RESUMEN

The aim of the work was to analyze the preferential solvation process, and determine the composition of the solvation shell of cyclic ethers using the calorimetric method. The heat of solution of 1,4-dioxane, 12-crown-4, 15-crown-5 and 18-crown-6 ethers in the mixture of N-methylformamide with water was measured at four temperatures, 293.15 K, 298.15 K, 303.15 K, and 308.15 K, and the standard partial molar heat capacity of cyclic ethers has been discussed. 18-crown-6 (18C6) molecules can form complexes with NMF molecules through the hydrogen bonds between -CH3 group of NMF and the oxygen atoms of 18C6. Using the model of preferential solvation, the cyclic ethers were observed to be preferentially solvated by NMF molecules. It has been proved that the molar fraction of NMF in the solvation shell of cyclic ethers is higher than that in the mixed solvent. The exothermic, enthalpic effect of preferential solvation of cyclic ethers increases with increasing ring size and temperature. The increase in the negative effect of the structural properties of the mixed solvent with increase in the ring size in the process of preferential solvation of the cyclic ethers indicates an increasing disturbance of the mixed solvent structure, which is reflected in the influence of the energetic properties of the mixed solvent.

Asunto(s)

Éteres Corona , Agua , Temperatura , Agua/química , Éteres Cíclicos , Éteres Corona/química , Solventes

RESUMEN

The solution enthalpy of 15-crown-5 and 18-crown-6 ethers in the mixture of formamide (F) and water (W) was measured at four temperatures: 293.15 K, 298.15 K, 303.15 K, 308.15 K. The standard molar enthalpy of solution, ΔsolHo, depends on the size of cyclic ethers molecules and the temperature. With increasing temperature, the values of ΔsolHo become less negative. The values of the standard partial molar heat capacity Cp,2o of cyclic ethers at 298.15 K have been calculated. The Cp,2o=f(xW) curve shape indicates the hydrophobic hydration process of cyclic ethers in the range of a high-water content in the mixture with formamide. The enthalpic effect of preferential solvation of cyclic ethers was calculated and the effect of temperature on the preferential solvation process was discussed. The process of complex formation between 18C6 molecules and formamide molecules is observed. The cyclic ethers molecules are preferentially solvated by formamide molecules. The mole fraction of formamide in the solvation sphere of cyclic ethers has been calculated.

RESUMEN

Predicting the equilibrium solubility of organic, crystalline materials at all relevant temperatures is crucial to the digital design of manufacturing unit operations in the chemical industries. The work reported in our current publication builds upon the limited number of recently published quantitative structure-property relationship studies which modelled the temperature dependence of aqueous solubility. One set of models was built to directly predict temperature dependent solubility, including for materials with no solubility data at any temperature. We propose that a modified cross-validation protocol is required to evaluate these models. Another set of models was built to predict the related enthalpy of solution term, which can be used to estimate solubility at one temperature based upon solubility data for the same material at another temperature. We investigated whether various kinds of solid state descriptors improved the models obtained with a variety of molecular descriptor combinations: lattice energies or 3D descriptors calculated from crystal structures or melting point data. We found that none of these greatly improved the best direct predictions of temperature dependent solubility or the related enthalpy of solution endpoint. This finding is surprising because the importance of the solid state contribution to both endpoints is clear. We suggest our findings may, in part, reflect limitations in the descriptors calculated from crystal structures and, more generally, the limited availability of polymorph specific data. We present curated temperature dependent solubility and enthalpy of solution datasets, integrated with molecular and crystal structures, for future investigations.

RESUMEN

The enthalpies of solution of the cyclic ethers 1,4-dioxane, 12-crown-4 and 18-crown-6 in mixtures of ethanol and water have been measured within the whole mole fraction range at T = 298.15 K. The enthalpy of solvation has been calculated. In pure ethanol and pure water, the solvation enthalpy of the investigated cyclic ethers depends linearity on the number of -CH2CH2- groups in the cyclic ether molecules. Based on the analysis of the preferential solvation model proposed by Waghorne, it can be concluded that the 1,4-dioxane, 15C5 and 18C6 molecules are preferentially solvated by water molecules in the range of low water content in these mixtures. The effect of base-acid properties of ethanol-water mixtures on the enthalpy of solution of cyclic ethers in these mixtures has been analyzed. The enthalpy of solution of cyclic ethers correlates with the acidic properties of ethanol-water mixtures in the range of high and medium water content. The results presented are compared with analogous data obtained for the methanol-water and propan-1-ol-water mixtures.

