Ion solvation, enthalpy

Fig. 8. Comparison of total single-ion solvation enthalpies and gas-phase solvation...

Cieplak, P. and P. Kollman (1990). Monte Carlo Simulation of Aqueous Solutions of Li+ and Na+ Using Many-Body Potentials. Coordination Numbers, Ion Solvation Enthalpies, and the Relative Free Energy of Solvation. I. Chem. Phvs. 921111 6761. 

Essentially this expression states that the ratio of heats of solvation in water for any pair of ions is equal to their ratio in any other solvent. They use the Csl assumption to arrive at the single ion solvation enthalpies given in Appendix 2.11.6. According to these results, a maximum in A/fg%v(ion) occurs in all cases at an alcohol mol fraction of around 0.25. 

Single Ion Solvation Enthalpies in Water, — Afls°oiv(i > )/kcal mol , at 25°C (Molal Scale)... 

Ion solvation enthalpies scale inversely with ion radius. Ionic radii increase from Li to Na to K yet the enthalpies of solution decrease (see Figure 22.14). Write an electrostatic expression to rationalize this trend. 

D. S. Gill, J. P. Singla, R. C. Paul, and S. P. Narula, J. Chem. Soc. Dalton Trans., 522 (1972). Thermochemical studies and ion solvation enthalpies in formamide, A-methylformamide and iV, A-dimethylforma-mide. 

It is possible to detemiine the equilibrium constant, K, for the bimolecular reaction involving gas-phase ions and neutral molecules in the ion source of a mass spectrometer [18]. These measurements have generally focused on tln-ee properties, proton affinity (or gas-phase basicity) [19, 20], gas-phase acidity [H] and solvation enthalpies (and free energies) [22, 23] ... 

If a substance is to be dissolved, its ions or molecules must first move apart and then force their way between the solvent molecules which interact with the solute particles. If an ionic crystal is dissolved, electrostatic interaction forces must be overcome between the ions. The higher the dielectric constant of the solvent, the more effective this process is. The solvent-solute interaction is termed ion solvation (ion hydration in aqueous solutions). The importance of this phenomenon follows from comparison of the energy changes accompanying solvation of ions and uncharged molecules for monovalent ions, the enthalpy of hydration is about 400 kJ mol-1, and equals about 12 kJ mol-1 for simple non-polar species such as argon or methane. 

In 1906, Matignon reported an enthalpy of solution of -21.54 kcal mol-1 (-90 kj mol ) for neodymium trichloride in ethanol (178). His ethanol may have been less than perfectly anhydrous, and the value for pure ethanol somewhat less negative than this, perhaps —80 or — 70 kJ mol 1. Certainly a value in this region is considerably less negative than his value for the enthalpy of solution of neodymium trichloride in water, —148 kj mol-1. The difference may reasonably be attributed to less favorable solvation qf the constituent ions in ethanol than in water. Ion solvation would be expected to be even less favorable in isopropanol, so it is not surprising to find an enthalpy of solution of about + 40 kJ mol-1 for neodymium trichloride in this alcohol. This estimate must be considered as only approximate, as it is derived from... 

The enthalpies of solution and solubilities reviewed here provide much of the experimental information required in the derivation of single-ion hydration and solvation enthalpies, Gibbs free energies, and entropies for scandium, yttrium, and lanthanide 3+ cations. 

In the course of our polarographic studies on organic cations we determined the half-wave potentials, 1/2, for various arylmethylium ions [1-11]. The aim of the present work is to extract from these values some new information concerning the relative magnitude of their solvation enthalpies in three very different solvents. A comparison of our results [obtained in methanesulphonic acid (MSA) and dichloromethane (DCM)] with those of Volz and Lotsch [12] [obtained in cyanomethane (CM) solutions] yields some useful conclusions. 

Here R+ is the carbenium ion from equation (1), AHS6 are solvation enthalpies, (R ) is the ionization potential of radical R, and Z includes all electro-energetic terms which do not depend upon the nature of R. If, to a first approximation, TAS0 and AHS(R ) are taken to be independent of the nature of R and are incorporated into Z, then ... 

