Reversible formation of alcohol solvates and their potential use for heat storage

In this study, CaCl2- and MgCl2–alcohol solvates of different stoichiometric quantities of ethyl alcohol (EtOH) and methyl alcohol (MeOH) were synthesized and characterized via coupled thermogravimetric–differential scanning calorimetry, thermogravimetric–mass spectrometric evolved gas analysis (TG-MS), spectroscopic analysis (Raman) methods as well as by X-ray diffraction. Correlations between the obtained calorimetric, thermodynamic, kinetic, and crystallographic data were carried out. The CaCl2–alcohol systems seem suitable for heat storage based on the feasible recovery of the salt. However, Raman spectroscopic results revealed that the MgCl2–EtOH solvates were instable compounds. Irreversible transformation of MgCl2–alcohol solvates related to the formation of alkyl chloride appeared upon heating, as proven by TG-MS and Raman spectroscopic measurements. Pure salt–alcohol solvates could not be prepared under technically applicable conditions. The samples contained at least traces of water. Appearance of side reactions resulting in magnesium oxychlorides, oxyhydroxides, and possible release of HCl with cycling may contribute to corrosion of reactor components. Based on these considerations, MgCl2–alcohol solvate systems are not recommended for heat storage.

To date, the reaction of inorganic salts with water to form salt hydrates under heat release (1) has mainly been studied and optimized [23][24][25][26].Alcohols are known to undergo similar reactions with salts (2).These reactions have only been studied in a casual manner in terms of their application in thermal energy storage [27,28], although the high vapor pressure and low freezing point of EtOH and MeOH facilitate low-temperature applications unlike the aqueous adducts.
During the exothermic association reaction (discharging mode), the solid inorganic salt MX reacts with the gaseous alcohol R-OH forming the salt-alcohol solvate (alcoholate) MXÁzR-OH, where z is the stoichiometric coefficient.The chemical energy stored in the reaction is released as heat of reaction D r H 2 .The reverse reaction (charging mode) is endothermic.Heat energy must be supplied to initiate the dissociation of the MXÁzR-OH adduct.The gaseous product R-OH can be collected and condensed.Therefore, the storage volume can be reduced and the reaction products can be stored separately.As a result, a long-term storage without sensible heat losses is technically feasible.A schematic diagram of the operating principle of a closed thermochemical heat storage system is shown in Fig. 1.
Fast reaction kinetics, precise thermodynamic control, and low regeneration temperatures make salt-alcohol pairs an advantageous option.Recently, the suitability of EtOH solvates of CaCl 2 , MgCl 2 , and their mixtures as heat storage materials for practical implementations has been studied [29].The EtOH sorption ability of CaCl 2 was better than that of MgCl 2 .At high EtOH vapor pressures, overstoichiometric EtOH uptake occurred.The CaCl 2 -EtOH reaction system exhibited convenient sorption properties coupled with good multi-cyclic stability.Consequently, it has a great potential for low-grade thermal energy storage.However, poor reversibility appeared in case of the MgCl 2 -EtOH system with increasing number of cycles, probably caused by irreversible decomposition that strongly reduced the material's sorption performance [29].
Due to the variety of synthesis procedures and conditions, an apparent discrepancy of the stoichiometry of salt alcoholates is found in the published literature.Moreover, present work on the energy storage density is limited.The aim of this work was to prepare and characterize saltalcohol solvates of different stoichiometry based on anhydrous CaCl 2 and MgCl 2 by use of commercially available, specially non-purified chemicals with methods that can be relatively easily transposed into practice.The salt-alcohol solvates were prepared by direct synthesis from CaCl 2 /MgCl 2 and EtOH/MeOH.This preparation method is based on the operating principle of a thermochemical heat store and can easily be adapted to practice.The focus was laid not only on the detailed comparative study of the calorimetric behavior, but also on the mapping of the heat-induced changes in these solvates.Besides coupled thermogravimetric-mass spectrometric (TG-MS) [30,31], thermogravimetric analysis/differential scanning calorimetric (TGA/DSC) [32] techniques are generally applied in heat storage studies.We believed that Raman spectroscopy [33,34] and X-ray powder diffraction (XRD) [35] methods are informative in sample characterization.Use of above techniques on structural changes in salt alcoholates is quite unique.We tried to find relationships between structural changes in the salt-alcohol solvates and the cyclic stability during repeated alcoholation/dealcoholation reactions observed earlier [29].
In solid state, the solvent molecules could be in a welllocalized form and in solution they could be located in the first, second, and so on solvation shell.The interaction energies between the solvent molecules and solvated ions can vary over a broad range.Conducting single-crystal diffraction experiments for each composition was not Dealcoholation (thermal energy charging) process Alcoholation (thermal energy discharging) process possible and may be the subject of another work.Throughout this article, the term built-in form refers to the solvent molecules coordinated to the salt ions and also to the molecules loosely bound to the complexes studied, which dissociate upon crystallization or slightly elevating the temperature.

Experimental
Synthesis of salt-alcohol solvates  1).The alcohol/salt stoichiometric ratios varied between 2 and 8 mol of EtOH or MeOH per mole anhydrous CaCl 2 or MgCl 2 .Samples of 10 g of the respective salt were placed into a round bottom flask that was purged with dry nitrogen (N 2 ) prior to each alcoholate synthesis reaction.The flask was equipped with a magnetic stirrer and two thermometers.A stoichiometric amount of liquid alcohol was dropped onto the vigorously agitated salt under dry N 2 atmosphere, while both the temperature of the sample and the vapor phase were measured.Changes in textural and optical properties of the alcoholated salts were recorded.Preparation and denomination of the samples obtained are summarized in Table 1.The prepared samples were stored in closed glass vessels in a vacuum desiccator over P 2 O 5 .

