Solvation enthalpies of neutral solutes in water and octanol (original) (raw)
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Thermodynamics of the solvation of non-electrolytes in C8 monofunctional organic solvents
Thermochimica Acta, 2001
Some results are reported concerning the enthalpies of solvation (D solv H 0 ) of some hydrocarbons (hexane, cyclohexane), ethers (tetrahydrofuran), ketones (cyclopentanone, pentan-3-one), alcohols (butan-1-ol, butan-2-ol), and water in octan-2-ol and hexylacetate. The D solv H 0 values were calculated by combining the calorimetric heats of solution here determined with the known enthalpies of vaporisation.
Journal of Molecular Liquids, 2009
Data have been assembled from the published literature on the enthalpies of solvation for 91 organic vapors and gaseous solutes in 2-propanol, for 73 gaseous compounds in 2-butanol, for 85 gaseous compounds in 2-methyl-1-propanol and for 128 gaseous compounds in ethanol. It is shown that an Abraham solvation equation with five descriptors can be used to correlate the experimental solvation enthalpies to within standard deviations of 2.24 kJ/mole, 1.99 kJ/mole, 1.73 kJ/mole and 2.54 kJ/mole for 2-propanol, 2-butanol, 2-methyl-1-propanol and ethanol, respectively. The derived correlations provide very accurate mathematical descriptions of the measured enthalpy of solvation data at 298 K, which in the case of ethanol span a range of 136 kJ/mole. Division of the experimental values into a training set and a test set shows that there is no bias in predictions, and that the predictive capability of the correlations is better than 3.5 kJ/mole.
Comparison Study of Polar and Nonpolar Contributions to Solvation Free Energy
Journal of Chemical Information and Modeling, 2017
In this study, we compared the contributions of polar and non-polar interactions to the solvation free energy of a solute in solvent, which is decomposed into four different terms based on the nature of interactions: (i) electrostatic solvation free energy term counting for the work done to move solute charges from fixed points in some reference environment to their configuration positions in solvent; (ii) solute-solvent van der Waals dispersion interactions; (iii) change on solvent-solvent interactions and solvent entropy due to reorganization of solvent around solute cavity in solvent; and (iv) compensation of electrostatic forces acting on the dielectric surface boundary between solvent and solute. We compared these contributions to each other for a data set of 573 proteins, which were prepared using CHARMM22 and AMBER force fields. In addition, we compared the calculated with experimental hydration free energies for a data set of 642 small molecules, which were prepared using general AMBER force field. Our results indicated the significance of each term to the total solvation free energy.
Thermochimica Acta, 2007
Data have been assembled on the enthalpies of solvation of 373 compounds in water and 138 compounds in 1-octanol. It is shown that an Abraham solvation equation with five descriptors can be used to correlate the experimental solvation enthalpies to within standard deviations of 3.68 kJ/mol (water) and 2.66 kJ/mol (1-octanol). The derived correlations provide very accurate mathematical descriptions of the observed enthalpies of solvation, which in the case of water span a range of 150 kJ/mol. Division of the experimental values into a training set and a test set shows that there is no bias in predictions and that the predictive capability of the correlations is better than 4 kJ/mol.
Thermochimica Acta, 2011
Data have been assembled from the published literature on the enthalpies of solvation for 91 organic vapors and gaseous solutes in 2-propanol, for 73 gaseous compounds in 2-butanol, for 85 gaseous compounds in 2-methyl-1-propanol and for 128 gaseous compounds in ethanol. It is shown that an Abraham solvation equation with five descriptors can be used to correlate the experimental solvation enthalpies to within standard deviations of 2.24 kJ/mole, 1.99 kJ/mole, 1.73 kJ/mole and 2.54 kJ/mole for 2-propanol, 2-butanol, 2-methyl-1-propanol and ethanol, respectively. The derived correlations provide very accurate mathematical descriptions of the measured enthalpy of solvation data at 298 K, which in the case of ethanol span a range of 136 kJ/mole. Division of the experimental values into a training set and a test set shows that there is no bias in predictions, and that the predictive capability of the correlations is better than 3.5 kJ/mole.
Solvation enthalpies of the proton in polar and non-polar solvents: Theoretical study
Acta Chimica Slovaca, 2013
In spite of the importance of proton transfer in solution-phase processes, there is still no systematic theoretical study of proton solvation enthalpies. We have investigated the solvation enthalpies of the proton in seven solvents of various polarities (benzene, chloroform, acetone, methanol, ethanol, DMSO, water) using the Integral Equation Formalism Polarized Continuum Model (IEF-PCM). All computations were performed at the B3LYP and BHLYP levels of theory with aug-cc-pVDZ, aug-cc-pVTZ and aug-cc-pVQZ basis sets. Our calculations have shown that the B3LYP and BHLYP functionals provide similar solvation enthalpies. Finally, differences in the solvation enthalpy of the proton values stemming from the various basis sets do not exceed 6 kJ mol-1, with exception of DMSO and chloroform. Distance between H+ and the acceptor atom of the solvent molecule is the shortest in the case of water. It has been also found that the B3LYP distances are slightly longer.
