodium from saturated aqueous NaCl solutions at 70°C. The molar Na⁺/Ca⁺⁺ ratio can reach values near 2/5. The sodium is partly incorporated into the water channels of the hemihydrate structure and partly substituting calcium ions. This process is reflected in a stepwise increase of dehydration temperatures, changes in Raman spectra and X-ray diffraction patterns [147, 148]. The highest ratio corresponds to the structure of the known double salt, sodium penta-hydrate, 5CaSO₄⋅Na₂SO₄⋅3H₂O [147].

Double salts – chlorides

Besides sylvite, the mineral carnallite, KCl·MgCl₂·6H₂O, represents the other most important potash mineral. It was the first mineral mined for potash production in Germany. An important property is the incongruent dissolution behaviour, where solid KCl is liberated and the solution enriches in MgCl₂. Only at high concentration of MgCl₂ carnallite dissolves without decomposition. Crystal structure is built of Mg(H₂O)₆ octahedra and K⁺ ions are situated in the holes of the chloride ion packing similar to perovskite lattice types. Potassium can be substituted by other large one-valence ions like NH₄⁺, Rb⁺, Cs⁺ or Li(H₂O)⁺, (H₃O)⁺ and Cl^– by Br^– and I^–. These substitutions change the lattice symmetry from ortho-rhombic in original carnallite to monoclinic [149]. Contents of bromide, rubidium and caesium as trace elements provide important information for an understanding of the genesis of evaporitic deposits and solutions [150]. If not too fast heated, during drying of carnallite a di-hydrate, KCl·MgCl₂·2H₂O, can be obtained [151], which can also exist in equilibrium with molten MgCl₂ hydrate saturated with KCl at T>160°C [152]. Carnallite melts incongruently in its own crystal water at 167°C. Thereby about 75% of the KCl is settling down as solid phase in the molten MgCl₂ hydrate. When water is evaporated from this melt KCl⋅MgCl₂⋅2H₂O and 1.5KCl·MgCl₂·2H₂O will crystallize [152, 153]. In presence of an excess of MgCl₂ in the hydrous melts solid solution formation between MgCl₂·2H₂O and KCl⋅MgCl₂⋅2H₂O was reported [152]. Interestingly, the chloridic double salt with calcium, the mineral tachhydrite, 2MgCl₂·CaCl₂·12H₂O, melts also incongruently nearly at the same temperature as carnallite [151]. From the chemical composition it can be compared with carnallite, where KCl is substituted by CaCl₂ and a twofold amount of MgCl₂·6H₂O. The coordination principle is the same in tachhydrite as in carnallite, however, crystal symmetry is changed to trigonal-rhombohedral [154, 155]. Whereas melting of tachhydrite liberates solid CaCl₂ or its monohydrate [156] incongruent dissolution in water below 55°C yields solid MgCl₂·6H₂O, bischofite, above 55°C tachhydrite dissolves congruently, that is without decomposition. Due to the components MgCl₂ and CaCl₂ both carnallite and tachhydrite are very hygroscopic. Conditions for deliquescence and efflorescence of these minerals were modelled and discussed [34]. An interesting geological occurrence of tachhydrite exists in the salt formation in Teutschenthal (Germany), where crystals of tachhydrite are in intimate contact with kieserite, MgSO₄⋅H₂O, without any microscopically detectable calcium sulfate between the crystals [157].

At temperatures above 363 K (90°C) a double salt, 2CaCl₂⋅MgCl₂⋅6H₂O, can be crystallized [156].

Double salts of MgCl₂ or CaCl₂ containing NaCl are not known. During solid-gas equilibration at enhanced temperature some incorporation of sodium was postulated by Orekhova et al. [158, 159]. However, double salts between KCl and CaCl₂ are known as 2KCl⋅CaCl₂⋅2H₂O [160, 161] and KCl⋅CaCl₂ (Bäumlerite) [162] crystallizing from very highly concentrated CaCl₂ solution at enhanced temperatures.