Simultaneous thermogravimetric analysis and differential scanning calorimetry
Experimental alcohol/salt ratios of the salt-alcohol solvates prepared at different theoretical stoichiometric molar ratios, according to Table 1, and associated heat absorptions were identified by simultaneous TGA/DSC.The horizontal furnace TGA/DSC 1 from Mettler Toledo was used.Specimens of 10-15 mg were uniformly placed into alumina crucibles of 70 lL volume and kept isothermally at 30 °C for 15 min to stabilize the sample.Dynamic runs were carried out by scanning the specimen from 30 to 180 °C with a temperature ramp of 3 °C min -1 at atmospheric pressure.During subsequent isothermal   MeOH methanol, EtOH ethanol.Calculation of theoretical stoichiometric molar ratios at standard conditions (20 °C, 1 atm) using the respective molar masses and densities of the different substances stabilization at the final temperature, the mass was recorded until the reaction went to completion.The TGA/DSC instrument was purged with N 2 at a flow rate of 50 mL min -1 during the entire measurement.Effects of buoyancy forces and temperature changes on the TGA and DSC signal were eliminated by automatic blank curve correction.For subtraction of the blank curve from the measurement curve, a blank curve was recorded under the same temperature conditions as the measurement curve, but using empty reference and sample crucibles.All measurements were performed under well-controlled laboratory conditions.The TGA/DSC device was calibrated with high purity metal standards (gallium, indium, lead, aluminum, and gold) over the temperature range of interest.
For temperature and enthalpy calibration, the calibration substances were subjected to the same temperature and heating conditions as the samples analyzed.The measured onset temperature, which is assigned the start of the melting process, was compared with the reference melting point of the respective standard.The determined enthalpy of fusion was also validated by comparison with the reference value.The blank curve reproducibility was better than ± 10 lg over the whole temperature range.The standard deviation of the enthalpy reproducibility was given as \ 5%.Mass changes and heat powers were determined with a precision of ± 0.1 lg and ± 1 mW.The Mettler Toledo STARe software 11.00a was used for data processing.
Selected samples of the series of salt-alcohol solvates were characterized complimentary by Raman spectroscopy, TG-MS, and XRD.In case of the first two methods, N 2 atmosphere was used for sample handling.

Raman spectroscopy
Raman spectra of samples in closed glass ampoules held at room temperature were recorded with a dynamically aligned Bio-Rad (Digilab) dedicated FT-Raman spectrometer equipped with a Spectra Physics Nd-YAG laser (1064 nm) and high sensitivity liquid-N 2 -cooled Ge detector.The excitation laser power used was about 250 mW at the samples.Increase in laser power up to 500 mW does not result in a significant change of the spectra that indicates the stability of the samples during the measurement.The resolution of the Raman instrument was ca. 4 cm -1 , and a backscattered geometry was used.For each spectrum, 256 individual spectra were averaged.

Thermogravimetric-mass spectrometric evolved gas analysis
The simultaneous thermogravimetric and mass spectrometric evolved gas analyses (TG-MS) were recorded on a Setaram LabsysEvo thermal analyzer, in high purity (99.9999%) helium atmosphere, with a flow rate of 80 mL min -1 .The measurements were done with a heating rate of 20 °C min -1 ; the samples were weighed into 100 lL aluminum crucibles in inert (dry N 2 ) atmosphere.Exposure to moisture during sample transfer was prevented, and MS recording of evolved volatiles was enabled by closing the crucibles with aluminum lids pierced with a 700 micron hole by crimping.The measurements were performed in the 25-500 °C temperature range.The obtained results were baseline corrected and then evaluated with the thermal analyzer's processing software (AKTS Calisto Processing, ver.1.41).Parallel with the thermogravimetric measurements, the analysis of the evolved gases/volatiles was carried out on a Pfeiffer Vacuum OmniStar TM gas analysis system coupled to the abovedescribed TGA.The gas splitter and transfer line to the mass spectrometer were preheated to 250 °C.The scanned m/z interval was 5-79 amu, with a scan speed of 20 ms amu -1 .The mass spectrometer was operated in electron impact mode.During a measurement, the total ion current (TIC), the discrete ion current of all scanned masses (75 masses) and the analog spectra on each scan cycle (1 scan cycle was 1.5 s long) were obtained in parallel.

X-ray powder diffraction
XRD patterns were obtained in a Philips model PW 3710 based PW 1050 Bragg-Brentano parafocusing goniometer using CuK a radiation (k = 0.15418 nm), graphite monochromator and proportional counter.Samples were placed in an Anton-Paar HTK2000 high-temperature oven chamber.Prior to measurements, the chamber was flushed with high purity N 2 (99.9999%) and during XRD measurements a continuous slow N 2 flow was ensured.The temperature steps chosen for the ''temperature programmed'' XRD were based on the results of the TG-MS measurement of the appropriate sample.First, XRD patterns were recorded at room temperature, which was followed by stepwise in situ heating.The heating rate was 5 °C min -1 .XRD patterns were recorded at every temperature step.Finally, the samples were allowed to cool down (20 °C min -1 ) and the room temperature XRD pattern was again measured.For the phase analysis, reference cards from the ICDD PDF-4 (2010) database were used.

Preliminary results of the direct synthesis of saltalcohol solvate
During the formation of the different salt alcoholate solvates, heat was evolved and resulted in a change in the temperature.The temperature change in the sample and the surrounding vapor phase was calculated from the difference between the respective initial and final temperatures.The initial sample temperatures ranged from 24 to 27 °C.The increase in temperature in both the sample and vapor chamber laid in the interval of 15-25 °C for almost all salt-alcohol pairs (Fig. 2).Rather fast reactions occurred and maximum temperatures of 40-56 °C were reached within a time of less than 3 min.In general, the temperature rise was higher in the solid phase than in the vapor phase, except for some outliers.For CaCl 2 -alcohol solvates, higher temperature differences were measured when compared with that for MgCl 2 -alcohol solvates.Since the temperature difference of both CaCl 2 -MeOH and CaCl 2 -EtOH solvates varied between 20 °C and 25 °C, irrespective of the alcohol/salt molar ratio, similar stoichiometric compositions of the different salt-alcohol solvates were considered.It is assumed that some heat dissipated into the ambient.
Lower temperature differences compared with that of CaCl 2 hydrates [36,37] indicate lower heat outputs and corresponding enthalpies of reaction.Heterogeneous compounds were observed particularly for samples of higher stoichiometry.Some compositions consisted of two phases: a solid salt alcoholate phase and an excess liquid alcohol phase containing dissolved salt particles.Thus, further analysis was carried out using the solid salt-alcohol solvates only.

Results of TGA/DSC measurements of all samples
The composition of the different salt-alcohol solvates synthesized according to Table 1 was analyzed by TGA/ DSC technique.The respective alcohol/salt molar ratios were derived from the change in sample mass recorded as a function of both time and temperature.

