Solubility Equilibria in the System Mg(OH)2âMgCl2âH2O from 298 to

Article pubs.acs.org/jced

Solubility Equilibria in the System Mg(OH)2−MgCl2−H2O from 298 to 393 K Melanie Pannach, Sebastian Bette, and Daniela Freyer* Institut für Anorganische Chemie, TU Bergakademie Freiberg, 09599 Freiberg, Leipziger Straße 29, Germany ABSTRACT: The solubility equilibria of the ternary system Mg(OH)2− MgCl2−H2O were determined at temperatures from 298 to 393 K applying equilibration periods of up to 3.5 years. As a result, four thermodynamically stable magnesium chloride hydroxide hydrates (Sorel phases) exist in the ternary system within the investigated temperature range. These are the 3-1-8 phase [3Mg(OH)2·MgCl2·8H2O], the 9-1-4 phase [9Mg(OH)2·MgCl2· 4H2O], the 2-1-4 phase [2Mg(OH)2·MgCl2·4H2O], and the 2-1-2 phase [2Mg(OH)2·MgCl2·2H2O]. The also known 5-1-8 phase [5Mg(OH)2·MgCl2· 8H2O] was found to be metastable in the solid−liquid system. With this work, a reliable solubility data set is now available, for example, to prove the long-term stability of magnesia building materials in the presence of saltbearing media, a challenging demand on the material in a special application as a barrier construction material in salt formation.

1. INTRODUCTION The ternary system Mg(OH)2−MgCl2−H2O represents the scientific basis for magnesia cement applications. Magnesia cement, also called Sorel cement, has been known since 1867 when S. Sorel1 described the formation of a new cement formed by the reaction of caustic magnesium oxide and concentrated magnesium chloride solution. Depending on the MgO content and the MgCl2 concentration of the admixed liquid, various magnesium chloride hydroxide hydrates (also known as Mg oxychlorides or Sorel phases) with the general composition MgkCll(OH)m·nH2O, or written as x-y-z phases according to the double salt hydrate notation xMg(OH)2·yMgCl2·zH2O, are formed during hydration, setting, and hardening of the cement according to reactions 1 and 2.

to permanent humidity changes typical of outdoor situations. For a selected range of applications, however, the material owns specific properties that make using magnesia cement more favorable than ordinary Portland cement (OPC). In particular, magnesia cement has special importance as a building material in salt formation because of its well-known stability in salt-bearing media and salt solutions, especially in magnesium chloridecontaining brines. The potential application of magnesia cement or concrete in barrier constructions, as part of plug and sealing systems for toxic and nuclear waste repositories in salt rock formations,9,10 requires understanding its long-term stability. Therefore, knowledge regarding the thermodynamic solubility equilibria in the solid−solution system (magnesium chloride hydroxides−salt solution system) is of essential importance. 1.1. Survey of Literature Data. Two phases, 3-1-8 and 5-18, have been known for a long time as relevant binder phases of magnesia cement. Their stoichiometries were determined in the 1950s.11−13 Further, magnesium chloride hydroxide phases have been unambiguously characterized as well. These are the 9-1-4, 21-4, 2-1-2, and 3-1-0 phases.14−19 The crystal structures of the magnesium chloride hydroxides, except the 3-1-8 and 3-1-0 phases, were determined over the past decade: the 5-1-8 phase in 2007,20 the 9-1-4 phase in 2010,14 and the 2-1-4 and 2-1-2 phases in 2012,19 whereas the crystal structures of the 3-1-8 and 3-1-0 phases had already been solved in 195313 and 1954,16 respectively. As these phases crystallize as fine intergrowth needles in nm−μm dimensions, and single crystals were neither grown successfully nor found in nature, all crystal structures were accordingly solved from high-resolution X-ray powder diffraction data.

3MgO + MgCl2 + 11H 2O → 2[Mg 2Cl(OH)3 ·4H 2O] = 3Mg(OH)2 ·MgCl2·8H 2O = 3‐1‐8 phase

(1)

5MgO + MgCl2 + 13H 2O → 2[Mg3Cl(OH)5 ·4H 2O] = 5Mg(OH)2 ·MgCl2·8H 2O = 5‐1‐8 phase

(2)

The 3-1-8 phase was also found as a mineral in veinlets of dolomitic marble and named after the locality as korshunovskite.2,3 Since the beginning of the 19th century, magnesia cement has been used for manufacturing special cement flooring, the socalled stone wood floor in living areas as well as highly resilient industrial floors.4,5 In addition, the material is also used as bonding agents for grinding wheel production, artificial stone manufacturing (lithographic stone), as artificial ivory (e.g., billiard balls), and in decorative interior plasters with embedded stone aggregates.6−8 Magnesia cement is not resistant to water or © XXXX American Chemical Society

Received: November 4, 2016 Accepted: March 9, 2017

A

DOI: 10.1021/acs.jced.6b00928 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

Appropriate solubility data for the magnesium chloride hydroxide phases reported in the literature for the system Mg(OH)−MgCl2−H2O are given as either OH−14,21−25 or H+ molalities26 depending on the MgCl2 solution concentration. Conversion between these two units is only possible by knowledge of the ionic product of water in dilute to concentrated MgCl2 solutions. The determination of the OH− solution concentration according to the dissolution reaction of basic magnesium chloride hydrates (e.g., reaction 3) is a more direct way to quantify the solubility equilibria. 3Mg(OH)2 ·MgCl2·8H 2O → 4Mg 2 + + 6OH− + 2Cl− + 8H 2O (3)