Double salts with magnesium sulfate

With Na₂SO₄ the three minerals blödite (astrakanite), Na₂SO₄·MgSO₄⋅4H₂O, löweite, 6Na₂SO4⋅7MgSO₄⋅15H₂O and vanthoffite, 3Na₂SO₄⋅MgSO₄, are known since the time of van’t Hoff [1]. The exact composition of löweite was established not before the chemical analysis of Kühn [163], which was confirmed by crystal structure analysis later [164]. In 1982 a new mineral named konyaite, Na₂SO₄⋅MgSO₄⋅5H₂O, was found in Konya basin in Turkey [165]. Crystals of this double salt can be grown easily from equimolar Na₂SO₄-MgSO₄ aqueous solution evaporating at temperatures around 300 K (27°C) and a humidity of 79%. Unfortunately, the solution composition from which the crystals grew was not determined, but the crystal structure was [166]. Asalt even richer in water, Na₂SO₄⋅MgSO₄⋅10H₂O, was prepared during evaporation of solutions at similar conditions [167, 168]. At low temperature from blödite at 263 K (–10°C) at a humidity buffered with ice within two weeks a Na₂SO₄⋅MgSO₄⋅16H₂O was obtained and the structure solved [169].

With K₂SO₄ magnesium sulfate forms the well-known minerals picromerite (schönite), K2SO₄⋅MgSO₄⋅6H2O, leonite, K₂SO₄⋅MgSO₄⋅4H₂O, and langbeinite, K₂SO₄⋅2MgSO₄. Lowest possible formation temperature of langbeinite from solutions of the seawater system is 325 K (52°C) [31]. According to earlier sources it was 310 K (37°C) [170]. As a quite common mineral in salt deposits of the Zechstein period knowledge about the possibilities of langbeinite formation are quite important for an understanding of the evaporitic deposits.

Leonite and picromerite are intermediate products in processes for K₂SO₄ production from KCl and MgSO₄ hydrates [171]. The crystal structure of leonite has some disorder at room temperature [172], which could be the reason for the reported ability to uptake some sodium sulfate [4]. However, Braitsch [3] discards the possibility of solid solution formation with arguments from a chemical analysis of large natural leonite crystals and the different crystal structures of bloedite (Na₂SO₄⋅MgSO₄⋅4H₂O) and leonite. To the best knowledge of the author, variable composition of leonite was never investigated in detail. At 100 K leonite structure becomes ordered, passing an intermediate structure between 269 and 121 K [173]. The successive structural transitions were confirmed by vibrational spectroscopy [174]. Standard formation enthalpies are known for leonite, picromerite (schönite) and langbeinite in [35].

Double salts with calcium sulfate

The most stable double salt with Na₂SO₄ represents glauberite, Na₂SO₄⋅CaSO₄. Crystallization of glauberite is often retarded and thus labile phases like 2Na₂SO₄⋅CaSO₄⋅2H₂O (labile salt) [175] or the sodium penta-salt [176] crystallize from solutions.

2Na₂SO₄⋅CaSO₄⋅2H₂O was found as a mineral eugsterite [177]. A salt with composition 5Na₂SO₄⋅CaSO₄⋅6H₂O was found in the area of the Aral lake by Slusareva [178]. Hill had already investigated the metastable phase equilibria of these double salts in the systems Na₂SO₄-CaSO₄-H₂O [176]. As mentioned above, the Na-penta salt can be considered as an end member of the calcium substitution in CaSO₄ hemi-hydrate [147].

K₂SO₄ forms two double salts with CaSO₄: syngenite, K₂SO₄⋅CaSO₄⋅H₂O, and görgeyite, K₂SO₄⋅5CaSO₄⋅H₂O. The latter was discovered as mineral in year 1953 [179], however as a synthetic compound it was known since vant’ Hoff. At hydrothermal conditions near 200°C a anhydrous double salt, K₂SO₄⋅CaSO₄, can crystallize [90].

No double salts are known between calcium and magnesium sulfate.

More complex salts

Salts containing more than three different ions are polyhalite, K₂SO₄⋅MgSO₄⋅2CaSO₄⋅2H₂O, kainite, 4KCl⋅4MgSO₄⋅11H₂O, dansite (in Russian literature “nona salt”), 9Na₂SO₄⋅MgSO₄⋅3NaCl and the so-called Na-polyhalite, 3Na₂SO₄⋅2K₂SO₄⋅25CaSO₄⋅15H₂O. Crystal structure analyses exist for kainite [180], and polyhalite [181, 182]. In old literature, but also up to now for kainite the formula KCl⋅MgSO₄⋅3H₂O is used, but the correct water content according to chemical analysis [183] and crystal structure is 2.75. For d’ansite only structure analyses are reported for the zinc [184] the manganese and iron [185] analogues. Kainite and polyhalite are very abundant evaporitic minerals.