Thermal analysis of the CaCl 2 -MeOH system
The total amount of MeOH uptake of CaCl 2 was calculated to be 4 (Table 2).Samples decomposed in 2-3 overlapping stages over a temperature domain of 33-145 °C (Fig. 3).The majority of MeOH was given off below 100 °C.Peak temperatures of 57-94 °C, 80-114 °C, and 95-118 °C were measured (Table 2).As depicted in Fig. 3, the mass loss rate of the first decomposition step, also referred to as rate of dealcoholation, increased with increasing stoichiometric ratio despite similar sample masses.At theoretical alcohol/salt ratios of 2 and 3, a compound doubly   [27,[38][39][40][41][42].With excess of absolute MeOH, even a CaCl 2 trimethanolate and CaCl 2 tetramethanolate could be synthesized (Table 2).Note that samples prepared at higher theoretical alcohol/salt ratios were a heterogeneous mixture of two phases: a solid phase and a liquid phase.The latter was discarded from analysis.The formation of CaCl 2 -MeOH complexes of different stoichiometry has been claimed by other researchers.Gmelin [43] has reported the existence of CaCl 2 monomethanolates and CaCl 2 trimethanolates that have been identified in a study conducted by Gerhold and Kahovec.Bonnell [44] and Menschutkin [45] have obtained CaCl 2 trimethanolates, too.MeOH solvates of CaCl 2 with a molar ratio of 4 [43][44][45][46][47][48] and 6 [49] have been characterized by different analysis methods.The standard reaction enthalpies were deduced from the DSC curve by peak integration and varied between 101 and 151 kJ mol -1 (Table 2).

Thermal analysis of the CaCl 2 -EtOH system
In CaCl 2 -EtOH solvate complexes (Fig. 4), the EtOH was evolved in up to four inseparable steps with peak temperature around 68-72 °C, 87-92 °C, 108-121 °C, and 150-177 °C, respectively (Table 2).The decomposition started roughly at 33 °C and was completed at around 178 °C.The major amount of EtOH was desorbed below  1), (measurement temperature interval: 30-180 °C; heating rate of 3 °C min -1 ; 20 min isotherm at 180 °C) 100 °C.The rate of deethanolation varied over a broad range and was not reproducible, as shown in Fig. 4. The calculated percentage of mass loss (Table 2) clearly showed that in each sample only 2 mol of EtOH was adsorbed per mole of anhydrous CaCl 2 .However, a variation in the standard enthalpy of reaction with values between 92 and 130 kJ mol -1 was observed.We suppose that CaCl 2 could only hold up to two molecules of EtOH under the experimental conditions studied, irrespective of the initial stoichiometric ratio we used during sample preparation.These data are consistent with the findings described in the literature [27,50].According to published literature data, EtOH forms also a number of other CaCl 2 ethanolates.Monoethanolates [43,50,51], triethanolates [43-45, 52, 53], and tetraethanolates [53][54][55] of CaCl 2 have been reported previously.Their existence was not proven in this study.

Thermal analysis of the MgCl 2 -MeOH system
The experimental alcohol/salt ratios of the MgCl 2 methanolates prepared, as listed in Table 1, varied significantly from the theoretical ones.Methanolates of lower methanolation states were formed when the absolute MeOH was added in excess under the conditions studied.

Energy analysis
Variation in the coordination numbers of the salt-alcohol solvates cited in the references and this study could be the result of the diversity in the synthesis procedures and conditions.According to Bart and Roovers [70], the synthesis method plays a crucial role with regard to the nature of the reactants.Salt-alcohol solvates can be prepared by direct synthesis, solution crystallization and elimination of excess solvent, or recrystallization, for example.The alcohol coordinates to the alkaline earth metal cations to form complexes with variable stoichiometry and structural properties that exist either in the solid state or in solution.
Note that the calculated enthalpies of reaction and dissociation (Table 2) tend to be inaccurate as no quantitative differentiation between alcohol molecules and possible H 2 O molecules coordinated to one molecule of salt could be made at this stage.Samples contained probably trace amounts of H 2 O that were neglected, since the freshly A linear regression analysis of the experimentally obtained standard enthalpies of dissociation as a function of the number of alcohol molecules revealed that the applied model fitted well the data (Fig. 7).The coefficient of determination was close to one for each reaction system.Evidently, the standard enthalpy of dissociation increased along the homologous series of the alcohols studied.EtOH solvates possessed higher enthalpies of formation compared to MeOH solvates.The Ca 2? ion is of higher charge than the Mg 2? ion and has a greater polarizing power.Therefore, the standard enthalpies of formation of CaCl 2alcohol complexes are comparatively higher.
The best fit (R 2 = 0.9981) was achieved with the CaCl 2 -EtOH reaction system confirming its chemical stability.The calculated enthalpies of dissociation for all CaCl 2 -2EtOH solvates analyzed were in the range of 1303-1394 kJ mol -1 (Table 2) and coincided with formation enthalpy and dissociation enthalpy data given in the published literature [27,50] (Table 3).CaCl 2 -EtOH complexes of higher ethanolated states as reported by Parker et al. [53] and those of lower ethanolated states [50,51] could not be synthesized under the experimental conditions applied.
The CaCl 2 -MeOH system exhibited an acceptable fit with a coefficient of determination of R 2 = 0.9967.The enthalpies of dissociation of the CaCl 2 -2MeOH solvates with values of 1314 kJ mol -1 and 1337 kJ mol -1 , respectively, (Table 2) were slightly higher than literature values [27,38,42,51] (Table 3).Enthalpies of reaction and formation/dissociation for CaCl 2 -3MeOH and CaCl 2 -4MeOH have not been determined yet.The enthalpy of dissociation of CaCl 2 -3MeOH and of CaCl 2 -4MeOH was estimated at 1520 kJ mol -1 and 1771 kJ mol -1 / 1754 kJ mol -1 , respectively.The enthalpy of dissociation of CaCl 2 -MeOH complexes increased with the number of alcohol molecules indicating a linear relationship.However, a nonlinear relationship could be observed when the standard enthalpies of reaction per mole of alcohol D r H 0 z were compared (Table 2).At levels of methanolation higher than 2, the enthalpies of reaction per mole MeOH varied between 31 and 38 kJ mol -1 (Table 2), which are close to the standard enthalpy of vaporization of MeOH (37.965 kJ mol -1 ) [67].Apparently, up to 2 molecules could be chemically bound as a ligand.In MeOH solvates of CaCl 2 with higher molar ratios, the inclusion of MeOH molecules or physical incorporation of molecules into the crystal lattice could have caused the decrease in the standard enthalpy of reaction per mole of MeOH.
A linear relation between the enthalpy of dissociation and the number of alcohol molecules was also found for MgCl 2 -EtOH and MgCl 2 -MeOH solvate complexes (Fig. 7).The R-squared values of the MgCl 2 -EtOH and MgCl 2 -MeOH systems were R 2 = 0.9979 and R 2 = 0.9971, respectively.Reference data are rarely available.Iyimen-Schwarz [51] has collected the enthalpies of formation and dissociation of MgCl 2 -1EtOH, MgCl 2 -1MeOH, and MgCl 2 -3MeOH solvates from thermal cycling tests by DSC technique (Table 4).When compared with values determined in this study, a significant deviation appeared.The values given by Iyimen-Schwarz [51] are mean values averaged over 13 cycles for MgCl 2 -MeOH solvates and 10 cycles for MgCl 2 -EtOH solvates.Maximum formation and dissociation enthalpy values of 31 kJ mol -1 /-34 kJ mol -1  and 111 kJ mol -1 /--103 kJ mol -1 were derived for MgCl 2 -1MeOH and MgCl 2 -3MeOH.In the present study, the enthalpies of dissociation of MgCl 2 -1MeOH, MgCl 2 -2MeOH, MgCl 2 -3MeOH, and MgCl 2 -4MeOH ranged from 891 to 1639 kJ mol -1 (Table 2).The calculated enthalpies of reaction per mole MeOH of 43-77 kJ mol -1 (Table 2) were higher than the standard enthalpy of vaporization of MeOH.However, the enthalpies of reaction per mole EtOH with values of 30-45 kJ mol -1 (Table 2) were lower in comparison with the enthalpy of vaporization of EtOH (42 kJ mol -1 ) [67].Iyimen-Schwarz [51] has obtained even lower enthalpies of formation and dissociation and, respectively, enthalpies of reaction per mole of EtOH for MgCl 2 -EtOH solvates (Table 3).These inconsistent and unreliable experimental results are ascribed to inherent instability issues of the MgCl 2 -alcohol solvates.
To assess the suitability of the various salt-alcohol solvate systems for low-temperature heat storage, the gravimetric energy density was calculated from the experimentally measured standard enthalpies of reaction.The gravimetric energy density is a thermochemical material characteristic.As a key performance metric, it is used to evaluate and compare the energy storage performance of thermal energy storage systems.The gravimetric energy density of the investigated salt-alcohol systems was similar, irrespective of the reaction pair combination and ranged from 354 to 943 kJ kg -1 , respectively.From the energetic point of view, salt-alcohol solvate systems for heat storage are not as good as salt-water solvate systems, as the measured standard enthalpies of reaction and associated gravimetric energy densities are roughly half the respective values of the latter ones [76].