Knowledge of the H+ solution concentration (representative for pH value), however, is more favorable in light of further using Mg(OH)2 and magnesium chloride hydroxide hydrates as pH buffers for salt solutions.27 Relating to radioactive waste repositories, the application of these phases will buffer the pH value in a range close to neutrality and sequestrate dissolved inorganic carbon species (e.g., HCO3, CO32). Therefore, the geochemical conditions in the near field of a nuclear waste repository can be controlled directly, which has a major impact on the potential radionuclide retention processes.28,29 The solubility diagrams (Mg(OH)2 vs MgCl2 molalities) in Figure 1 for 293−323 K and in Figure 2 for 323−393 K

Figure 2. Solubility data of the system Mg(OH)2−MgCl2−H2O at temperatures from 323 to 393 K given in the literature.

predict solubility equilibria in the system Mg(OH)2−MgCl2− H2O, a subsystem of the hexary oceanic salt system. The available solubility data for the system Mg(OH)2− MgCl2 −H 2 O show that Mg(OH) 2 occurs at low MgCl2 concentrations as stable phase at all temperatures (298−393 K). In detail, at 293 K, D’Ans and Katz23 observed Mg(OH)2 as the solid phase up to a magnesium chloride concentration of 4.5 mol MgCl2 (kg H2O)−1 (Figure 1). In contrast, Robinson and Waggaman21 reported that Mg(OH)2 transforms into the 3-1-8 phase already at a MgCl2 concentration between 1.9 and 2.2 mol kg−1. The authors21 applied equilibration periods of six months and D’Ans and Katz23 only 3−4 days, which is the reason they do not observe the 3-1-8 phase. At 323 K, Nakayama24 also observed Mg(OH)2 and the 3-1-8 phase as stable phases with an invariant point between 2.0 and 2.5 mol kg−1 MgCl2, which is comparable to the MgCl2 concentration of the invariant point at 298 K given by Robinson and Waggaman.21 At 373 K (Figure 2), Nakayama25 did not observe the formation of the 3-1-8 phase; instead, he obtained the 2-1-4 phase (given as “2-1-6” by the author; most likely the composition was not analyzed exactly), Mg(OH)2, and the invariant point between both at 5.4 mol MgCl2 (kg H2O)−1. The author reports a steep increase in Mg(OH)2 solubility until the invariant point is reached. Within the crystallization field of the “2-1-6” phase, Nakayama25 analyzed a rapid decrease of the hydroxide concentration with increasing MgCl2 concentration. In the course of our previous investigations at 393 K,14 we found the 9-1-4 phase in 5 mol MgCl2 (kg H2O)−1 solution and the invariant point between Mg(OH)2 and the 9-1-4 phase at 4.1 mol MgCl2 (kg H2O)−1. In higher concentrations of MgCl2 solution (5.8 and 6.7 mol kg −1), the 2-1-4 phase was observed. The 2-1-2 phase could be identified as the stable phase in the most concentrated MgCl2 solutions (7.1−7.9 mol kg−1) at 393 K. A significantly lower Mg(OH)2 solubility was determined for the latter two phases than for the 9-1-4 phase (Figure 2). In summary, the literature data do not provide a consistent picture of the crystallization fields of Sorel phases based on their dependence on temperature. To obtain reliable solubility data, we conducted an extensive experimental program to determine temperature-dependent solubility data in the system Mg(OH)2− MgCl2−H2O over the last 4 years. The resulting data are presented as isotherms [Mg(OH)2 vs MgCl2 molalities] at 298, 313, 333, 353, 373, and 393 K in this paper.

Figure 1. Solubility data of the system Mg(OH)2−MgCl2−H2O at temperatures from 293 to 323 K given in the literature. The cyan symbols refer to the quaternary system Na+, Mg2+//Cl−, OH−−H2O.32

summarize the data from the literature.14,21−25 In addition, there are more publications that insufficiently describe solubility investigations in the system Mg(OH)2−MgCl2−H2O as hydroxide solution concentrations12,15,17,18,30 or numerical values31 are missing. The data of D’Ans et al.32 at 293 K are often regarded as solubilities in the system Mg(OH)2−MgCl2− H2O due to the authors own statement in the summary: “... investigations were done in the “System H2O−Mg(OH)2− MgCl2” ”. However, the authors used in the experiments “just enough” sodium hydroxide solution to precipitate the solid phases in different concentrations of MgCl2 solutions. Thus, the solutions contained unknown NaCl concentrations, and the resulting data belong rather to the system Na+, Mg2+//Cl−, OH−−H2O (these data are also plotted in Figure 1 for comparison as cyan symbols). These data were falsely used for model parametrization of thermodynamic databases33,34 to B



Article

Table 1. Substances Used in This Study Including Chemical Abstracts Service (CAS) Registry Numbers, Suppliers, and Purities As Given by the Suppliers substance

formula

CAS no.

supplier

purity

purification method

magnesium chloride hexahydrate magnesium oxide (type M2923) magnesium hydroxide (type M725)

MgCl2·6H2O MgO Mg(OH)2

7791-18-6 1309-48-4 1309-42-8

Fluka, Sweden Magnesia, Germany Magnesia, Germany

≥99.5% ≥99.0% ≥99.0%

none none none

Table 2. Initial Compositions of the MgCl2 Solutions, Quantities of Solid Starting Materials, and Measured Solubilities after 3 Years in the Equilibrated System Mg(OH)2−MgCl2−H2O at T = 298 K and Ambient Pressure P = 0.1 MPaa initial solid−solution system added solid phase g (MgO (MgO (MgO (MgO Mg(OH)2 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8

equilibrated solid−solution system after >3 years m(MgCl2)