Both from kainite and from polyhalite it is known that without seeding crystallization even from very supersaturated solutions can last several weeks up to months. Also the dissolution kinetics of polyhalite is very slow at ambient temperature.

Stable equilibria in the quinary oceanic system

Jacobus Henricus van’t Hoff was the first, who developed a systematic approach to build up the phase relations in the oceanic system from simple binary systems step-wise adding a component until reaching the hexary system. The papers from his 12 years work on this topic together with his pupils and colleagues are collected in the famous book [1]. Van’t Hoff used thermodynamic relationships and appropriate experimental methods. His main experimental methods were dilatometry, calorimetry, differential vapour pressure measurements and solubility determinations. In a very effective way he had combined these methods to fix invariant points and thus to establish the formation conditions for most of the various salt hydrates and double salts of the hexary oceanic system.

Starting with the binary systems, the lower and upper temperature for the stability of hydrates was determined, then in ternary systems the formation of double salts and its hydrates had to be studied. Figure 2 shows schematically temperature limits for some salt phases. With these temperature limits van’t Hoff constructed so-called paragenetic schemes (Fig. 3) to show the possible simultaneous existence of salt phases and their saturated solution in higher component systems at a chosen temperature in a compact form.

Fig. 2:

Lower and upper temperature limits for some hydrates and double salts.

Fig. 3:

Phase paragenesis scheme after van’t Hoff [1] for the quinary system at 298 K.

The isothermal invariant points were denoted with letters P, Q, R, Z (Fig. 3), which are still in use today (although in the original papers the letters denote not all the same phase equilibria).

Later this work was continued by a large number of scientists and groups as D’Ans, Autenrieth, Serowy, Kurnakov, Zdanovskii, Lepeshkov, Yanat’seva, Soloveva, Balarev, Emons, Holldorf and others, who exclusively applied the solubility method, mostly in order to improve accuracy of solubility description, to refine some invariant points or to extend the temperature range. D’Ans compiled and evaluated the work on oceanic solubility equilibria until 1933 [41].

Taking into account solubility data up to the year 1962 Braitsch [3] discussed certain solubility equilibria in the quinary and hexary system with particular importance for the genesis of evaporitic deposits. Thereby, he points out differences of certain mineral parageneses due to more new investigations since D’Ans and van’t Hoff. One of these parageneses was blödite-kainite at 35°C, which should not exist according to van’t Hoff and also not according to the polythermal 4-salt lines given in [41].

Collections of experimental data until 1972 can be found in the volumes [187]. In the Supplement volume of the “Potassium” volume Gmelins Handbook of Inorganic Chemistry [24] literature of oceanic systems is compiled until 1970.

The most recent evaluation of experimental solubility data from binary until the quinary system was performed by Usdowski and Dietzel [31]. For the multi-component systems in general it covers a temperature range from 0 to 100°C. However, for the subsystem K⁺, Mg⁺⁺, Cl^–, SO₄^––H₂O the authors stated, that the basis of experimental data did not justify to draw reliable solubility diagrams above 55°C. For the quinary system under the condition of saturation with NaCl Usdowski et al. [188] provided an extensive comparison of their invariant points with those given by D’Ans [41] and Autenrieth [170] (Table 5). The methodology for estimating these invariant points is illustrated in Fig. 4 and Fig. 5.

Fig. 4:

Trigonal-prismatic composition space of the quinary system Na^+, K⁺, Mg⁺⁺/Cl^–, SO₄^––H₂O. Two polythermal 4-salt lines intersect at T₁ giving a polythermal invariant 5-salt point.

Fig. 5:

Magnesium concentration of paragenesis R (halite, carnallite, kieserite, kainite) and Q (halite, carnallite, sylvite, kainite) at various temperatures. Intersection of the R and Q branch yields the polythermal invariant point No. 10 in Table 5. Symbols=experimental data, lines=equations from [31].