Raman spectroscopic characterization of selected samples
In a starting series of experiments, we increased the laser energy systematically from 200 mW to 500 mW and we found that spectra of sample 6 were not influenced by the change of the excitation energy.This observation let us to conclude that the alcohol solvate sample was stable enough For the calculation of missing enthalpy data, a standard enthalpy of formation of -641 kJ mol -1 for solid MgCl 2 , of -201 kJ mol -1 for gaseous MeOH and of -235.1 kJ mol -1 for gaseous EtOH was used, respectively [82] to be characterized by Raman spectroscopy.Raman spectra of several CaCl 2 -alcohol solvates recorded at room temperature can be seen in Fig. 8. Comparing the spectrum of pure CaCl 2 , (line d in Fig. 8a) to those of the different alcohol solvates (lines a, b, e, f in Fig. 8), it could be suggested that the very low wavenumber region was characteristic for CaCl 2 and bands from * 450 to * 3100 cm -1 belonged to the alcohols.Although Raman spectroscopy is not actually sensitive to H 2 O, a very weak band at about 3400 cm -1 indicated that CaCl 2 contained a certain amount of coordinated H 2 O (line e in Fig. 8b).As a comparison, the Raman spectrum of moisture exposed CaCl 2 is also shown (line h in Fig. 8b).
The MeOH in sample 1 (line a in Fig. 8a) was mainly in built-in form as m as of CH 3 at 2952 cm -1 and m s of CH 3 at 2844 cm -1 and was shifted with * 10 cm -1 comparing to those of the MeOH (2942 cm -1 and 2833 cm -1 , line c in Fig. 8a).Bands at 3013 cm -1 and 2982 cm -1 in sample 1 indicated the presence of coordinated OCH 3 (line a in Fig. 8a), while the band at * 3400 cm -1 indicated the presence of certain built-in H 2 O (cf. line a and line d in Fig. 9a) that was somewhat more intensive than in the starting salt.As the Raman technique gives average information about the material, we could not decide, whether the sample consisted of the mixture of non-coordinated, OH-coordinated, and MeOH-coordinated units of the CaCl 2 or units, which had MeOH and H 2 O in the same coordination sphere.
Contrary to sample 1, sample 4 contained a large amount of free MeOH besides the coordinated MeOH witnessed by the doublet of m s CH 3 band at 2844 cm -1 and 2836 cm -1 ; the latter belongs to non-coordinated MeOH.In addition, a double or triple band in the region * 3300 cm -1 indicated that -OH and non-coordinated MeOH existed in this sample (cf.lines b and c in Fig. 8a).The appearance of a significant amount of noncoordinated MeOH in sample 4 (prepared at an alcohol/salt ratio of 6:1) confirmed that the maximum number of the coordinated MeOH could be 4, as it was already indicated previously (Sect.Wavenumber/cm -1 Raman intensity/arb.units Wavenumber/cm -1 Raman intensity/arb.units Regarding the alcohol solvates obtained from CaCl 2 and EtOH, no real difference between the Raman spectra of the samples prepared at different alcohol/salt ratios could be found (cf.line e and f in Fig. 8b) in accordance with the findings described in Sect.3.1.Certain relative intensity changes in 2973 cm -1 / * 2878 cm -1 of m as CH 3 /m s CH 3 in the region of EtOH (cf.lines e, f and g in Fig. 8b) indicated the presence of the EtOH incorporated into the crystal lattice.The shift of the mC-O bands at 1096 cm -1 and 1051 cm -1 to 1088 cm -1 and 1048 cm -1 , respectively, could support this idea.However, the appearance of a weak band at 3396 cm -1 confirmed again the presence of coordinated H 2 O.
Raman spectra of several MgCl 2 -MeOH solvates are depicted in Fig. 9.Samples in these series were very similar to each other.The splitting of the mC-O band at 1033 cm -1 indicated MeOH incorporation into the crystal structure.The intensity ratios of MeOH bands/Mg-Cl bands increased with the amount of introduced MeOH.The m as of CH 3 appeared at 2950 cm -1 and m s of CH 3 appeared at 2848 cm -1 in the MgCl 2 -MeOH solvates.These values were shifted compared to that of free MeOH (cf.lines a, b, c and d in Fig. 9b), which indicated that MeOH definitely existed in built-in form in the samples 10, 12, and 13.
Regarding the MgCl 2 -EtOH systems (Fig. 10), the spectra resembled the spectrum of EtOH showing only negligible band shifts; it was difficult to decide, whether the EtOH existed in a real built-in form or not (cf.line a and c in Fig. 10a).Surprisingly, a new C-H stretching band at high wavenumber appeared (better visualized for sample 18, and only as a shoulder for sample 15) that might be assigned to mC-H of halogen substituted methyl group.As a comparison, the library IR spectrum of ethyl chloride is also shown (Fig. 10b).It seems plausible that ethyl chloride was formed in MgCl 2 -EtOH systems.Simultaneously, the Ca-Cl vibrational band at 241 cm -1 almost completely disappeared (cf.line b in Fig. 10a).
As a conclusion, MgCl 2 -EtOH solvates underwent certain decomposition resulting in ethyl chloride formation.Pure salt-alcohol solvates could not be obtained, because of the presence of a certain amount of salt hydrates.