MgCl2 solution b

1.0013) 1.0013)b 1.0030)b 1.0000)b 0.0125 0.1018 0.1010 0.0308 0.0937 1.2068 0.1032 0.1527 0.1080 0.0304 1.2592 0.1104

g

mol (kg [H2O])−1

199.99 197.85 197.53 200.00 152.19 157.25 150.00 70.02 150.01 239.11 150.03 200.00 149.99 70.01 240.01 150.09

0.5 0.5 1.0 1.0 1.5 1.5 2.0 2.5 3.0 3.5 3.5 4.0 4.0 5.0 5.0 5.5

102·m(Mg(OH)2) mol (kg [H2O])−1

0.511 0.512 1.036 1.077 1.521 1.526 2.042 2.556 3.066 3.512 3.584 3.973 4.131 5.057 5.070 5.456

0.045 0.031 0.070 0.106 0.113 0.134 0.165 0.143 0.175 0.187 0.183 0.205 0.202 0.237 0.264 0.294

solid phase Mg(OH)2 Mg(OH)2 Mg(OH)2 Mg(OH)2 Mg(OH)2 3-1-8c 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8

Standard uncertainties for temperature and Mg(OH)2 solubility are u(T) = 0.2 K and u(m(Mg(OH)2)) = 0.0003 mol kg−1, respectively. Relative standard uncertainties for pressure and MgCl2 solubility are ur(P) = 0.05 and ur(m(MgCl2)) = 0.002, respectively. bClear OH− supersaturated MgCl2 solutions (see section 2.2). cMetastable solid phase. a

2.2. Achievement of Equilibrium States. The adjustment of equilibrium in the system Mg(OH)2−MgCl2−H2O at temperatures of 298−393 K was investigated by the isothermal saturation method both from supersaturation as well as from undersaturation. For both cases, suspensions were prepared by the addition of solid phases to MgCl2 solutions with defined concentrations (Tables 2−7). In the case from supersaturation, equilibrium is approached by a decrease of Mg(OH)2 supersaturation. The degree of supersaturation is effected, for instance, by the use of (high) reactive MgO as initial solid phase. The MgO transforms into the thermodynamically stable phase, which is associated with a decrease of Mg(OH)2 supersaturation and a slight change of the initial MgCl2 solution concentration if a chloride-containing solid phase is formed. Approaching from undersaturation was conducted by adding the expected thermodynamically stable solid phase to MgCl2 solutions within the potential crystallization field at the investigated temperature. For solid−solution equilibrium to be reached, this phase is partly dissolved to reach (buildup) the equilibrium Mg(OH)2 solution concentration (accompanied by small changes in the initial MgCl2 solution concentration due to the incongruent dissolution of magnesium chloride hydroxides) at the appropriate temperature. In addition to the suspensions, clear Mg(OH)2 supersaturated solutions were prepared for monitoring of the solid phase formation approaching equilibrium from supersaturation at 298 and 313 K. More details are described below. 2.2.1. Equilibration at 298, 313, and 333 K. For producing clear Mg(OH)2 supersaturated solutions, 1 g of MgO (type

2. EXPERIMENTAL SECTION 2.1. Materials and Solutions. All substances used in the experiments were reagent grade (Table 1). Magnesium chloride solutions were prepared by dissolution of magnesium chloride hexahydrate, MgCl2·6H2O, in deionized water, which had been previously boiled for 10−15 min to remove dissolved carbon dioxide. In this paper, all solution concentrations are given in terms of molality, m, which has the unit mol-solute (kg-H2O)−1, abbreviated to mol kg−1. Commercially available magnesium hydroxide, Mg(OH)2, and highly reactive magnesium oxide, MgO, were used as initial solid phases for equilibration. In addition, some of the MgO was annealed in an electrical furnace at 1600 °C (1873 K) for 48 h to significantly decrease the reactivity. The sintered solid was ground in an agate mortar before usage. Annealed MgO is denoted as “MgO-1600”. Besides magnesium oxide and hydroxide, magnesium chloride hydroxide hydrates, in particular the 3-1-8, 9-1-4, and 2-1-4 phases, were also used as starting materials for the equilibration experiments. The 3-1-8 phase was synthesized by the conversion of MgO (type M2923) in 3.3 mol kg−1 MgCl2 solution for 1 year at 298 K. The 9-1-4 and 2-1-4 phase were obtained by the hydrothermal reaction of MgO (type M2923) with 5.0 and 6.0 mol kg−1 MgCl2 solution, respectively, at 373 K for 14 days. The solid phases were filtered from the mother liquor and washed several times with cold (T < 278 K) deionized water. Finally, each solid was suspended in cold (T < 278 K) ethanol to remove adhered water and dried at room temperature. C



Article

Table 3. Initial Compositions of the MgCl2 Solutions, Quantities of Solid Starting Materials, and Measured Solubilities after 3 Years in the Equilibrated System Mg(OH)2−MgCl2−H2O at T = 313 K and Ambient Pressure P = 0.1 MPaa initial solid−solution system added solid phase (MgO (MgO (MgO (MgO (MgO 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8

equilibrated solid−solution system after >3 years

g

g

1.0011)b 1.0035)b 1.0009)b 1.0003)b 1.0004)b 0.1505 0.0535 0.1005 0.1514 0.1513 0.1546 0.0978 0.1503

199.99 200.00 200.00 200.00 199.70 150.02 70.03 101.33 150.01 150.01 150.00 97.25 150.00

102·m(Mg(OH)2)

m(MgCl2)

MgCl2 solution mol (kg [H2O])

−1

mol (kg [H2O])−1

0.5 1.0 1.0 1.5 1.5 2.0 2.5 2.5 3.0 3.5 4.0 5.0 5.5

0.535 1.026 1.050 1.545 1.581 2.048 2.574 2.607 3.060 3.549 4.075 5.045 5.595

solid phase

0.042 0.083 0.163 0.258 0.202 0.356 0.392 0.402 0.376 0.400 0.410 0.597 0.530

Mg(OH)2 Mg(OH)2 Mg(OH)2 Mg(OH)2 Mg(OH)2 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8

a Standard uncertainties for temperature and Mg(OH)2 solubility are u(T) = 0.2 K and u(m(Mg(OH)2)) = 0.0005 mol kg−1, respectively. Relative standard uncertainties for pressure and MgCl2 solubility are ur(P) = 0.05 and ur(m(MgCl2)) = 0.002, respectively. bClear OH− supersaturated MgCl2 solutions (see section 2.2).

