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Abstract  

A simple method for determination of the hydrate numbers of saturating multi-hydrate salts in developed. The method demonstrated for scandium sulfate is based upon estimation of the enthalpy of solution of the hydrates from the solubility smoothing equations. It is shown that in the Sc2(SO4)3–H2O system, contrary to common opinion, the equilibrium solid phases are: Sc2(SO4)3.6H2O at 273–295 K, Sc2(SO4)3.5H2O at 295–333 K and Sc2(SO4)3.4H2O at 333–373 K. The solubility smoothing equations for the hexa-, penta- and tetrahydrate of scandium sulfate are given.

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The enthalpies of crystallization of NaCl, KCl, LiCl·H2O, MgCl2·6H2O, CaCl2·6H2O and BaCl2·2H2O from aqueous solution were determined by means of different calculation methods on the basis of the earlier-measured differential and integral enthalpies of solution of the above salts. The obtained crystallization enthalpies are discussed and compared with the appropriate literature data.

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Abstract  

Assuming that the correlation between the cocrystallization coefficients and solubilities of the co-crystallizing ethylsulphates in the [(Ln,M) (H2O)9] (C2H5SO4)3–H2O system is valid when M is changed from lanthanides into the title elements, the solubilities of the ethylsulphates of trivalent Y, Pm, Pu, Am and Cm in water at 288–318 K have been determined from the matrix. The solubilities of Y, Pm, Pu, Am and Cm ethylsulphates and of all the lanthanide ethylsulphates are given in the form of smoothing equations of the lg molality=A+B/T type. From the B parameters of the solubility equations the enthalpies of solution have been estimated. The crystallization behaviour of yttrium in the ethylsulphate system is between that of holmium and that of erbium.

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The enthalpy of solution of KCl in H2O (1∶2000 mol·mol−1) at 298.15 K was measured in an interlaboratory test in the G.D.R. The test material was prepared in the ASMW laboratories. The purity found on the high-precision coulometric titration of chloride was 0.9999±0.0001 g·g−1. The consensus value of the enthalpy of solution in the test was ΔH s 298.15=17.47±0.07 kJ·mol−1. This result is in good agreement with experimental values from recognized international scientific laboratories. The test material is applicable as a CRM.

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Abstract  

The heats of solution of tetrabutylammonium bromide have been measured in mixtures of formamide (FA) with methanol (MeOH) and ethylene glycol (EG) at 313.15 K by calorimetric method. The standard enthalpies of solution in binary mixtures have been extrapolated to infinite dilution by Redlich–Rosenfeld–Meyer type equation using the literary data at 298.15 K and the present paper data at 313.15 K. The Debye–Hückel limiting law slope A H required for calculation of the ∆sol H 0 value has been obtained with application the new additive scheme of determination of the physic-chemical characteristics of binaries. The scheme is tested on the example of Bu4NBr solutions in FA–MeOH mixture at 298.15 K. Its application yields the ∆sol H 0 value very closed on the ones determined with the real (non-additive) characteristics of binaries. The standard enthalpies of solution extrapolated by Redlich–Rosenfeld–Meyer type equation are in a good agreement with the ones computed in terms of the Debye–Hückel theory in the second approximation. The heat capacities characteristics of Bu4NBr have been calculated in H2O–FA, MeOH–FA and EG–FA mixtures using the literary and present data. The sequence of solvents H2O > FA > EG > MeOH located on their ability to solvophobic solvation found by us earlier for enthalpic characteristics is confirmed by the ∆C p 0 values. The comparison of thermochemical characteristics of Bu4NBr solutions in aqueous and non-aqueous mixtures containing FA has been carried out. The own structure of water remains in the region of small additions of formamide to co-solvents. It considerably differs the H2O–FA mixture from the investigated non-aqueous systems.

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Abstract  

The standard molar enthalpies of solution at infinite dilution
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of glycylglycine, dl-alanyl-dl-alanine and glycylglycylglycine in aqueous solutions of potassium chloride and ethanol as well as of glycylglycine and glycylglycylglycine in the solutions containing urea and water have been determined by calorimetry at the temperature 298.15 K. Changes of solution enthalpy, expressed in a form so-called heterotactic interaction coefficients,
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were used for analysis of interactions occurring between the investigated solutes in water. The group contributions illustrating the interactions of KCl, urea and ethanol with selected functional groups in the peptide molecules, namely CH2, “pep,” and “ion” groups, were calculated and discussed.
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The enthalpies of solution in 2 N nitric acid at 298 K were measured for alkali metal borate glasses and crystals. From the data obtained, their enthalpies of formation from the oxides and the heats of crystallization of the glasses were calculated.

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Abstract  

The thermodynamic energy relationship between two crystal modifications of cimetidine was investigated and compared with differences in their processing properties with respect to transformation from one modification to the other.The crystal energies of the two modifications A and D were found to be almost identical and therefore the polymorphs are regarded as virtually isoenergetic crystals. This statement is based on DSC measurements of the melting points and of the enthalpies of fusion for the two crystal forms, which enable the calculation of the Gibbs free energy functions. Furthermore, the statement is supported by measurements of the enthalpies of solution in two different solvents. Both DSC and solution experiments reveal a slightly higher stability of the D modification with respect to the A form. In addition, tribomechanical treatment also indicates modification D to be the more stable one, as well as the higher density of the D form. No transformation during DSC at low heating rate was found which could be used in a stability consideration.As the explicit crystal structures of the two modifications are resolved, it was possible to calculate crystal energies theoretically as well. The theoretical results showed a remarkable difference in the crystal energies at zero degree Kelvin. Furthermore, they were just contradicting experimental findings by stating A being more stable than D. Possible reasons for this discrepancy and the feasibility of today's calculation methods with respect to prediction of stability properties are discussed.

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Abstract  

Enthalpies of solution of dibenzo-18-crown-6 at infinite dilution have been measured in tetrachloromethane, benzene, chloroform, pyridine, acetone, dimethylformamide, dimethylsulfoxide, acetonitrile, benzyl alcohol and propylene carbonate at temperature of 298.15 K. Values for enthalpy of solvation and solute–solvent interaction in the solvents were determined. Correlation of enthalpies of solvation with the enthalpy of cavity formation and contribution for the different types of solute–solvent interaction was obtained. In benzene, tetrachloromethane, pyridine, DMF and DMSO polar conformer of DB18C6 dominates. Its effective dipole moment amounts to 3.7 D. Conformation dynamics of the solute reduces the effective polarity of such solvents as acetone, chloroform and propylene carbonate in which population of polar conformer of dibenzo-18-crown-6 decreases. Condensation of two benzene rings to 18-crown-6 results in increasing molecule polarity and exothermic contribution of dipole–dipole interaction in polar solvent media. The specific interaction with acetonitrile and chloroform becomes weaker from 18-crown-6 to dibenzo-18-crown-6.

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The excess enthalpies were investigated for 12 binary and 4 ternary systems, including mixtures and solutions of electrolytes and non-electrolytes. The excess enthalpies of mixtures and integral heats of solutions were measured with an isoperibol calorimeter at 35 °C. Heats of fusion and heat capacities as functions of temperature were measured with a Perkin Elmer Corp., DSC-2. Integral heats, heats of fusion and heat capacities allow investigations of excess enthalpies of solutions. For modelling of the experimental results, the modified Redlich-Kister equation was used with good success.

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