Enthalpy of solvation

The layer of solvent molecules not directly adjacent to the metal is the closest distance of approach of solvated cations. Since the enthalpy of solvation of cations is nomially substantially larger than that of anions, it is nomially expected that tiiere will be insufBcient energy to strip the cations of their iimer solvation sheaths, and a second imaginary plane can be drawn tlirough the centres of the solvated cations. This second plane is temied the outer Helmholtz plane (OHP). 

Before considering different theoretical approaches to determining the free energies and other thermodynamic properties of ionic solvation, it is important to be aware of a problem on the experimental level. There are several methods available for obtaining these quantities for electrolyte solutions, both aqueous and nonaqueous some of these have been described by Conway and Bockris162 and by Padova.163 For example, enthalpies of solvation can be found via thermodynamic cycles, free energies from solubilities or galvanic cell potentials. However the results... 

The standard enthalpy of solvation of any substance A, therefore, is the difference between its standard enthalpy of formation in solution and its standard enthalpy of formation in the gas phase ... 

Figure 5.1 Thermochemical cycle (f = 298.15 K), showing how solution and gas-phase bond dissociation enthalpies are related. AS0 VH° are standard enthalpies of solvation.

Figure 1.4 Thermodynamic cycle comparing the lattice enthalpy, A Hi, and the enthalpy of hydration, AHb, for AgF. The enthalpy of solvation, AHs, is equal to the difference between A Hi and AHb...

Fig. 8. Free enthalpies of solvation AGsoi for the halides of silver and sodium in water...

In Table 6 the differences of free enthalpies of solvation for several anion ligands in a donor solvent D and in AN are given. HMPA shows very weak solvation whereas water is a very strong solvating agent for anions. The free enthalpies of solvation of halide and pseudohalide ions are by 4 to 15 keal/mol more negative than in aprotic donor solvents. 

In a donor solvent the iodide ions is much more strongly solvated than the neutral donor and hence the donor properties of the iodide ion are lowered in solution. This event has been described as the thermodynamic solvatation effect. It becomes increasingly important with an increase of the ratio of the free enthalpy of solvation to the free enthalpy of the ligand exchange reaction. 

For a given metal ion the free enthalpies for (a) and (b)remain constant and hence the value for E° of a given redox system in various solvents is determined by the corresponding free enthalpies of solvation of the cation. 

As a first approximation to entropic terms AS° in the respective solvent may be considered as constant E° for a given ion in different solvents is then determined by the enthalpy of solvation AH°. 

A considerable amount of work has been done on the development of water-ion potential energy functions." " Most of these functions are of the standard Lennard-Jones plus Coulomb form, with parameters selected to give the experimental free energy or enthalpy of solvation. ... 

The qualitative trend predicted by this equation is that, when the heat of solution is negative (the dissolution is exothermic, i.e., heat is evolved, the enthalpy of solvation is more negative than the lattice enthalpy is positive), the solubility diminishes with increasing temperatures. The opposite trend is observed for endothermic dissolution. An analogue of Eq. (2.58), with H replacing G, and the same tables [12] can be used to obtain the required standard enthalpies of solution of ionic solutes. No general analogues to Eqs. (2.53)-(2.55) are known as yet. 

Step 1 in Fig. 2 corresponds to the measured enthalpy of the reaction in a polar solvent, step 4 the enthalpy of forming the adduct in the gas phase, steps 2 and 3 the enthalpies of desolvating the acid and base respectively, and step 5 the enthalpy of solvating the adduct. The enthalpy measured, AHi, is the sum of many contributions, i.e.,... 

Lack of thermodynamic data prevents our making full use of the relationship between standard redox potential and free enthalpy of solvation. For many simple redox systems the standard redox potential has been found to be related to the donicities of the respective ligands. This question has been extensively discussed in a previous paper (3). 

Here Vij denotes the distance between atoms i and j and g(i) the type of the amino acid i. The Leonard-Jones parameters Vij,Rij for potential depths and equilibrium distance) depend on the type of the atom pair and were adjusted to satisfy constraints derived from as a set of 138 proteins of the PDB database [18, 17, 19]. The non-trivial electrostatic interactions in proteins are represented via group-specific dielectric constants ig(i),g(j) depending on the amino-acid to which atom i belongs). The partial charges qi and the dielectric constants were derived in a potential-of-mean-force approach [20]. Interactions with the solvent were first fit in a minimal solvent accessible surface model [21] parameterized by free energies per unit area (7j to reproduce the enthalpies of solvation of the Gly-X-Gly family of peptides [22]. Ai corresponds to the area of atom i that is in contact with a ficticious solvent. Hydrogen bonds are described via dipole-dipole interactions included in the electrostatic terms... 

TABLE 2. Enthalpy of solvation of components of the Schlenk equilibrium in Et20 and THE (kJmol ) ... 

It is noteworthy that enthalpy depends mainly on electronic interactions, while entropy depends mainly on trans-lation and rotation solvation affects both enthalpy and entropy. Enthalpies and entropies of solvation usually tend to oppose each other. For charged species, the more negative (favorable) the enthalpy of solvation, the more negative (unfavorable) the entropy of solvation. 

Scheme 4. PET in electron donor — acceptor (EDA) complexes (D+ A--) . .vertical" ion pair (D+ A- ), adiabatic ion pair AGS free enthalpy of solvation (cf. [42]).

The enthalpies of solvation of four geometric isomers of 2,5-dimethyl-l-phenyl-1-thioxophosphorinan-4-one in chloroform, nitrobenzene, and methanol were calculated using the enthalpies of vaporization of the isomers determined by the modified Solomonov-Konovalov method from the enthalpies of solution of the compounds in CCI4 and />-xylene and molar refractions <2000RCB1522>. 

Since the electrostatic component of AGsoi depends on the permittivity of the solvent, e, and on the cavity size (represented by means of the vector normal to the cavity surface, n), the electrostatic contribution to the enthalpy of solvation, A//ele, can be determined as indicated in Eq. 4-10 ... 

A series of 16 molecules, which include different monofunctional compounds, were chosen to determine the enthalpy of solvation in water. Besides four hydrocarbons (hexane, heptane, octane and cyclohexane) and water, the series of molecules include alcohols (2-methylpropan-2-ol, 1-butanol and 2-butanol), ethers (diethylether, tetrahydrofuran and tetrahydropyran), amines (propylamine, butylamine, diethy-lamine and dibutylamine) and piperidine. This choice allows us to examine the differences between different functional groups, as well as the influence of the molecular size on the enthalpic contributions for a given series of monofunctional compounds. Free energies of hydration as well as the corresponding enthalpies taken from the data compiled by Cabani and coworkers [26] are shown in Table 4-1. 

The enthalpy change for the reaction is favorable because (1) electrostatic repulsion between the negative charges in ATP exceeds that in the reaction products, (2) the reaction products are resonance stabilized, and (3) the enthalpies of solvation of the products are larger than that for ATP. The entropy change for the reaction is favorable because of the release of a phosphate group. Note that this implies that ATP hydrolysis is strongly temperature-dependent [cf. Eq. 10.7)]. 

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