Table 4. Initial Compositions of the MgCl2 Solutions, Quantities of Solid Starting Materials, and Measured Solubilities after 205− 318 Days in the Equilibrated System Mg(OH)2−MgCl2−H2O at T = 333 K and Ambient Pressure P = 0.1 MPaa initial solid−solution system added solid phase

m(MgCl2)

MgCl2 solution g

(MgO Mg(OH)2 Mg(OH)2 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 2-1-4

equilibrated solid−solution system after 205−318 days

1.0006) 0.1004 0.1110 0.2307 0.2551 0.2679 0.2893 0.2730 0.3378 0.1984 0.6237 0.4122

g b

200.01 199.95 200.06 200.01 200.02 200.01 200.00 200.02 200.02 200.00 200.04 200.07

−1

mol (kg [H2O])−1

mol (kg [H2O]) 5.5 1.0 2.0 2.0 2.5 3.0 3.5 4.0 5.0 5.5 6.0 6.3

102·m(Mg(OH)2)

5.338 1.024 1.997 2.003 2.512 2.999 3.482 4.019 4.949 5.352 5.780 5.787

1.253 0.096 0.339 0.311 0.663 0.841 0.859 0.930 1.177 1.336 1.624 0.713

solid phase 3-1-8c Mg(OH)2 Mg(OH)2 Mg(OH)2 Mg(OH)2 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8c 2-1-4

Standard uncertainties for temperature and Mg(OH)2 solubility are u(T) = 0.2 K and u(m(Mg(OH)2)) = 0.001 mol kg−1, respectively. Relative standard uncertainties for pressure and MgCl2 solubility are ur(P) = 0.05 and ur(m(MgCl2)) = 0.002, respectively. bClear OH− supersaturated MgCl2 solutions (see section 2.2). cMetastable solid phase. a

M2923) was added to 200 g of MgCl2 solution with m = 0.5−5.5 mol kg−1 for T = 298 and 313 K and m = 1.0−6.3 mol kg−1 for T = 333 K. Each mixture was stirred for 30 min and filtered through a microfiber glass filter (Munktell, Germany; pore size = 1.0 μm) afterward. In this way, very high concentrations of dissolved magnesium hydroxide, up to approximately 0.06 mol kg−1 in 1.5 mol kg−1 MgCl2 solution at 298 K, were obtained as clear starting solutions. In experiments starting from a suspension, the samples were prepared by mixing Mg(OH)2 (type M725), 3-1-8 or 2-1-4 phases with MgCl2 solutions of m = 0.5−6.3 mol kg−1. Both the clear solutions and suspensions were stored in sealed 250 mL PP bottles that were placed in water baths. At 298.0 ± 0.2 K and 313.0 ± 0.2 K, temperature control was realized by insert thermostats (type EB; Julabo Labortechnik GmbH, Germany). During the following 3−4 years of equilibration time, the samples were constantly stirred using a magnetic stirrer. Some of the samples were prepared in triplicate batches to allow separate

analyses at various times. Thus, unopened samples (in sealed PP bottles) were useable for analysis even after 3 years of storage without CO2 contact. A comparison with occasionally opened sample bottles showed partial carbonization as indicated by the beginning formation of chlorartinite, Mg2(CO3)(H2O)(OH)Cl· H2O, in the solid phase starts after 4 years. For the 333 ± 0.2 K experiments, the initially clear solutions and suspensions stored in sealed PP bottles were placed in a water bath thermostat (type 1083; Gesellschaft für Labortechnik, Germany). An oscillating insert of the thermostat provided mixing of the samples by constant shaking. Sampling was performed after 6 months of equilibration. Details are listed in Tables 2−4. 2.2.2. Monitoring of Phase Formation and Transformation in Clear Hydroxide Supersaturated Solutions at 298 and 313 K. Starting from clear Mg(OH)2 supersaturated solutions, the decrease of OH− concentration in MgCl2 solution from 0.5 to 5.0 mol kg−1 at 298 and 313 K was monitored as a function of time by D



Article

Table 5. Initial Compositions of the MgCl2 Solutions, Quantities of Solid Starting Materials, and Measured Solubilities after 14−53 Days in the Equilibrated System Mg(OH)2−MgCl2−H2O at T = 353 K and Ambient Pressure P = 0.1 MPaa initial solid−solution system added solid phase 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 3-1-8 (MgO MgCl2·6H2O H2O (MgO MgCl2·6H2O H2O (MgO MgCl2·6H2O H2O (MgO MgCl2·6H2O H2O (MgO MgCl2·6H2O H2O (MgO MgCl2·6H2O H2O (MgO MgCl2·6H2O H2O (MgO MgCl2·6H2O (MgO MgCl2·6H2O (MgO MgCl2·6H2O (MgO MgCl2·6H2O slurrye slurrye

equilibrated solid−solution system after 14−53 days m(MgCl2)