TG-MS behavior of selected samples
As preliminary investigations, the starting materials, i.e., the CaCl 2 and MgCl 2, were analyzed, and it was found that both of them contained roughly around 2% H 2 O of crystallization (more precisely, CaCl 2 contained 2.28% H 2 O, while the H 2 O content of MgCl 2 was 2.36%).
Four samples were chosen for the TG-MS measurements, namely 3 (containing CaCl 2 and MeOH), 7 (which contain CaCl 2 and EtOH), 12 (containing MgCl 2 and MeOH), and finally 18 (composed of MgCl 2 and EtOH).In Fig. 11, the mass loss and some selected ion currents are plotted against the temperature obtained from sample 3, while on the right side of the figure the chemical species corresponding to the chosen m/z values are also shown.On the TG curve, between 55 and 300.5 °C, three mass loss steps can be seen.The first mass loss step is between 55 and 123.5 °C, with a mass loss of 23.3%, the second step is between 123.5 and 178 °C, with 19.2% mass lost, and the last step between 178 and 300.5 °C resulting in 9.5% mass loss; the total mass loss during the measurement was 52.9%.Comparing the shape of the three ion currents, three overlapping peaks can be observed on each curve, which correspond to the appropriate mass loss step.The m/z 32 is the molecular ion of MeOH, the m/z 31 is the base peak of MeOH (and also a characteristic marker of aliphatic alcohols), while the m/z 15 corresponds to the methyl ion (CH 3 ).From the TG-MS measurement, it can be concluded that the evolution of MeOH starts at very low temperatures (around 30 °C).Some H 2 O was also detected during the measurement (the ion current of H 2 O is not shown), but its concentration variation was within one order of magnitude compared to the concentration variation of MeOH, which was 2 orders of magnitude (m/z 31), thus confirming that the major volatile component formed during the measurement was MeOH.Moreover, no chlorinated compounds (e.g., methyl chloride) were detected.It can be seen that above 300 °C all three ion current curves are still decreasing, which implies that some MeOH is still lost, but this causes a very small mass loss, not detected by TG.
Considering the TG-MS trace of the sample 7 (Fig. 12), only two mass loss steps can be distinguished.The first mass loss step is between 55 and 95 °C, with a mass loss of 1.9%, while the second step is between 95 and 241 °C, with a mass loss of 48%; the total mass loss during the measurement was 50.3%.It can be seen that the formation of EtOH vapors (m/z 46 is the molecular ion of EtOH, m/ z 45 is a deprotonated EtOH, while m/z 31 is the base peak of EtOH) begins at very low temperatures, as already observed in the case of sample 3.
A very small amount of H 2 O can be detected between 152 and 235 °C, which means that the H 2 O is much more strongly bound to the CaCl 2 than the EtOH.No alkyl chloride was detected proving the thermal stability of the CaCl 2 -alcohol solvate.Comparing the results of sample 3 and 7, it can be concluded that EtOH (sample 7) is lost in a narrower temperature range than MeOH (sample 3).Moreover, the temperature value, where the alcohol is practically lost decreases from 300.5 °C (sample 3) to 240 °C (sample 7).This means that during cyclic measurements, EtOH can be liberated with less energy than MeOH.
Replacing CaCl 2 with MgCl 2 , significant differences appear in the TG-MS trace of sample 12 and 18.On the mass loss curve of the sample 12 (Fig. 13), three mass loss steps can be identified.The first larger step is between 55 and 198 °C, with a mass loss of 40%, the second, smaller mass loss is between 198 and 306 °C (10% mass lost), and the last, very small step is from 306 °C up to the end of the measurement (up to 500 °C, 1.2% mass loss).Evaluating the ion current curves of four masses (m/z 32 molecular ion of MeOH, m/z 31 base peak of MeOH, while m/z 52 is the molecular ion of methyl chloride, with the 37 Cl isotope, and m/z 50 is the molecular ion of methyl chloride, with the 35 Cl isotope), it can be seen that MeOH is liberated below 180 °C.Above this temperature, besides MeOH, which is still the major component of the volatiles and some traces of H 2 O, the formation of methyl chloride begins.This confirms that at higher temperatures, MgCl 2 hydrolyzes and converts the MeOH into methyl chloride.Around the upper end of the temperature scale, a small amount of methyl chloride is still released between 410 and 490 °C (0.7% mass is lost).Replacing the MeOH in the magnesium salt with EtOH (sample 18, Fig. 14), again a complicated, multi-step decomposition pattern is obtained.On the mass loss curve of sample 18, four decomposition steps can be identified (first step between 50 and 100 °C, mass loss 4.2%; second, larger step between 100 and 170 °C, mass loss 46.3%, third step between 170 and 212 °C, mass loss 13.3% and fourth step between 212 and 288 °C, mass loss 10.3%).It can be seen that below 140 °C, the mass loss is caused mainly by the evaporation of EtOH (m/z 46 is the molecular ion of the EtOH, m/z 45 is a deprotonated EtOH, while m/z 31 is the base peak of EtOH) and some traces of H 2 O (ion curves not shown).Above 140 °C, the formation of ethyl chlorides (m/z 66 is the molecular ion of ethyl chloride, with the 37 Cl isotope, while m/z 64 is the molecular ion of ethyl chloride, with the 35 Cl isotope) can be detected.Ethyl chloride is formed over a broad temperature range (between 140 and 370 °C), additionally the concentration of ethanol decreases above 290 °C, while the ethyl chloride is still formed.The formation of ethyl chloride confirms the fact that MgCl 2 hydrolyzes and converts the EtOH into the corresponding alkyl chloride.Comparing the TG-MS results of sample 12 and 18, it can be seen that the alkyl chloride formation starts at lower temperatures in the case of sample 18 (140 °C).The end temperature value of the mass loss end is shifted toward lower temperatures (mass loss end temperature of sample 12 is 306 °C, while the corresponding value is 288 °C in the case of sample 18).