RESUMEN

In this work enthalpies of dissolution in water of polyethylene glycols (PEGs) having an average molecular weight of 1000 and 1400, Pluronic-F127, phenacetin as well as the composites prepared from them were measured using solution calorimetry at 298.15 K. Intermolecular interaction energies of polymer-phenacetin were calculated on the basis of an additive scheme. It was shown that for mixtures with high content of polymer (>90 wt%) Pluronic-F127 has the highest solubilizing effect, while for mixtures with (4-6):1 polymer: phenacetin ratio the best solubilizing agent is PEG-1400. Infrared-spectra showed a decrease of the number of self-associated molecules of phenacetin with increasing of polymer content in the composites. The obtained results enabled us to identify the features of intermolecular interactions of polymers with a model hydrophobic drug and may be used for optimizing the conditions for preparing solid dispersions based on hydrophilic polymers.

RESUMEN

The following values of the enthalphy of solution of well-characterized samples of guanine were obtained from measurements in an adiabatic solution calorimeter: [Table: see text] The following enthalpies of protonation were calculated for guanine: [Table: see text].

RESUMEN

The value for the enthalpy of solution of SRM 1655 (KCl), ΔH°(500 H2O, 298.15 K) = (235.86±0.23)J·g-1 or (17.584±0.017)kJ·mol-1, was obtained from measurements in an adiabatic calorimeter, and confirmed by measurements in an isoperibol calorimeter. Enthalpy of solution measurements are reported in the temperature range 296 K to 358 K, at molalities between 0.005 and 0.18 mol-kg-1, for various ranges of particle size up to 1000 µm, and in CO2- saturated solutions. Observations were also made relative to the hygroscopicity, removal of occluded H2O, and homogeneity of SRM 1655. Between 296 K and 303 K, ΔC p =-(2.076±0.087) J·g-1·K-1or -(154.8±6.4) J·mol-1·K-1 for the reaction of SRM 1655 in 500 H2O.

RESUMEN

Systematic errors in an isoperibol calorimeter of a widely-used design, amounting to about 0.5 percent of the endothermic enthalpy of solution of SRM 1655 (KCl) in H2O, were found to be the result of errors in heat leak corrections due to inadequate stirring and commonly used calorimetric procedures. Other systematic errors were found in measurements of the enthalpy of solution of the exothermic reaction of tris(hydroxymethyl)aminomethane in aqueous HCl solution. Recommended procedures are summarized.

RESUMEN

The following values were obtained from measurements of the enthalpy of solution of a well characterized sample of crystalline adenine in various solvents: [Table: see text] Using these measured values the following enthalpies of protonation and proton dissociation were calculated: [Table: see text].

RESUMEN

Enthalpy of solution measurements of four potassium salts in H2O were made in either an adiabatic or an isoperibol calorimeter or both. The following summarizes the measured and recommended values: [Table: see text] The value for KIO4 has been corrected for the hydrolysis of the periodate ion. The ΔC p = -(82.5 ± 4.3)J·mol-1·K-1 for the unhydrolysed reaction. For the reaction of KBr in H2O, ΔC p was measured as -(166.6 ± 7.2) J·mol-1·K-1 in the temperature range 298 K to 319 K. Measurements with the isoperibol calorimeter are also reported for the endothermic reaction of tris(hydroxymethyl)aminomethane, SRM 724a, in aqueous NaOH (0.05 mol·L-1). Comparisons of measurements by different calorimeters on the same samples reveal unidentified calorimetric errors for endothermic reactions which are greater than the imprecision of the measurements.

RESUMEN

An adiabatic solution calorimeter was used to measure the enthalpy of solution in water of various adenine samples for which a large amount of analytical information is reported. The experimental imprecision of the measurements was 1.1 percent. However, it was necessary to assign an overall uncertainty of 3 percent because of impurity uncertainties. Thus, the best value for the enthalpy of solution is Δ H ∘ ( ∞ , 298.15 K ) = ( 33.47 ± 1.00 ) k J m o l - 1 ΔC p = (78.7 ± 10.4) J·mol-1·K-1 in the range 298 to 328 K at 5 mmol·kg-1, and the enthalpy of dilution is -(316 ± 208) kJ·mol-1 (mol·kg-1)-1 in the range 1 to 7 mmol·kg-1. The entropy of solution for adenine was calculated to be ΔS° (298.15 K) = (72.1 ± 3.9) J·mol-1·K-1, and the partial molar heat capacity at infinite dilution, C° p2 = (226 ± 11) J·mol-1·K-1. The density of adenine was measured by displacement as 1.47 g·mL-1 with an estimated uncertainty of 1 percent.