MgCl2 solution

102·m(Mg(OH)2)

solid phase

mol (kg [H2O])−1

g

g

0.8662 0.9911 0.5808 1.0012 0.7770 1.0152 0.6908 0.5363 1.0489 0.5878 0.3684)c 1.0490 0.5817 0.3610)c 1.0484 0.5892 0.3661)c 1.0483 0.5905 0.3605)c 1.0480 0.5896 0.3710)c 1.0489 0.5872 0.3651)c 1.0483 0.5888 0.3640)c 0.5644 1.4620)d 0.5606 1.4226)d 0.5661 1.5664)d 0.5683 1.5064)d

31.8806 31.2544 30.8977 29.3019 25.9989 26.0500 26.0718 26.0017 25.0412

0.5 2.0 3.0 4.0 4.5 5.5 5.8 6.5 1.0

0.668 2.059 3.161 4.034 4.496 5.468 5.796 6.491 1.150

0.038 0.398 1.867 1.984 2.412 3.125 3.711 1.643 0.142

Mg(OH)2 Mg(OH)2 3-1-8b 3-1-8 3-1-8 3-1-8 3-1-8b 2-1-4 Mg(OH)2

25.0026

2.0

2.121

0.430

Mg(OH)2

25.0475

3.0

3.159

1.415

Mg(OH)2

25.0025

3.0

3.209

1.319

Mg(OH)2

25.0156

6.0

5.815

3.447

3-1-8 + 2-1-4

25.0149

6.0

5.840

3.567

3-1-8 + 2-1-4

24.9943

6.5

6.062

2.295

2-1-4

24.9920

5.0

5.104

2.959

3-1-8

24.9847

6.0

5.932

1.934

2-1-4

24.9842

6.0

5.964

2.123

2-1-4

24.9053

6.5

6.453

1.551

2-1-4

4.3 4.3

3.493 3.497

2.143 2.112

3-1-8 3-1-8

25.0270 g 29.2962 g

mol (kg [H2O])

−1

a Standard uncertainties for temperature and Mg(OH)2 solubility are u(T) = 0.2 K and u(m(Mg(OH)2)) = 0.005 mol kg−1, respectively. Relative standard uncertainties for pressure and MgCl2 solubility are ur(P) = 0.05 and ur(m(MgCl2)) = 0.002, respectively. bMetastable solid phase. cInitial solid phase as a mixture of MgO, MgCl2·6H2O, and H2O in the molar fraction of 9MgO:MgCl2·6H2O:7H2O. dInitial solid phase as a mixture of MgO and MgCl2·6H2O in the molar fraction of 2MgO:MgCl2·6H2O. eSlurry of 3-1-8 phase and 4.3 mol kg−1 MgCl2 solution.

of initial solid phase (MgO or Mg(OH)2) or a magnesium chloride hydroxide hydrate phase to reach equilibrium from super- as well as undersaturation (details are listed in Tables 5−7) were placed in Teflon cups with a volume of 30 mL situated in TiPd autoclaves. Constant mixing of the samples was provided by rotating each autoclave around its vertical axis with alternating directions. The autoclaves were placed in a metal block thermostat at constant temperatures of 353 ± 0.3, 373 ± 0.3, and 393 ± 0.3 K for 14−53 days. Because of the special construction of the autoclaves (fused silica filter insert; Ø: 29−30 mm; thickness: 1.5−1.8 mm; pore diameter: 10−16 μm; Vogelsberger Quarzglastechnik GmbH, Germany), the subsequent separation of the solid phase from the solution took place

using a pH electrode (ROSS, type Orion 8103BN; Thermo Scientific, USA; 3 M KCl reference solution) for 37−105 days. The electrode potential was recorded regularly during aging of the solutions and correlated to a measured one in clear n mol kg−1 MgCl2 reference solutions with precisely known OH− concentration (calibration solutions). After the formation of a precipitate in the solution could be observed, a small portion of the solid was filtered and identified using X-ray powder diffraction and Raman spectroscopy. 2.2.3. Equilibration at 353, 373, and 393 K. For the determination of the phase equilibria at temperatures of 353, 373, and 393 K of the MgCl2 solution, a mixture of MgCl2·6H2O and water (for solutions with m > 5.5 mol kg−1) and a defined amount E