Temperature programmed XRD characterization of selected samples
Figure 15 shows the XRD patterns of four selected saltalcohol solvate samples recorded by use of stepwise heating.The temperature steps were also chosen according to the TG curve obtained from the TG-MS measurement.
Sample 3 in its starting form (Fig. 15) consisted of MeOH-coordinated phases.The XRD diffraction pattern was somewhat more complicated at 55 °C than at 25 °C, which indicated the formation of a more complex system by the mild heating.These samples consisted of certain CaCl 2 xMeOH (H 2 O) phases.The most likely composition was the average CaCl 2 /MeOH/H 2 O = 1:1:3 ratio in the coordination sphere, which could be estimated from TG.Either a mixture of 1:4 and 1:2 coordinated samples of CaCl 2 -H 2 O and CaCl 2 -MeOH or samples with mixed coordination spheres could be imagined.At 120 °C, 177 °C, and 301 °C samples were CaCl 2 -like materials having amorphous part.At 500 °C, only the CaCl 2 phase existed.The sample after re-cooling was CaCl 2 with some Ca-hydroxide-carbonate.The XRD patterns at 120 °C, 177 °C, and 301 °C were very similar in spite of the large amount of MeOH that went away in the above temperature steps, which implies that a significant amount of MeOH was not really bound in the coordination sphere of the salt.
The diffraction pattern of sample 7 (Fig. 15b) did not change significantly during the heating up to 53 °C in accordance with the TG-MS results, which shows that no significant amount of EtOH or H 2 O was removed during the first some minutes.The starting solvate formed at 1:3 molar ratio could be imagined as a mixture of crystals with coordination sphere 1:2 and 1:4 or crystals with coordination sphere 1:2 and alcohol inclusions.The situation was complicated with the presence of H 2 O. Consequently, besides the pure CaCl 2 ÁzEtOH and pure CaCl 2 ÁzH 2 O crystals, CaCl 2 Á(x-y)EtOHÁyH 2 O structures could exist.Based on the XRD results it was impossible to give a more detailed description.Pure CaCl 2 phase could also be found in the region of 55-500 °C.Probably easily decomposable alcoholic, a relatively stable alcoholic and hydrous phases existed parallel in this sample.The final state was alcoholand H 2 O-free CaCl 2 .
The composition of the initial sample 12 was rather complicated (Fig. 15c).A volatile part was removed at 56 °C that resulted in a slight change in the diffraction pattern.One might think that physisorbed solvent was removed during the heating from 25 up to 56 °C and the solid consisted of mainly several H 2 O-coordinated MgCl 2 and MeOH-coordinated MgCl 2 forms.Regarding the starting solvate, the TG-MS indicated parallel removing of MeOH and a significant amount of H 2 O.The TG-MS gave a MgCl 2 /MeOH ratio of 3.12 only instead of 4. Since the molar mass of H 2 O is smaller than that of MeOH, the difference might come from the presence of structures with coordination number 4, but both ligands were in the coordination sphere in different variation.Consequently, the starting solvate itself had to be a mixture.The slight changes in the XRD pattern of sample 12 at 56 °C compared to that at 25 °C resulted from the small variation of   [77] and MgCl 2 Á6H 2 O system.Although the transformation of the MgCl 2 hydrates has long been known [78], clarifying of its mechanism is still in the focus of interest [79].
Possible problems of the applications of saltalcohol solvates in heat storage systems, based on the results of various techniques used Our results revealed that the MgCl 2 -EtOH solvates were instable compound; decomposition of these samples during storage, handling, and analysis was assumed.Raman spectroscopic measurements proved the appearance of ethyl chloride from MgCl 2 -EtOH without heating.Upon heating, both MeOH and EtOH solvates of MgCl 2 were involved in alkyl chloride release, as proven by TG-MS measurements.The alkyl chloride ion current increased slightly from the beginning of the heating and intensified above 140 °C.Although the heating rate can influence on the mechanism of the decomposition, TG-MS measurements by use of high heating rate indicated the formation of alkyl chloride as well as Raman spectroscopy without any heating.On the other hand, XRD measurements showed transformation of MgCl 2 in all MgCl 2 -alcohol solvate samples, which released alkyl chloride.
These observations were in accordance with the literature.Micro-calorimetric analysis has been conducted by Iyimen-Schwarz [51], who has made pioneering scientific contributions to the usability of salt-alcohol solvates based on CaCl 2 /MgCl 2 and MeOH/EtOH for thermal energy storage.Iyimen-Schwarz [51] calculated the energy density from the forward and reverse reaction enthalpy determined in dynamic DSC measurements under vacuum at controlled temperatures and alcohol vapor pressures.Based on observations on both the shape of the MgCl 2 -MeOH solvate's measurement curve and the change of the physical appearance, Iyimen-Schwarz suspected the decomposition of the MgCl 2 -MeOH solvate and the release of CH 3 Cl with successive cycling, analogous to the reaction of MgCl 2 with H 2 O, affecting the reproducibility of the measurements.According to his work, EtOH can be desorbed easily from the MgCl 2 -ethanolate system.Iyimen-Schwarz further assumed that in this system the MgCl 2 is likely to decompose into C 2 H 5 Cl and HCl, due to unreproducible results of the measured dissociation enthalpy.Results of our TG-MS and Raman spectroscopic measurements fully support the conclusions of Iyimen-Schwarz and explain the poor cyclic stability of MgCl 2 -EtOH solvates [51].
For alkyl chloride formation in our system different pathways could be assumed.It has been reported that gas phase reaction of EtOH and HCl with formation of H 2 O over ZnCl 2 /Al 2 O 3 as a catalyst, is a suitable method for ethyl chloride preparation [80].According to a possible explanation, the starting material (MgCl 2 ) contains MgCl 2 ÁzH 2 O in the presence of traces of H 2 O. Heating the MgCl 2 hydrates results in MgOH x Cl x with probable formation of hydrochloric acid (HCl).At the same time, reaction of alcohols and HCl can lead to the formation of methyl chloride (CH 3 Cl), or ethyl chloride (C 2 H 5 Cl) via MgCl 2 acting as a catalyst.Deliberation of H 2 O carries on the decomposition of MgCl 2 .A direct interaction between the ligands (i.e., Cl -and alcohol) in the coordination sphere of the Mg 2? can also be assumed.This idea does not need the H 2 O to assist in the alkyl chloride formation, but it can lead to the gradual hydrolysis of MgCl 2 .Although we do not have direct evidence for the first or second pathway, we believe that the direct way is more likely, since ethyl chloride formation appeared even at room temperature, while HCl release from MgCl 2 Á2H 2 O was reported only at 167 °C [79].It is worth to note that the final temperature in the cyclic stability test of MgCl 2 ÁzEtOH was 180 °C [29].
The results presented in Sect.3.3-3.5 demonstrate that pure salt-alcohol solvates could not be prepared under technically applicable conditions; the samples contained at least traces of H 2 O. H 2 O could be introduced by both the preparation procedure and by the starting materials, despite the use of commercial absolutized solvents and N 2 atmosphere for the preparation.The starting salts also contained 1-2% of H 2 O, which resulted in the presence of salt hydrates besides the mixture of salt-alcohol solvates.The possibility of the formation of salt hydrates implies the possibility of HCl formation during the thermal treatment above 167 °C [81].In practical applications, an inert atmosphere cannot be maintained during the whole process, so that traces of H 2 O cannot be avoided.Additionally, the use of high purity grade H 2 O-free substances might be too expensive at technical scale.The appearance of side reactions and the release of HCl with cycling can contribute to corrosion of the reactor components.The decomposition of MgCl 2 -alcohol solvates results further in a degradation of the overall performance of the thermal energy storage system.In conclusion, the reaction system MgCl 2 ÁR-OH is not suitable for practical implementation, due to its instability and irreversibility.From the energetic point of view, this compound is also not favorable as the measured enthalpies of reaction and associated energy densities are lower than that of CaCl 2 -alcohol solvates and salt-H 2 O systems [76].