At a given temperature the relative portions of salts in a solution of the quinary oceanic system Na⁺, K⁺, Mg⁺⁺, Cl^–, SO₄^––H₂O can be presented by a trigonal prism with the salt components at the corners as shown in Fig. 4. An isothermal invariant point at temperature T₁ (4-salt point) is the result of crossing 3-salt lines in this point. If these salt phases exist also at another temperature T₂ they will cross at another composition for T₂ and so on for more temperatures. Plotting the projections of all these 4-salt points on the planes of the prism will yield lines, which then meet in a polythermal invariant 5-salt point. Thus, the experimenter has to take the analytically determined equilibrium concentrations for all the ions, to calculate the salt portions from it and to plot into diagrams consisting of the planes of the trigonal prism. Ideally all these plots should show a crossing point at the same Cl-SO₄ ratio. Due to the analytical errors this will not be the case. Dependent on the available analytical techniques, the accuracy for the various ions will differ and calculations of ion ratios will magnify these uncertainties. One should mention here, that accurate determinations of ion concentrations by the commonly applied ICP or AAS technique requires considerable calibration efforts to ensure relative errors below 1.0% for each element in a salt solution. In Fig. 4 the water content is not included. To consider the water content a fourth dimension would be necessary. However, the water content will also change from point to point and so the absolute ion concentration in mol/kg H₂O. Consequently, also the absolute ion concentrations must meet in a common point, when plotted the relevant equilibrium concentrations. This is usually done, because in this way magnifying effects of errors by calculation of ion ratios are avoided. Therefore, one chooses the concentration of that ion, which changes its concentration sensible enough in relation to the analytical accuracy. This is mostly the magnesium or sulfate ion. All the invariant points shown in Table 2 were determined by plotting the isothermal invariant points as a function of temperature and a suited composition variable as for instance the magnesium concentration in Fig. 5.

As shown in Table 5 quite a number of changes are proposed by Usdowski et al. [188]. These changes concern the phase assemblages as well as some temperature changes. Thus, 13 invariant points given by D’Ans were cancelled and 11 new invariant points are introduced. The most relevant changes concern the inclusion of dansite in the equilibria, the decrease of the upper stability limit for kainite from 83°C (van’t Hoff) to 78°C. The drawback of this compilation is that for the reader the procedure of arriving at these conclusions is not transparent. Also unpublished results had been involved and no accuracy statements have been given for fitting equations derived for the equilibria in subsystems.

Even if one knows all invariant and univariant equilibria of all subsystems the curvature of the bi-variant solubility planes or hyper-planes can only be roughly estimated. Such knowledge is for example important to elucidate concentration gradients (driving forces) or mass balances in geochemistry of evaporites, when evaluating diagenetic processes [3]. However, solubility experiments were performed nearly exclusively for uni- or non-variant equilibria not within the planes of the diagrams.

When assessing our present knowledge about the solubility equilibria in this quinary system we have to point out, that in the areas unsaturated of NaCl only a relatively small number of data is available and a complete evaluation of the system in a temperature range 273–373 K is absent until today.

CaSO₄-containing solutions and the hexary system

The situation becomes worse when adding calcium and one has to switch to the hexary system. Until today, there exists no further detailed assessment of the hexary system of oceanic salts based on experimental data. Many years ago, an IUPAC initiative was started to systematically evaluate solubility data from simple binary up to the complete hexary system. Such a huge task could not be full-filled in an acceptable period of time on a voluntary basis. However, some volumes of data evaluations of subsystems appeared [26, 189].

Except gypsum all other CaSO₄-containing phases require a quite long time (months, years) to equilibrate with its solution. This slow kinetics of equilibration causes a quite large scatter in the reported experimental results. Thus, stability limits between gypsum and anhydrite or the extension of the polyhalite field around ambient temperatures are not fixed accurately enough.

Equilibria between a pair of solids are well defined, if conditions can be realized in the experiments to approach the equilibrium from both sides. In case of the gypsum/anhydrite pair at room temperature only gypsum can nucleate within laboratory time scales, not anhydrite. So the accurate temperature-concentration conditions for the conversions are still too uncertain [131]. In water the transition temperatures are given between 42 and 58°C. Dissolved electrolytes depress the transition temperature due to decreasing the water activity. An actual question is, at which temperature the dehydration process of gypsum will set in, when the solution is saturated with NaCl. Is it above room temperature or below and how much? This problem concerns the use of CaSO₄-based building materials in rock salt environments. Also a new experimental attempt [190] could not reduce the uncertainty, since the new experiments suffered from similar deficiencies as in [191] earlier as discussed by us [131].