The direct access to the electrical-energetic properties of an ion-in-solution which polarography and related electro-analytical techniques seem to offer, has invited many attempts to interpret the results in terms of fundamental energetic quantities, such as ionization potentials and solvation enthalpies. An early and seminal analysis by Case etal., [16] was followed up by an extension of the theory to various aromatic cations by Kothe et al. [17]. They attempted the absolute calculation of the solvation enthalpies of cations, molecules, and anions of the triphenylmethyl series, and our Equations (4) and (6) are derived by implicit arguments closely related to theirs, but we have preferred not to follow their attempts at absolute calculations. Such calculations are inevitably beset by a lack of data (in this instance especially the ionization energies of the radicals) and by the need for approximations of various kinds. For example, Kothe et al., attempted to calculate the electrical contribution to the solvation enthalpy by Born s equation, applicable to an isolated spherical ion, uninhibited by the fact that they then combined it with half-wave potentials obtained for planar ions at high ionic strength. 

M. D. Tissandier, K.A. Cowen, W.Y. Feng, E. Gundlach, M. H. Cohen,A. D. Earhart, J. V Coe, T. R. Tuttle Jr. The Proton s Absolute Aqueous Enthalpy and Gibbs Free Energy of Solvation from Cluster-Ion Solvation Data. J. Phys. Chem. A 1998, 102, 7787-7794. 

Fig. 1. Born-Haber cycle for the formation of solvated ions from an ionic crystal [M+X ]w. U lattice energy, Affsoiv. enthalpy of ion solvation...

In the absence of data for the solvation enthalpies in nonaqueous media, an attempt has been made to plot the half-wave potentials for a given ion in different solvents (expressed in the bis(biphenyl)chromium(I) scale) vs the donicity of the solvent molecules. The Fig. 22 to 26 reveal a relationship between Et /2 and DN, i.e., the half-wave potential becomes more negative with increasing donicity of the solvent. 

Table 9. Comparison of experimental gas-phase solvation enthalpies for the ions H+, F, and Cl- by different solvents.

To round off the discussion of solvation enthalpies, reference is made here to the articles by Case 38> and by Friedman and Krishnan 27>, who have reviewed the thermodynamic aspects of ion solvation extensively. 

Above 0.1m the values of aqueous salts flatten out and may become negative except where small ions are involved. In 90% acetonitrile, the 0l curve becomes steeper with increasing concentration. This is probably caused by depletion of the free water in the solvent, so that ion solvation by acetonitrile becomes significant. It is also possible that, because ionic electric fields are significant at greater distances in a solvent of lower dielectric constant, the overlapping of ionic co-spheres has a much more drastic effect on the enthalpy than it does in water for this salt. 

The calorimetric measurements in metal oxide-aqueous electrolyte solution systems are, beside temperature dependence of the pzc measurements, the method for the determination of the enthalpy of the reaction in this system. Because of the low temperature effects in such systems they demand very high precision. That is why these measurements may be found only in a few papers from the last ten years [89-98]. A predominant number of published measurements were made in the special constricted calorimeters (bath type), stirring the suspension. The flow calorimeters may be used only for sufficiently large particles of the solid. A separate problem is the calculation of the enthalpy of the respective reactions from the total heat recorded in the calorimeter. A total thermal effect consists of the heat of the neutralization in the liquid phase, heat connected with wetting of the solid, heat of the surface reaction and heat effects caused by the ion solvation changes (the ions that adsorb in the edl). Considering the soluble oxides, one should include the effects connected with the transportation of the ions from the solid to the solution... 

Of course, it is impossible to determine a a priori. If Eq. (17) is physically reasonable, it should be possible to choose a in such a way that a plot of AG (g) — AG°(g) versus DN can be represented by a fairly smooth curve, as shown in Fig. 4 for a s 0.5. This value appears reasonable low a values are highly improbable since the free Cl ion is undoubtedly a much stronger base than the coordinated Cl ion in [CoCl ] on the other hand, solvation enthalpies of the complex anions will compensate only in part, so that a is necessarily < 1. Equation (17) can be used to estimate free energies of formation or stability constants of [CoCl4] in other solvents, provided that the donicities and the values AG(sv)(Cl ) (Table V) are known. Values A( (g) — A( °(g) required for this purpose may be interpolated or extrapolated from Fig. 4. 

---