Hydration entropy and enthalpy

The roles of hydration enthalpies and entropies in determining the solubilities of ionic compounds... 

The study of hydrated biomolecules has only just begun and very few papers have been published on that topic [103-107]. Hydration enthalpies and entropies for the peptides LHRH and bradykinin [108] and for the protein BPTI [106,107] have been measured and the amount of water addition to folded vs unfolded states of cytochrome c has been studied [103-105]. At this point it is too early to draw general conclusions but it is already apparent that only a few water molecules can have significant structural consequences [107,108]. 

Finally, let us point out that the absolute standard electrode potential value of the couple H+w/H2(g) is actually about 4.5 V. This value cannot be verified since we cannot measure an absolute potential. It was obtained by using thermodynamic cycles, taking into account some thermodynamic data such as the proton hydration enthalpy and entropy. These last ones have been approached by considering the quadrupole model of water (see Chap. 1). It is quite evident that the value of 4.5 V differs considerably from the conventional one (0.00 V). However, it does not change the redox phenomena provision since only the standard electrode potential differences are taken into account. 

The heats of solution of americium metal in aqueous hydrochloric add solutions at 298.15 0.05 K were redetermined in 1972 by Fuger, Spirlet, and Muller [241] with pure ameridum metal prepared by distillation. Combined with earlier results, the standard enthalpy of formation of Am (aq) at 198 K of — 616.7 1.3kJmol was calculated [242]. For hydration enthalpy and entropy, see Choppin [359]. 

Solvent effects also play an important role in the theory separating enthalpy and entropy into external and internal parts (134-136) or, in other terms, into reaction and hydration contributions (79). This treatment has been widely used (71, 73, 78, 137-141). The most general thermodynamic treatment of intermolecular interaction was given by Rudakov (6) for various states of matter and for solution enthalpy and entropy as well as for kinetics. A particular case is hydrophobic interaction (6, 89, 90). 

After this computer experiment, a great number of papers followed. Some of them attempted to simulate with the ab-initio data the properties of the ion in solution at room temperature [76,77], others [78] attempted to determine, via Monte Carlo simulations, the free energy, enthalpy and entropy for the reaction (24). The discrepancy between experimental and simulated data was rationalized in terms of the inadequacy of a two-body potential to represent correctly the n-body system. In addition, the radial distribution function for the Li+(H20)6 cluster showed [78] only one maximum, pointing out that the six water molecules are in the first hydration shell of the ion. The Monte Carlo simulation [77] for the system Li+(H20)2oo predicted five water molecules in the first hydration shell. A subsequent MD simulation [79] of a system composed of one Li+ ion and 343 water molecules at T=298 K, with periodic boundary conditions, yielded... 

The acidity dependence of the activation enthalpies and entropies, All and AS. of the hydration of 1,3-cyclohexa- and 1,3-cyclooctadienes was ascribed30 to a dielectric solvation effect in dilute acids, which is overcome by increasing solvent structure as the availability of water decreased in concentrated acids. This suggestion was one of the early premises of a more recent interpretation31 of acidity effects in terms of water activity and solvation of cationic species. 

Often, it is difficult to distinguish definitely between inner sphere and outer sphere complexes in the same system. Based on the preceding discussion of the thermodynamic parameters, AH and AS values can be used, with cation, to obtain insight into the outer vs. inner sphere nature of metal complexes. For inner sphere complexation, the hydration sphere is disrupted more extensively and the net entropy and enthalpy changes are usually positive. In outer sphere complexes, the dehydration sphere is less disrupted. The net enthalpy and entropy changes are negative due to the complexation with its decrease in randomness without a compensatory disruption of the hydration spheres. 

Both U02(TTA)2 and Th(TTA)4 have two molecules of hydrate water when extracted in benzene, and these are released when TBP is added in reactions Eqs. (4.11) and (4.12). The release of water means that two reactant molecules (e.g., U02(TTA)2 2H2O and TBP) formed three product molecules (e.g., U02(TTA)2 TBP and 2H2O). Therefore, AS is positive. Since TBP is more basic than H2O, it forms stronger adduct bonds, and, as a consequence, the enthalpy is exothermic. Hence, both the enthalpy and entropy changes favor the reaction, resulting in large values of log K. 

From measurements of the temperature dependency of the equilibrium constant, thermodynamic parameters may be deduced (section 3.4). Very few enthalpy and entropy constants have been derived for the distribution reaction MAj(aq) MA2(org) of neutral complexes such investigations give information about hydration and organic phase solvation. 

This chapter consists of a description of the structure of liquid water and the nature of ions in aqueous solution. The discussion is largely restricted to the interactions between monatomic ions with liquid water in which they become hydrated by acquiring a hydratiun sphere or shell-Additionally, a few diatomic and polyatomic anions are dealt with, including the important hydroxide ion. The hydration of ions derived from the s- and p-block elements of the Periodic Table, and the derivation of values of their enthalpies and entropies of liydralioii, are described in considerable detail. 

