Enthalpy formation of fluorene: Is a challenging problem for theory or experiment?

The large discrepancy between the experimental enthalpy of formation of fluorene and theoretical value calculated by G3(MP2) method was revealed more than ten years ago. Three years later, a new...

Experimental determination of the solubility constant of kurnakovite, MgB3O3(OH)5·5H2O

In this study, I present experimental results on the equilibrium between boracite [Mg3B7O13C1(cr)] and kurnakovite [chemical formula, Mg2B6O11·15H2O(cr); structural formula, MgB3O3(OH)5·5H2O(cr)] at 22.5 ± 0.5 °C from a long-term experiment up to 1629 days, approaching equilibrium from the direction of supersaturation, Mg3B7O13C1(cr) + H+ + 2B(OH)4− + 18H2O(1) ⇌ 3MgB3O3(OH)5·5H2O(cr) + C1−. Based on solubility measurements, the 10-based logarithm of the equilibrium constant for the above reaction at 25 °C is determined to be 12.83 ± 0.08 (2σ). Based on the equilibrium constant for dissolution of boracite, Mg3B7O13C1(cr) + 15H2O(1) = 3Mg2+ + 7B(OH)4− + C1− + 2H+ at 25 °C measured previously (Xiong et al. 2018) and that for the reaction between boracite and kurnakovite determined here, the equilibrium constant for dissolution of kurnakovite, MgB3O3(OH)5·5H2O(cr) = Mg2+ + 3B(OH)4− + H+ + H2O(1) is derived as −14.11 ± 0.40 (2σ). Using the equilibrium constant for dissolution of kurnakovite obtained in this study and the experimental enthalpy of formation for kurnakovite from the literature, a set of thermodynamic properties for kurnakovite at 25 °C and 1 bar is recommended as follows: ΔHf0 = −4813.24 ± 4.92 kJ/mol, ΔGf0 = −4232.0 ± 2.3 kJ/mol, and S0 = 414.3 ± 0.9 J/(mol·K). Among them, the Gibbs energy of formation is based on the equilibrium constant for kurnakovite determined in this study; the enthalpy of formation is from the literature (Li et al. 1997), and the standard entropy is calculated internally with the Gibbs-Helmholtz equation in this work. The thermodynamic properties of kurnakovite estimated using the group contribution method for borate minerals based on the sums of contributions from the cations, borate polyanions, and structural water to the thermodynamic properties from the literature (Li et al. 2000) are consistent, within their uncertainties, with the values listed above. Since kurnakovite usually forms in salt lakes rich in sulfate, studying the interactions of borate with sulfate is important to modeling kurnakovite in salt lakes. For this purpose, I have re-calibrated our previous model (Xiong et al. 2013) describing the interactions of borate with sulfate based on the new solubility data for borax in Na2SO4 solutions presented here.

Molecular Mobility and Stability Studies of Amorphous Imatinib Mesylate

The proposed study examined the characterization and stability of solid-state amorphous imatinib mesylate (IM) after 15 months under controlled relative humidity (60 ± 5%) and temperature (25 ± 2 °C) conditions. After 2 weeks, and 1, 3, 6, and 15 months, the samples were characterized using differential scanning calorimetry (DSC), thermogravimetric analysis (TGA), X-ray powder diffractometry (XRPD), attenuated total reflectance-Fourier transform infrared spectroscopy (ATR-FTIR) and scanning electron microscopy (SEM). Additionally, the amorphous form of imatinib mesylate was obtained via supercooling of the melt in a DSC apparatus, and aged at various temperatures (3, 15, 25 and 30 °C) and time periods (1–16 h). Glass transition and enthalpy relaxation were used to calculate molecular-relaxation-time parameters. The Kohlrausch–Williams–Watts (KWW) equation was applied to fit the experimental enthalpy-relaxation data. The mean molecular-relaxation-time constant (τ) increased with decreasing ageing temperature. The results showed a high stability of amorphous imatinib mesylate adequate to enable its use in solid dosage form.

Modeling Superionic Behavior of Plutonium Dioxide

The Bredig transition to the superionic phase indicated with the $$\lambda $$-peak in $${C_p}$$ was highly expected for plutonium dioxide ($${\rm{Pu}}{{\rm{O}}_2}$$) as other actinide dioxides. However, least-square fit and local smoothing techniques applied to the experimental enthalpy data of PuO2 in 1980s could not detect a $$\lambda $$-peak in specific heat that might be due to too scattered and insufficient experimental data. Therefore, this issue has not been yet put beyond the doubts. In the current article, a superionic model of $${\rm{Pu}}{{\rm{O}}_2}$$ is developed with partially ionic model of a rigid ion potential. Thermophysical properties were calculated in constant pressure–temperature ensemble using molecular dynamics simulation. The Bredig transition with vicinity of a $$\lambda $$-peak in specific heat was successfully observed for the model system at about 2,100 K. Moreover, the experimental enthalpy change was well reproduced before and after the estimated transition temperature.

Determination of Energy Characteristic and Microporous Volume by Immersion Calorimetry in Carbon Monoliths

Activated carbon monoliths disc and honeycomb type were prepared by chemical activation of coconut shell with zinc chloride at different concentrations, without using a binder. The structures were characterized by N2adsorption at 77 K and immersion calorimetry into benzene. The experimental results showed that the activation with zinc chloride produces a wide microporous development, with micropore volume between 0,38 and 0,79 cm3g-1, apparent BET surface area between 725 and 1523 m2g-1and immersion enthalpy between 73,5 and 164,2 Jg-1. We compared the experimental enthalpy with calculated enthalpy by equation Stoeckli-Kraehenbuehl finding a data dispersion from which can infer that the structures are not purely microporous; this fact is ratified with similar behavior that the evidence t the product EoWo.

Thermodynamics of SmCl3 and TmCl3: Experimental Enthalpy of Fusion and Heat Capacity. Estimation of Thermodynamic Functions up to 1300K

Heat capacities of solid SmCl3 and TmCl3 were measured by differential scanning calorimetry in the temperature range from 300 K up to the respective melting temperatures. The heat capacity of liquid SmCl3 was also investigated. These results were compared with literature data and fitted by a polynomial temperature dependence. The temperature coefficients were given. Additionally, the enthalpy of fusion of SmCl3 was measured. Furthermore, by combination of these results with the literature data on the entropy at 298.15 K, S0m (LnCl3, s, 298.15 K) and the standard molar enthalpy of formation of ΔformH0m (LnCl3, s, 298.15 K), the thermodynamic functions were calculated up to T = 1300 K.