Article

5130 SB, Germany (20 kV accelerating voltage), after coating the sample with gold.

after transfer of the autoclaves from the thermostat into a special high-temperature centrifuge by centrifugation at 3000 rpm for 40 min at 353 ± 1.0, 373 ± 1.0, and 393 ± 1.0 K. The entire equipment is described in detail by Freyer et al.35 The resulting solid phases were immediately removed from the autoclave for analysis. Sampling of the liquid phase was performed after the autoclave (the part with the separated solution) cooled down. 2.3. Analytics. The resulting solution concentrations were determined by chemical analysis of the Mg2+, Cl−, and OH− content. The magnesium concentration was determined by complexometric titration with 0.05 M sodium ethylendiaminetetraacetate (Na-EDTA) in a NH3/NH4Cl buffered solution (pH 9−10) using Erio T as the indicator36 (uncertainty ur ± 0.3%). The chloride concentration was determined by applying the Mohr method37 with 0.1 M AgNO3 in a Na2CO3/NaHCO3 buffered solution (pH 8−9) using K2CrO4 as the indicator (uncertainty ur ± 0.2%). Because of the very small amounts of hydroxide present in the solutions, determination of the OH− content was carried out without further dilution of the sample. Careful evaluation and validation of several analytical approaches using synthetic samples revealed that the OH− determination in MgCl2 solutions by a direct optical titration method employing phenolphthalein as indicator leads to vast underestimation of the total hydroxide solution concentrations. The equivalence point of phenolphthalein is at pH 8−9. At this pH value, a significant portion of the OH− is still bound as MgOH+, which results in an underdetermination of the total hydroxide content. An appropriate determination of the total OH− content can be assured by potentiometric titration. We used the potentiometric acid-back-titration in the following procedure: after addition of a defined volume of 0.01 M HCl, the resulting excess of the acid was determined by potentiometric back-titration with 0.01 M NaOH and a pH glass electrode (ROSS, type Orion 8103BN; Thermo Scientific, USA). For each solution sample, the OH− concentration determination was carried out at least three times (uncertainty ur ± 0.5%). The given MgCl2 concentrations were calculated from the analyzed chloride content, and the Mg(OH)2 concentration was calculated from the analyzed OH− content. When the analyzed Mg2+ concentrations deviated more than 0.5% (sum of uncertainty ur(Mg2+) and ur(Cl−)) from charge balance, which was scaled on the analyzed total anion concentration, the data set was discarded, and the analysis was repeated if possible. The resulting solid phases were identified using X-ray powder diffraction (XRPD). Patterns for phase identification were taken at room temperature with a laboratory powder diffractometer in Bragg−Brentano geometry (D8 Discover (Bruker, Germany), Cu Kα radiation, Vantec 1 detector). The samples were prepared as flat plates. Phase identification was carried out by matching the measured reflection positions with either reference data available in the PDF-Database38 or with calculated reflection positions from published crystal structure data sets using POWDERCELL.39 In addition, Raman spectroscopy, using a Fourier transform spectrometer [RFG 100/S (Bruker, Germany) with Nd:YAG laser; wavelength, 1064 nm; laser power, 200−300 mW], was employed as a complementary technique for fast phase identification of very small sample amounts (solid phase screening during equilibration period) using measured spectra of pure magnesium chloride hydroxide hydrate phases19 as references. Scanning electron microscopy pictures (SEM) of the solids were taken exemplarily after removal of the adhered mother liquor and drying of the samples with a TESCAN Vega

3. RESULTS AND DISCUSSION 3.1. Equilibration Process at 293 and 313 K. 3.1.1. Phase Formation and Transformation. Approaches to the equilibrium states in the ternary system Mg(OH)2−MgCl2−H2O were carried out from both suspensions (equilibration approaches from super- as well as undersaturation depending on the initial solid phase) and clear Mg(OH)2 supersaturated MgCl2 solutions (see section 2.2). In the clear starting solutions in particular, the solid phase formation was observed as a function of time at 298 and 313 K (Figure 3). The decrease of OH− concentrations in

Figure 3. Time-dependent evolution of the Mg(OH)2 solution concentration and phase formation in (a) 1.51 mol kg−1 MgCl2 solution at 298 K and (b) in 2.04 mol kg−1 MgCl2 solution at 313 K: (circles) electrode potential U [mV] on left axis, (squares) correlated Mg(OH)2 concentration on right axis, (open symbols) data outside calibration range, and (filled symbols) data within calibration range.

these solutions (from 0.5 to 5.0 mol MgCl2 kg−1) was monitored by a pH electrode over 37−105 days. The electrode potential was recorded regularly during aging of the solutions and correlated to that measured in n mol kg−1 MgCl2 reference solutions with precisely known OH − concentration; hence, the OH − concentration could be derived directly from the measured potentials. An uncertainty of ±0.05 mol OH− resulted from a potential drift of ±3 mV during the time of data recording. The precipitation of solids was visually observed typically after 5−98 h at 298 K (15−49 h at 313 K) and correlates with a down-drop of the OH− solution concentration. Subsequent phase transformations were identified by X-ray powder diffraction as a F



Article

Figure 4. X-ray powder diffraction pattern of solid phases formed from clear hydroxide supersaturated MgCl2 solutions dependent on time in (a) 1.51 mol kg−1 MgCl2 solution at 298 K and (b) 2.04 mol kg−1 MgCl2 solution at 313 K. The high signal-to-noise ratio results from small amounts of freshly separated solid phases and from fast scans.

well. These procedures correspond to approaches of equilibrium states from supersaturation. In contrast, approaching the equilibria from undersaturation requires a buildup of the hydroxide concentration in the appropriate MgCl2 solutions. For example, a synthesized 3-1-8 phase was used as initial solid and permanently suspended in 2.0−5.5 mol kg−1 MgCl2 solutions at 298 K. During sampling within the equilibration period and after final solid phase analysis, only the 3-1-8 phase could be identified. The formation of another phase, like the 5-18 phase, was not observed. 3.1.2. Solubility Data Resulting from Super- and Undersaturation Approaches. As a result of all the sample approaches, Mg(OH)2 represents the stable phase in the ternary system Mg(OH)2−MgCl2−H2O at 298 K in MgCl2 solutions up to 1.5 mol kg−1 (Figure 5, circles). The Mg(OH)2 solubility rises with increasing MgCl2 solution concentration. In more concentrated MgCl2 solutions (≥1.5 mol kg−1), the 3-1-8 phase occurs as the stable phase. The determined hydroxide solution concentrations in the approaches from supersaturation (Figure 5, blue symbols) exhibit a systematic deviation from the ones obtained by the approaches from undersaturation (Figure 5, black symbols). The extent of this gap (∼7.5 × 10−4 mol kg−1) clearly exceeds the uncertainty of the analytical method for the hydroxide determination (2 mol kg−1 MgCl2 solutions, a metastable region because the solubility values determined for the 3-1-8 phase in this work are lower in appropriate MgCl2 solutions

reason for the too low given hydroxide solution concentrations 14

by our earlier published data at 393 K

in the region of high

MgCl2 solution concentration (>5.5 mol kg−1, Figure 9b). J



Article

Figure 10. (a) X-ray powder diffraction patterns of the resulting solid obtained at the invariant point between Mg(OH)2 and the 9-1-4 phase at 373 K (lower pattern) and 393 K (upper pattern) and (b) SEM image of the resulting solid obtained at the invariant point between Mg(OH)2 and the 9-1-4 phase at 393 K.