RESUMEN

An adiabatic solution calorimeter was used to measure enthalpies of solution in water of 7 uracil samples in a concentration range of 3 to 45 mmol · kg-1 and over a temperature range of 298 K to 325 K. Analytical data for the samples are given. Our best value for the enthalpy of solution is Δ H 0 ( ∞ , 298.15 K ) = ( 29.33 ± 1.2 ) kJ mol - 1 and for the change in heat capacity for the reaction with temperature, Δ C p 0 = ( 57 ± 13 ) J mol - 1 K - 1 No change in the enthalpy of solution with concentration was found in this range. Values were calculated for the entropy of solution, ΔS° = (68.1 ± 4.2) J · mol-1 · K-1, and for the apparent molal heat capacity at infinite dilution, C p 2 0 = ( 178 ± 15 ) J m o l - 1 K - 1 .

RESUMEN

An adiabatic solution calorimeter was used to measure enthalpies of solution in water of 8 samples of thymine for which analytical data are reported. Our best values for the enthalpy of solution and the change in heat capacity are Δ H ∘ ( ∞ , 298.15 K ) = ( 24.32 ± 0.70 ) kJ ⋅ mol - 1 Δ C p ∘ = ( 106 ± 26 ) J ⋅ mol - 1 ⋅ K - 1 , 298 K < T < 328 K . These were used to calculate ΔS° = (53. 1 ± 3.6) J · mol-1 · K-1 for the entropy of solution, and C p 2 ∘ = ( 256 ± 26 ) J · mol-1 · K-1 for the apparent molal heat capacity at infinite dilution. No change in the enthalpy of solution with concentration was observed in the range of 5 to 35 mmol · kg-1.

RESUMEN

An adiabatic solution calorimeter was used for measuring enthalpies of solution in water of seven samples of cytosine for which analytical data are given. Our best values are: Δ H ∘ ( ∞ , 298.15 K ) = ( 27.2 ± 4.0 ) kJ ⋅ mol - 1 , and Δ C p = ( 76 ± 21 ) J ⋅ mol - 1 ⋅ K - 1 , 298 K < T < 324 K at ( 24 ± 2 ) mmol ⋅ kg - 1 . Evidence is given for an unidentified side reaction at low concentrations which is responsible for the large uncertainty assigned to the enthalpy of solution at infinite dilution. An approximate value of (1.44 ± 0.08) g · mL-1 for the density of cytosine was also measured. Values are calculated for ΔG° and ΔS° for the solution reaction.

RESUMEN

An adiabatic solution calorimeter was used to measure the enthalpies of the following reactions: BeO ( c ) + [ 96 HF + 338 H 2 O ] ( l ) → [ BeF 2 + 94 HF + 339 H 2 O ] ( l ) Δ H at 298.15 K = - 101.30 ± 0.20 kJ ⋅ mol - 1 BeO ( c ) + [ 48 HCL + 333 H 2 O ] ( l ) → [ BeCL 2 + 46 HCl + 333 H 2 O ] ( l ) Δ H at 352.58 K = - 54.19 ± 0.22 kJ ⋅ mol - 1 Δ H at 298.15 K = - 53.0 ± 2.0 kJ ⋅ mol - 1 Δ Cp is 12.5 ± 2.8 J ⋅ mol - 1 ⋅ K - 1 for the first reaction in the range 298 to 325 K. .

RESUMEN

The enthalpies of precipitation of silver halides and the enthalpies of solution of AgNO3, KCl, and KBr in H2O were measured in an adiabatic solution calorimeter. From the enthalpy measurements of KCl(c) and KBr(c) in AgNO3(aq), and of AgNO3(c) in KCl(aq), in KBr(aq), and in KI(aq), we calculated (in kJ · mol-1) -65.724, -84.826, and -111.124 for ΔH° pptn(298.15 K) for the averages of the chloride, bromide, and iodide reactions, respectively. A reevaluation of the data for the enthalpy of solution of AgNO3(c) has resulted in our selected best value, Δ H ° ( ∞ ) ( 298.15 K ) = 22.730 + 0.084 kJ ⋅ mol - 1 = 5.433 ± 0.020 kcal ⋅ mol - 1 A table of enthalpies of dilution of AgNO3(aq) is also given. The average standard entropy for the aqueous silver ion at 298.15 K is found to be S ° [ Ag + ( aq ) ] = 73.42 ± 0.20 J ⋅ mol - 1 ⋅ K - 1 = 17.55 ± 0.05 cal ⋅ mol - 1 ⋅ K - 1 . .