Accurate determination of solvation free energies of neutral organic compounds from first principles
Nature Communications
The main goal of molecular simulation is to accurately predict experimental observables of molecular systems. Another long-standing goal is to devise models for arbitrary neutral organic molecules with little or no reliance on experimental data. While separately these goals have been met to various degrees, for an arbitrary system of molecules they have not been achieved simultaneously. For biophysical ensembles that exist at room temperature and pressure, and where the entropic contributions are on par with interaction strengths, it is the free energies that are both most important and most difficult to predict. We compute the free energies of solvation for a diverse set of neutral organic compounds using a polarizable force field fitted entirely to ab initio calculations. The mean absolute errors (MAE) of hydration, cyclohexane solvation, and corresponding partition coefficients are 0.2 kcal/mol, 0.3 kcal/mol and 0.22 log units, i.e. within chemical accuracy. The model (ARROW FF) ...
Computation of hydration free energies of organic solutes with an implicit water model
Journal of Computational Chemistry, 2006
A new approach for computing hydration free energies ⌬G solv of organic solutes is formulated and parameterized. The method combines a conventional PCM (polarizable continuum model) computation for the electrostatic component ⌬G el of ⌬G solv and a specially detailed algorithm for treating the complementary nonelectrostatic contributions (⌬G nel ). The novel features include the following: (a) two different cavities are used for treating ⌬G el and ⌬G nel . For the latter case the cavity is larger and based on thermal atomic radii (i.e., slightly reduced van der Waals radii). (b) The cavitation component of ⌬G nel is taken to be proportional to the volume of the large cavity. (c) In the treatment of van der Waals interactions, all solute atoms are counted explicitly. The corresponding interaction energies are computed as integrals over the surface of the larger cavity; they are based on Lennard Jones (LJ) type potentials for individual solute atoms. The weighting coefficients of these LJ terms are considered as fitting parameters. Testing this method on a collection of 278 uncharged organic solutes gave satisfactory results. The average error (RMSD) between calculated and experimental free energy values varies between 0.15 and 0.5 kcal/mol for different classes of solutes. The larger deviations found for the case of oxygen compounds are probably due to a poor approximation of H-bonding in terms of LJ potentials. For the seven compounds with poorest fit to experiment, the error exceeds 1.5 kcal/mol; these outlier points were not included in the parameterization procedure. Several possible origins of these errors are discussed. Figure 3. (a) Calculated [⌬G nel (calc)] and experimental [⌬G nel (exp)] nonelectrostatic components of solvation free energies in kcal/mol for aromatic and nitrogen compounds. (b) Calculated [⌬G nel (calc)] and experimental [⌬G nel (exp)] nonelectrostatic components of solvation free energies in kcal/mol for aromatic and nitrogen compounds. ⌬G el calculated with DMol 3 . compounds: ethers and alcohols. (b) Calculated [⌬G nel (calc)] and experimental [⌬G nel (exp)] nonelectrostatic components of solvation free energies in kcal/mol for oxygen compounds: esters and carbonyl groups. (c) Calculated [⌬G nel (calc)] and experimental [⌬G nel (exp)] nonelectrostatic components of solvation free energies in kcal/mol for oxygen compounds: ethers and alcohols. ⌬G el calculated with DMol 3 .
The Journal of Chemical Physics
We demonstrate that with two small modifications, the popular dielectric continuum model is capable of predicting, with high accuracy, ion solvation thermodynamics (Gibbs free energies, entropies, and heat capacities) in numerous polar solvents. We are also able to predict ion solvation free energies in water-co-solvent mixtures over available concentration series. The first modification to the classical dielectric Poisson model is a perturbation of the macroscopic dielectric-flux interface condition at the solute-solvent interface: we add a nonlinear function of the local electric field, giving what we have called a solvation-layer interface condition (SLIC). The second modification is including the microscopic interface potential (static potential) in our model. We show that the resulting model exhibits high accuracy without the need for fitting solute atom radii in a state-dependent fashion. Compared to experimental results in nine water-co-solvent mixtures, SLIC predicts transfer free energies to within 2.5 kJ/mol. The co-solvents include both protic and aprotic species, as well as biologically relevant denaturants such as urea and dimethylformamide. Furthermore, our results indicate that the interface potential is essential to reproduce entropies and heat capacities. These and previous tests of the SLIC model indicate that it is a promising dielectric continuum model for accurate predictions in a wide range of conditions.