(Figure 9a). D’Ans and Katz23 experiments were probably too short (3−4 days) to observe the formation of the stable 3-1-8 phase in the more concentrated MgCl2 solutions. By later published data of D’Ans et al.,32 the formation of the 3-1-8 phase is described already in lower concentration MgCl2 solutions. The authors estimated the invariant point for Mg(OH)2/3-1-8 phase at 0.7−0.9 mol kg−1 MgCl2 solution according to the experimental values. However, the data refer to the quaternary system Na+, Mg2+//Cl−, OH−−H2O (see section 1.1). Therefore, direct comparisons with these data are not possible, although the solubility of the 3-1-8 phase at 5 mol kg−1, it clearly shifted to higher values compared to the data of the ternary system Mg(OH)2−MgCl2−H2O obtained in this study at 298 K (Figure 9a). According to the solubility data of this work, an increase in the solubility of the 3-1-8 phase can be observed with increasing temperature until the crystallization field of the 9-1-4 phase is reached. The crystallization field of the 3-1-8 phase extends with increasing MgCl2 concentration to the saturation of MgCl2· 6H2O up to 333 K. At higher temperatures and MgCl2 concentrations, the 2-1-4 phase appears. The invariant point of the 3-1-8 and the 2-1-4 phase is expected at ∼5.5 mol kg−1 MgCl2 at 333 K and was experimentally found at 353 K in a 5.8 mol kg−1 MgCl2 solution. As only one data point for the 2-1-4 phase was found at 333 K, it cannot be concluded that this phase might exist at lower temperatures as well. If this were the case, the values of the 3-1-8 phase at 323 K given by Nakayama24 at the highest MgCl2 solution concentration would be metastable states (Figure 9a). At temperatures >353 K, the 3-1-8 phase is replaced by the 9-14 phase (even if a formation of the 3-1-8 phase was observed at 373 K due to its higher solubility in comparison to the 9-1-4 phase at the same MgCl2 concentrations, the 3-1-8 phase is metastable at this temperature), wherefore the Mg(OH)2 field is limited by the 9-1-4 phase. The invariant point Mg(OH)2/9-1-4 phase was found at 373 and 393 K both at ∼4.0 mol kg−1 MgCl2 solution (Tables 6 and 7). The presence of Mg(OH)2 and the 91-4 phase in the solid phase at this solution concentration was indicated by the X-ray powder diffraction patterns (Figure 10a)

as well as by the appearance of needles (9-1-4 phase) and platelike (Mg(OH)2) crystals in the SEM images (Figure 10b). Nakayama25 did not find the 9-1-4 phase in his experiments at 373 K. The author describes Mg(OH)2 and the 2-1-4 phase as occurring solids. The invariant point of Mg(OH)2 and 2-1-4 phase given at 5.4 mol kg−1 MgCl2 is associated with a much higher OH− solution concentration than analyzed in this work for the 9-1-4 phase as well as for the invariant point of 9-1-4 and the 2-1-4 phase in the appropriate MgCl2 solution concentration range (Figure 9b). Therefore, the thermodynamically stable equilibrium was not reached by Nakayama.25 The invariant point between the 9-1-4 and 2-1-4 phases at 393 K could be found at 5.8 mol kg−1 MgCl2 (Figure 11). Within the

Figure 11. X-ray powder diffraction patterns of the resulting solid obtained from 5.8 mol kg−1 MgCl2 solution at 393 K: invariant point of the 9-1-4 and 2-1-4 phases.

crystallization field of the 2-1-4 phase, the solubility decreases with increasing MgCl2 concentration up to ∼7.5−8 mol kg−1 at 393 K. In more concentrated MgCl2 solutions, the 2-1-4 phase transforms into the 2-1-2 phase. The crystallization field of the latter phase is limited by MgCl2·4H2O. In general, the solubility of the Mg oxychloride phases in the ternary system Mg(OH)2−MgCl2−H2O increases with temperK



Article

ature. Maximum values range between 5 and 6 mol kg−1 MgCl2 solution concentrations (locations of appropriate invariant points). The also well-known 5-1-8 phase, 5Mg(OH)2·MgCl2·8 H2O, was not found as a stable phase in the system. In the course of the solubility studies, the intermediate formation of the 5-1-8 phase was observed in the temperature range of 298−313 K only from supersaturated states and always followed by conversion to the 3-1-8 phase. Therefore, the 5-1-8 phase represents a metastable phase in the Mg(OH)2−MgCl2−H2O system, which is in complete agreement with observations described in the literature.12,15,18,31,42,43