The models may be somewhat deficient, but estimates of the enthalpies and entropies of hydrated ions are accurate and are the subject of Section 2.4. 

Absolute values lor the enthalpies and entropies of hydration of ions were discussed in terms of their sizes and chai ses. 

The nature of ions in solution is described in some detail and enthalpies and entropies of hydration of many ions are defined and recalculated from the best data available. These values are used to provide an understanding of the periodicities of standard reduction potentials. Standard reduction potential data for all of the elements, group-bygroup, covering the s-and p-, d- and/- blocks of the Periodic Table is also included. Major sections are devoted to the acid/base behaviour and the solubilities of inorganic compounds in water. 

This book offers no solutions to such severe problems. It consists of a review of the inorganic chemistry of the elements in all their oxidation states in an aqueous environment. Chapters 1 and 2 deal with the properties of liquid water and the hydration of ions. Acids and bases, hydrolysis and solubility are the main topics of Chapter 3. Chapters 4 and 5 deal with aspects of ionic form and stability in aqueous conditions. Chapters 6 (s- and p-block). 7 (d-block) and 8 (f-block) represent a survey of the aqueous chemistry of the elements of the Periodic Table. The chapters from 4 to 8 could form a separate course in the study of the periodicity of the chemistry of the elements in aqueous solution, chapters 4 and 5 giving the necessary thermodynamic background. A more extensive course, or possibly a second course, would include the very detailed treatment of enthalpies and entropies of hydration of ions, acids and bases, hydrolysis and solubility. 

The enthalpy of solution has been discussed somewhat more quantitatively by Morris.30 He has pointed out the relation between the enthalpy of solution and the difference between the hydration enthalpy of the cation and that of the anion. This difference win be largest when the cation and anion differ most in size (Fig. 8.9). ]n these cases the enthalpy of solution tends to be large and negative and favors solution. When the hydration enthalpies (and the sizes) are more nearly alike, the crystal is favored. When entropy effects are added, a very nice correlation with the solubility and the free energy solution is found (Fig. 8.10). 

Hickel B, Sehested K (1985) Activation energy for the reaction H + OH" —> eaq. Kinetic determination of the enthalpy and entropy of solvation of the hydrated electron. J Phys Chem 89 5271-5274 Hoffman MZ, Hayon E (1973) Pulse radiolysis study of sulfhydryl compounds in aqueous solution. J Phys Chem 77 990-996... 

To summarise E° values for redox couples of the type M(s)/M"+(aq) can largely be rationalised in terms of the atomisation enthalpies, ionisation energies and hydration enthalpies. The entropy terms can be neglected in most cases. 

It can be anticipated that the computation of A//soi and AAsoi is more delicate than the prediction of AGsoi, which benefits from the enthalpy-entropy compensation. Accordingly, the suitability of the QM-SCRF models to predict the enthalpic and entropic components of the free energy of solvation is a challenging issue, which could serve to refine current solvation continuum models. This contribution reports the results obtained in the framework of the MST solvation model [15] to estimate the enthalpy (and entropy) of hydration for a set of neutral compounds. To this end, we will first describe the formalism used to determine the MST solvation free energy and its enthalpic component. Then, solvation free energies and enthalpies for a series of typical neutral solutes will be presented and analyzed in light of the available experimental data. Finally, collected data will be used to discuss the differential trends of the solvation in water. 

Table 4-1. Experimental free energy, enthalpy and entropy of hydration (kcal/mol) for the series of neutral molecules considered in this study...

Table 4-3. Enthalpic and entropic components of the electrostatic and non-electrostatic terms of the hydration free energy (kcal/mol). The non-electrostatic enthalpy and entropy were determined by subtracting the electrostatic enthalpy and entropy from the corresponding experimental data...

Table 1. Standard enthalpy and entropy of hydration for the protonated species indicated...

From a molecular view, the decrease of entropy upon hydrophobic hydration is not mitigated by a large hydration enthalpy and this translates into an increase in the free energy of water. A system will tend to minimize this increase in free energy through association of the hydrophobic moieties. This phenomenon that explains the salting in of a neutral hydrophobic molecule by hydrophobic ions is expected to amplify with the sizes of the hydrophobic moieties [49]. Attractive forces between two hydrophobic ions and repulsive forces between hydrophilic and hydrophobic... 

The data on decomposition temperatures of hydroxides and hydrated oxides, as well as the data on enthalpy and entropy of their formation, are listed in Table 2.7 according to [30-32]. 

TABLE 19.9 The Enthalpies and Entropies of Hydration for the Halide Ions... 

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