Cations medium/solvent effects

In chapter 2, Profs. Contreras, Perez and Aizman present the density functional (DF) theory in the framework of the reaction field (RF) approach to solvent effects. In spite of the fact that the electrostatic potentials for cations and anions display quite a different functional dependence with the radial variable, they show that it is possible in both cases to build up an unified procedure consistent with the Bom model of ion solvation. The proposed procedure avoids the introduction of arbitrary ionic radii in the calculation of insertion energy. Especially interesting is the introduction of local indices in the solvation energy expression, the effect of the polarizable medium is directly expressed in terms of the natural reactivity indices of DF theory. The paper provides the theoretical basis for the treatment of chemical reactivity in solution. 

Macroscopic solvent effects can be described by the dielectric constant of a medium, whereas the effects of polarization, induced dipoles, and specific solvation are examples of microscopic solvent effects. Carbenium ions are very strong electrophiles that interact reversibly with several components of the reaction mixture in addition to undergoing initiation, propagation, transfer, and termination. These interactions may be relatively weak as in dispersive interactions, which last less than it takes for a bond vibration (<10 14 sec), and are thus considered to involve "sticky collisions. Stronger interactions lead to long-lived intermediates and/or complex formation, often with a change of hybridization. For example, onium ions are formed with -donors. Even stable trityl ions react very rapidly with amines to form ammonium ions [41], and with water, alcohol, ethers, and esters to form oxonium ions. Onium ion formation is reversible, with the equilibrium constant depending on the nucleophile, cation, solvent, and temperature (cf., Section IV.C.3). 

Alkali halides are compounds with a strong ionic character and without a solvent they are stabilized by the strong electrostatic interaction between the cation and the anion. In the present study we limit attention to the dissociation potentials comparing the curves obtained for the free molecule and for the molecule in water solution. The solvent is treated only as a continuum medium and then in this way we cannot consider the formation of complexes between the ions and the water molecules which instead are extensively studied by means of Monte Carlo (MC) and molecular dynamics (MD) simulations (see for example [13]). Although it could be possible to include some water molecules with the sodium chloride as a more complicated solute, we have preferred to focus attention on the solvent effect on the electronic structure of the simplest solute, this effect being the most important in the next two examples. 

An explanation of the cation size and solvent effect can be found in the energy required to separate the two charges E = (qiq2)/(Dr ). This energy is related to the distance (r) between the ions and the dielectric constant of the medium (D). The separation of the charge centers (r) in the tetrabutylammonium phenoxide ion pair is larger than in the potassium phenoxide ion pair. Tetrabutylammonium phenoxide is therefore more energetic and the nucleophilic anion is more reactive. As a consequence, alkoxide anions exhibit enhanced reactivity under phase transfer conditions. 

Although we first think of acidity, and basicity in terms of aqueous solvents, it must be clear that many of the anions and cations we have discussed could not even exist in an aqueous medium. Solvents do have a significant effect on p/Q. In solution, f-BuO /f-BuOH is a stronger base than MeO /MeOH. ffowever, in the gas phase, methoxide is the stronger base. Methoxide is small and highly solvated, especially in protic solvents, whereas f-BuO is large and poorly solvated (naked, and hence more keen to reacquire a proton). 

If this electrostatic treatment of the substituent effect of poles is sound, the effect of a pole upon the Gibbs function of activation at a particular position should be inversely proportional to the effective dielectric constant, and the longer the methylene chain the more closely should the effective dielectric constant approach the dielectric constant of the medium. Surprisingly, competitive nitrations of phenpropyl trimethyl ammonium perchlorate and benzene in acetic anhydride and tri-fluoroacetic acid showed the relative rate not to decrease markedly with the dielectric constant of the solvent. It was suggested that the expected decrease in reactivity of the cation was obscured by the faster nitration of ion pairs. 

Oae found that for both base- and acid-catalyzed hydrolysis of phenyl benzenesul-fonate, there was no incorporation of 0 from solvent into the sulfonate ester after partial hydrolysis. This was interpreted as ruling out a stepwise mechanism, but in fact it could be stepwise with slow pseudorotation. In fact this nonexchange can be explained by Westheimer s rules for pseudorotation, assuming the same rules apply to pentacoordinate sulfur. For the acid-catalyzed reaction, the likely intermediate would be 8 for which pseudorotation would be disfavored because it would put a carbon at an apical position. Further protonation to the cationic intermediate is unlikely even in lOM HCl (the medium for Oae s experiments) because of the high acidity of this species a Branch and Calvin calculation (See Appendix), supplemented by allowance for the effect of the phenyl groups (taken as the difference in between sulfuric acid and benzenesulfonic acid ), leads to a pA, of -7 for the first pisTa of this cation about -2 for the second p/sTa. and about 3 for the third Thus, protonation by aqueous HCl to give the neutral intermediate is likely but further protonation to give cation 9 would be very unlikely. 

Water is the most common solvent used to dissolve ionic compounds. Principally, the reasons for dissolution of ionic crystals in water are two. Not stated in any order of sequence of importance, the first one maybe mentioned as the weakening of the electrostatic forces of attraction in an ionic crystal known, and the effect may be alternatively be expressed as the consequence of the presence of highly polar water molecules. The high dielectric constant of water implies that the attractive forces between the cations and anions in an ionic salt come down by a factor of 80 when water happens to be the leaching medium. The second responsible factor is the tendency of the ionic crystals to hydrate. 

Attention should be paid to the possible existence of several complexes having different stoichiometries. A necessary preliminary experiment consists of recording the fluorescence and/or excitation spectra under experimental conditions (nature of the solvent, composition of the medium, ionic strength, pH (if it has an effect on the stability constant), etc.) as close as possible to the medium in which a cation must be detected. The variations in the fluorescence intensity for an appropriate couple of excitation and emission wavelengths (or for several emission or excitation wavelengths) as a function of cation concentration must be analyzed in order to determine the stoichiometry and the stability constant of the complexes (Appendix B). As in the case of pH determination (see Section 10.2.1), ratiometric measurements are recommended. 

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