Conclusions
Reversible chemical reactions are highly efficient in terms of storage volume, storage period, and sensible heat losses to the environment compared to other energy storage technologies.Different CaCl 2 -and MgCl 2 -alcohol solvates (EtOH, MeOH) were synthesized and their suitability for heat storage was examined by employing combined thermogravimetric analysis and differential scanning calorimetry (TGA/DSC), spectrometric and spectroscopic analysis (TG-MS, Raman) methods as well as by using X-ray diffraction (XRD).Due to their chemical nature, the CaCl 2 -EtOH systems exhibited lower energy densities than CaCl 2 -H 2 O systems.Decomposition of MgCl 2 -EtOH solvates accompanied by ethyl chloride formation started already during storage.Upon heating, both MeOH and EtOH solvates of MgCl 2 were affected by alkyl chloride release, as proven by TG-MS measurements.Our results fully support the assumptions of Iyimen-Schwarz [51] and explain the poor cycle stability of MgCl 2 -EtOH solvates reported previously.We also demonstrated that pure saltalcohol solvates cannot be prepared under technically applicable conditions.Formation of salt hydrates implies the possibility of HCl formation during the thermal treatment.Appearance of side reactions and possible release of HCl with cycling may be conducive to corrosion.Conclusively, MgCl 2 -alcohol solvate systems are not recommended for heat storage, whereas CaCl 2 -alcohol systems are suitable demonstrating stable cycle performance.The information on the decomposition pattern and associated changes in the structural integrity of the salt alcoholates obtained in this study are essential for selecting and designing efficient thermochemical energy stores.The fundamental data on the heat content of the parent salt provide the basis for the development of new two-component thermochemical materials with advanced properties.
Acknowledgements Open access funding provided by MTA Research Centre for Natural Sciences (MTA TTK).Project no.TE ´T_12_DE-1-2013-0003 has been implemented with the support provided by the National Research, Development and Innovation Fund of Hungary, financed under the TE ´T_12_DE funding scheme.This study received funding from the German Federal Ministry of Education and Research (BMBF) within the framework of a bilateral research collaborative project between the Leuphana University of Lu ¨neburg and the Hungarian Academy of Sciences under grant agreement number 01DS14029.The authors thank Christina Apel for the preparation of salt-alcohol solvate samples.
Open Access This article is distributed under the terms of the Creative Commons Attribution 4.0 International License (http://creative commons.org/licenses/by/4.0/),which permits unrestricted use, distribution, and reproduction in any medium, provided you give appropriate credit to the original author(s) and the source, provide a link to the Creative Commons license, and indicate if changes were made.