(2) Malinko, S. V.; Lisitsyn, A. E.; Purusova, S. P.; Fitsev, B. P.; Khruleva, T. A. Korshunovskite, Mg2Cl(OH)3·nH2O, a new hydrous magnesium chloride. (in Russian). Zapiski Vserossiyskogo Mineralogicheskogo Obshchestva 1982, 111, 324−329. (3) Malinko, S. V. Korshunovskite Mg2Cl(OH)3·nH2O a new hydrous magnesium chloride. Int. Geol. Rev. 1983, 25, 1105−1110. (4) Scherer, R. Die kü nslichen Fußbö den, Wandbeläge und Deckenverkleidungen. Wien und Leipzig, A. Bartleben’s Verlag 1922, 1. (5) Walo Bretschinger AG. http://www.walo.ch/de/produkte/ bodenbelaege/fama-hartsteinholzbelaege/ (accessed Apr 11, 2016). (6) Holleman, A. F.; Wiberg, N. Lehrbuch der Anorganischen Chemie, de Gruyter Verlag, 102, 2007. (7) Grinding wheels - Master Abrasives, www.master-abrasives.co.uk, ISO 9001:2008 at http://www.master-abrasives.co.uk/downloads/ content/Master%20Grinding%20Wheels.pdf (accessed Apr 11, 2016). (8) Grecian Magnesite. http://www.grecianmagnesite.com/markets/ construction/abrasives (accessed Apr 11, 2016). (9) Nuclear Waste Management. http://www.bfs.de/EN/topics/ nwm/repositories/repositories_node.html (accessed Apr 11, 2016). (10) Müller-Hoeppe, N.; Buhmann, D.; Czaikowski, O.; Engelhardt, H.-J.; Herbert, H.-J.; Lerch, C.; Linkamp, M.; Wieczorek, K.; Xi, M. Vorläufige Sicherheitsanalyse für den Standort Gorleben (VSG): Integrität geotechnischer Barrieren, Teil 1 Vorbemessung−AP 9.2, GRS 287 (in German),2012. (11) de Wolff, P. M.; Walter-Levy, L. Structures et formules de quelques constituants du ciment Sorel. C. R. Acad. Sci. Paris 1949, 229, 1232−1234. (12) Cole, W. F.; Demediuk, T. X-ray, thermal, and dehydration studies on magnesium oxychlorides. Aust. J. Chem. 1955, 8, 234−251. (13) de Wolff, P. M.; Walter-Levy, L. The crystal structure of Mg2(OH)3(Cl,Br)·4 H2O. Acta Crystallogr. 1953, 6, 40−44. (14) Dinnebier, R. E.; Freyer, D.; Bette, S.; Oestreich, M. 9 Mg(OH)2· MgCl2·4 H2O, a High Temperature Phase of the Magnesia Binder System. Inorg. Chem. 2010, 49, 9770−9776. (15) Bianco, Y. The formation of basic magnesium chlorides at 50− 175°C by aqueous methods (in French). C. R. Acad. Sci. Paris 1951, 232, 1108−1110. (16) de Wolff, P. M.; Kortlandt, D. Crystal-structure determination from an x-ray powder diffraction pattern of β-Mg2(OH)3Cl. Appl. Sci. Res., Sect. B 1954, 3, 400−408. (17) Demediuk, T.; Cole, W. F.; Hueber, H. V. Magnesium and calcium oxychlorides. Aust. J. Chem. 1955, 8, 215−233. (18) Bianco, Y. The basic chlorides and bromides of magnesium (in French). Ann. Chim. (Paris) 1958, 3, 370−405. (19) Dinnebier, R. E.; Oestreich, M.; Bette, S.; Freyer, D. 2 Mg(OH)2· MgCl2·2 H2O and 2 Mg(OH)2·MgCl2·4 H2O, Two High Temperature Phases of the Magnesia Cement System. Z. Anorg. Allg. Chem. 2012, 638, 628−633. (20) Sugimoto, K.; Dinnebier, R. E.; Schlecht, T. Structure determination of Mg3(OH)5Cl·4H2O (F5 phase) from laboratory powder diffraction data and its impact on the analysis of problematic magnesia floors. Acta Crystallogr., Sect. B: Struct. Sci. 2007, 63, 805−811. (21) Robinson, W. O.; Waggaman, W. H. Basic Magnesium Chlorides. J. Phys. Chem. 1908, 13, 673−678. (22) Gjaldbaek, J. K. Untersuchungen über die Löslichkeit des Magnesiumhydroxids. II. Die Löslichkeitsprodukte und die Dissoziationskonstante der Magnesiumhydroxide. Z. Anorg. Allg. Chem. 1925, 144, 269−288. (23) D’Ans, J.; Katz, W. Magnesiumhydroxyd-Löslichkeiten, pHZahlen und Pufferung im System H2O−MgCl2−Mg(OH)2. KaliZeitschrift für Kali, Steinsalz- und Erdölindustrie sowie Salinenwesen 1941, 35, 37−41. (24) Nakayama, M. A New Basic Triple Salt Containing Magnesium Hydroxide Part IV. The Quinary System KCl-K2SO4-MgCl2-MgSO4Mg(OH)2-H2O at 50°. Bull. Agric. Chem. Soc. Jpn. 1960, 24, 362−371. (25) Nakayama, M. A. New Basic Triple Salt Containing Magnesium Hydroxide Part II. The Quaternary System KCl-MgCl2-Mg(OH)2-H2O at 100°. Bull. Agric. Chem. Soc. Jpn. 1959, 23, 46−48.

4. CONCLUSIONS Solubility equilibria in the ternary system Mg(OH)2−MgCl2− H2O were determined at temperatures of 298−393 K applying equilibration periods of up to 3.5 years. The adjustment of equilibrium was realized by isothermal saturation method from both super- as well as undersaturation. As a result, four thermodynamic stable magnesium chloride hydroxide hydrates (Mg oxychlorides, Sorel phases) were found within the investigated temperature range. These are the 3-1-8 phase [3Mg(OH)2·MgCl2·8H2O], 9-1-4 phase [9Mg(OH)2·MgCl2· 4H2O], 2-1-4 phase [2Mg(OH)2·MgCl2·4H2O], and 2-1-2 phase [2Mg(OH)2·MgCl2·2H2O]. Although the 3-1-8 phase represents the stable phase at lower temperature (298 to

Solubility Equilibria in the System Mg(OH)2âMgCl2âH2O from 298 to

Recommend Documents

Solubility Equilibria in the System Mg(OH)2âMgCl2âH2O from 298 to