Appendix 1: Calculations
The alcohol/salt molar ratios (levels of alcoholation) were derived from experimental data obtained by TGA.The percentage mass loss X is defined as the mass of alcohol m R-OH desorbed per unit mass of salt alcoholate m MX ZR- OH : The level of alcoholation z is defined as the ratio of the number of alcohol molecules evolved during the endothermic dissociation reaction n R-OH to the amount of anhydrous salt n MX and was calculated from the following equation: wherein m MX and m R-OH are the masses of the anhydrous salt and the alcohol desorbed, respectively.The principle of thermochemical heat storage using reversible gas-solid reactions is based on the conversion of thermal energy into chemical energy required to break the chemical bonds of the reactants.At constant pressure, the amount of heat energy that must be supplied to induce the decomposition reaction equals the endothermic heat of reaction, also designated enthalpy of reaction.An infinitesimal change in the temperature results in a change of the enthalpy by DC p dT.The enthalpy of reaction D r H under non-standard condition, in case of a reaction temperature T 1 different from the standard state temperature T 0 can be estimated from the standard reaction enthalpies and heat capacities of the reactants using Kirchhoff's law.The heat energy DQ is then expressed by: where DrH 0 (T 0 ) is the standard enthalpy of reaction at standard state conditions T 0 = 298.15K and p 0 = 1 bar, C p is the constant-pressure heat capacity and D Tr H is the enthalpy of transformation.
The enthalpy of reaction was obtained from DSC measurement by peak area integration.The energy that is liberated or absorbed as heat during a chemical reaction as a result of a temperature difference DT can be quantitatively determined by heat flux DSC.This technique measures the thermally induced heat flux transferred between the sample and an inert reference that are connected by a low-resistance heat flow path.Thermocouples below the symmetrically positioned sample crucible and empty reference crucible detect and compare the temperature of the specimen to the temperature of the reference as a function of time under same conditions.The heat flux between sample and reference is proportional to the temperature difference: wherein E (T) and DSC are the calorimetric sensitivity and the measured DSC signal.Integration of the peak area under the baseline-subtracted DSC signal over time yields D r H 0 : According to Hess Law, the value of the standard reaction enthalpy of the forward and reverse reaction must be equal and the same applies for the standard enthalpy of association and standard enthalpy of dissociation.The standard enthalpy of dissociation is the inverse of the standard enthalpy of formation.In general, the standard enthalpy of reaction is calculated from the difference of the standard enthalpy of formation of the products and the standard enthalpy of formation of the reactants: Solving Eq. ( 8) for D f H 0 products gives the standard enthalpy of formation and hence the standard enthalpy of dissociation.
Different mathematical approaches are used to determine the specific energy storage density.Two types of energy densities are known: gravimetric and volumetric energy density.The gravimetric energy density E m is defined as the capacity of heat energy stored at a defined temperature and pressure per unit mass of storage material and can be calculated from the ratio of the standard reaction enthalpy to the molar mass of the reactant M MX zR-OH : whereas the volumetric energy density E V is related to the volume of storage material and is described as: The storage material volume is derived from the material's mass and bulk density q.The higher the level of alcoholation, the lower is the density.The volumetric energy density is an important key energy storage metric for designing and operating storage systems.It is also preferred for performance comparison studies.Space can be a limiting factor for many practical applications.There are no data available for the mass and bulk densities of the salt-alcohol systems studied and hence only the gravimetric energy density could be calculated.
Appendix 2: Data processing-example of a TGA/DSC curve

Fig. 1
Fig. 1 Schematic diagram of the operating principle of a closed thermochemical heat storage system

Fig. 7
Fig.7Experimentally obtained standard enthalpies of dissociation of various salt-alcohol solvates plotted against the number of alcohol molecules evolved during the dissociation reaction, MeOH methanol, EtOH ethanol.Linear regression analysis was applied to determine the coefficient of determination (R-squared)

Fig. 12 Fig. 13
Fig. 12 TG-MS trace of the sample 7, on the right side the formulae of the corresponding fragments/ions

Fig. 14 Fig. 15
Fig. 14 TG-MS trace of the sample 18, on the right side the formulae of the corresponding fragments/ions

Figure 16
Figure16displays the DSC measurement curve and corresponding TGA signal normalized to the sample mass for sample 6.To determine the heat content of the sample and the associated enthalpy of reaction, the area under the DSC peak was integrated.Similar curves were recorded for all other samples.

Fig. 16
Fig. 16 Example of a recorded TGA and DSC curve.The TGA curve or change in mass (%) and DSC curve or heat flow (mW) are given as a function of time (min) and temperature (°C).The sample was heated Anhydrous CaCl 2 powder (Merck, Ph Eur), anhydrous MgCl 2 powder (Roth, C 98.5%), absolute MeOH (max.0.003% H 2 O, Merck), and absolute EtOH (max.0.01% H 2 O, Merck) were used as starting materials.The H 2 O content of pure CaCl 2 and MgCl 2 samples as received was determined by TGA.The powdered salts contained only traces of H 2 O. Alcohol-salt pairs in four combinations, i.e., MeOH/ CaCl 2 , EtOH/CaCl 2 , MeOH/MgCl 2 , EtOH/MgCl 2 , were prepared setting different alcohol/salt molar ratios (Table

Table
Preparation of salt-alcohol solvates by direct synthesis from liquid phase alcohols under neat conditions

Table 1 . MeOH methanol; EtOH ethanol; diamond- sample, square-vapor chamber
, gravimetric energy storage density.Step 1, step 2, step 3 and step 4 refer to the respective EtOH/MeOH release steps of the multi-stage dissociation reactions coordinated with MeOH was the only CaCl 2 -MeOH solvate found.This result is consistent with the findings of other authors initial , initial temperature; T final , final temperature; T Peak peak temperature; D r H°, reaction enthalpy; D r H°/z, reaction enthalpy per mole EtOH/MeOH; D f H°, enthalpy of formation; E m

Table 3
Data available in the literature on experimental and theoretical standard enthalpies of reaction per mole salt alcoholate D r H°and standard enthalpies of formation/dissociation D f H°of CaCl 2 -MeOH and CaCl 2 -EtOH complexes of different stoichiometry z [82]the calculation of missing enthalpy data, a standard enthalpy of formation of -796 kJ mol -1 for solid CaCl 2 , of -201 kJ mol -1 for gaseous MeOH and of -235 kJ mol -1 for gaseous EtOH was used, respectively[82]

Table 4
Data available in the literature on experimental and theoretical standard enthalpies of reaction per mole salt alcoholate D r H°and standard enthalpies of formation/dissociation D f H°of MgCl 2 -MeOH and MgCl 2 -EtOH complexes of different stoichiometry z r H°D f H°D r H°D f H°D r H°D f H°D r H°D f H°/ 3.1).It is worth noting that the Raman spectrum of sample 3 was very similar to that of sample 4, temperature region of 101-306 °C, the sample gradually lost the EtOH.According to TG-MS results, only EtOH was removed up to 171 °C; EtOH and a small amount of H 2 O went away between 171 and 306 °C.Nevertheless, ethyl chloride was also detected by TG-MS indicating a decomposition process.Increasing the temperature from 101 °C up to 212 °C resulted in a change in the XRD patterns with gradual formation of amorphous phases.MgOHCl, MgCl 2 (aq) probably existed at higher temperatures.The XRD pattern obtained at 500 °C indicated a cubic crystalline material.The pattern of MgO could be fitted well, but the coexistence of other cubic components could also be suggested.The ''temperature programmed'' XRD behavior of sample 18 was very similar to that of sample 12.Results of XRD measurements led us to the following conclusions.Mixtures of different phases containing alcohol and H 2 O existed in all of our salt-alcohol solvate samples.CaCl 2 was retrievable from its alcohol solvate by bake out of the alcohol solvate, but irreversible processes appeared in case of MgCl 2 -alcohol solvates.Upon increasing the temperature up to 500 °C, MgO was obtained with elimination of hydrogen chloride similarly to the MgCl 2 Á6H 2 OÁ1,4-C 4 H 8 O 2