Ion-solvation equilibrium

The reduction process of polycycles by lithium metal converts the neutral atoms to anions. The electron transfer is best achieved in ethereal solvents. This enables the stabilization of the lithium cation by coordination to the oxygen atoms of the solvent. The hydrocarbon anion and the cation are linked together by electrostatic forces in which the solvent molecules are also involved, therefore the ion-solvation equilibrium should be considered8. The limiting cases in this equilibrium are free ions and contact ion-pairs (CIP), and in between there are several forms of solvent separated ion-pairs (SSIP)9. In reality, anionic species of aromatic hydrocarbons in ethereal solvents exist between CIP and SSIP. Four major factors influence the ion-solvation equilibrium of lithium-reduced 7T-conjugated hydrocarbons, as observed by H and 7Li NMR spectroscopies8,10. 

The following observation emphasizes the influence of the temperature on ion-solvation equilibrium. The reduction product of 1 with lithium metal in methyltetrahydrofuran is temperature-dependent11. At —120 °C only the radical anion (1 ) could be observed by ESR, while at higher temperatures the paramagnetism disappears and the dianion (12 ) is detected. This reaction must be endothermic it therefore seems that disproportionation is driven by entropy and not by energy, due to ion-pair-solvation equilibrium. It is noteworthy that 12 cannot be observed by NMR spectroscopy due to its special electronic structure12. 

The stability of anionic systems is governed by several factors (a) carbon hybridization (b) effective overall n-conjugation (c) inductive effects (d) aromatic stabilization and (e) environmental factors, e.g., ion-solvation equilibrium. 

In the early 50 s, an ion pair model was introduced by Winstein to rationalize the mechanism and stereochemistry of solvolysis of sulfonates72). This research of carbocationic intermediates and the role of ion solvation equilibrium in reaction mechanisms represents a landmark in the study of charged species. These thermodynamically different ionic species were coined as free ions, contact ion-pairs (c.i.p.), and solvent-separated ion pairs (s.s.i.p.). The ion pair situation can be described as an equilibrium between thermodynamically distinct contact (c.i.p.) and solvent-separated ion pairs (s.s.i.p.) 2-l3 16 The situation should be represented by a continuum of ion-solvation equilibria states in which the two extreme states are the c.i.p. and the s.s.i.p. 2 76) (Eq. 12)... 

The detailed understanding of such a system also allows the assessment of the positions to be attacked, most likely by electrophiles. Moreover, the degree of the system s diatropicity may also be dependent on ion-solvation equilibrium. In this case it has already been shown that the system is not diatropic even if its total number of electrons may predict diatropicity 85). Mullen studied independently a family of anions including the dianion of acenaphthylene (82 ) and he also concludes that in THF solutions the lithium and potassium salts exist predominantly as contact ion pairs 83b). 

Cyclic voltametry of 3 showed two one-electron reduction waves (Ej,c = — 1.92 V and EpC = —2.43 V), the first step is reversible 82c). As already discussed, a careful 13C NMR study of 32" reports the charge densities at the various carbon atoms and the ion-solvation equilibrium of its dilithium salt 60,82cl, showing that the lithium salt exists mainly as a contact ion pair (c.i.p.). 

Acenaphtylene (82 ) dianion is obtained by alkali metal reduction (e.g. lithium or sodium) in an ether solvent. Early 3H NMR studies were reported by Lawler and Ristagno 10, n). This dianion was subject to detailed studies which concentrated on its mode of electron delocalization and the ion-pairing equilibrium. The mode of electron delocalization84b) is mainly deduced from H and 13C NMR chemical shifts in tetrahydrofuran (THF) (Fig. 8) while the ion-solvation equilibrium was deduced from 1H, 13C and 7Li shifts in THF-dg, 2-MeTHF, Et20 and THF-HMPA-d8 83a). 

Ion-Exchange Equilibrium. Retention differences among cations with an anion exchanger, or among anions with a cation exchanger, are governed by the physical properties of the solvated ions. The stationary phase will show these preferences ... 

Many organic reactions involve acid concentrations considerably higher than can be accurately measured on the pH scale, which applies to relatively dilute aqueous solutions. It is not difficult to prepare solutions in which the formal proton concentration is 10 M or more, but these formal concentrations are not a suitable measure of the activity of protons in such solutions. For this reason, it has been necessaiy to develop acidity functions to measure the proton-donating strength of concentrated acidic solutions. The activity of the hydrogen ion (solvated proton) can be related to the extent of protonation of a series of bases by the equilibrium expression for the protonation reaction. 

The driving force for diffusion of C+ from the membrane to the aqueous solution is the favorable solvation of the ion by water. As C+ diffuses from the membrane into the water, there is a buildup of positive charge in the water immediately adjacent to the membrane. The charge separation creates an electric potential difference ( ou,cr) across the membrane. The free-energy difference for C+ in the two phases is AG = —nFE(Mcr, where F is the Faraday constant and n is the charge of the ion. At equilibrium, the net change in free energy for diffusion of C+ across the membrane boundary must be 0 ... 

Kimura Y, Alfano JC, Walhout PK, Barbara PF (1994) Ultrafast transient absorption spectroscopy of the solvated electron in water. J Phys Chem 98 3450-3458 Li X, Sevilla MD, Sanche L (2003) Density functional theory studies of electron interaction with DNA can zero eV electrons induce strand breaks J Am Chem Soc 125 13668-13699 Lind J, Shen X, Eriksen TE, Merenyi G, Eberson L (1991) One-electron reduction of N-bromosuccin-imide. Rapid expulsion of a bromine atom. J Am Chem Soc 113 4629-4633 Marasas RA, lyoda T, Miller JR (2003) Benzene radical ion in equilibrium with solvated electrons. J Phys Chem A 107 2033-2038... 

Far from third-phase formation, Kanellakopulos et al. (118) showed in an earlier study that the extraction behavior of given electrolytes with the same cation is primarily influenced by the solvation properties of the associated anions. They found that the electrolyte phase distribution can be explained by single ion solvation, by comparing the equilibrium constants for the extraction of acids by undiluted TBP with the free energies of transfer for the anions (Table 7.3). 

The initial photoinduced radical ion pair equilibrium is further complicated by competing solvation equilibria. As shown in Eq. 2, the electron transfer typically... 

While in the methods treated before ion solvation represents the sum of various terms of ion-solvent interaction, spectroscopic methods are mainly, if at all, sensitive to the immediate environment of an ion. Due to this the coordination model, representing the primary solvation shell, is not only used for highly charged ions but also for univalent ions. The precise results of the direct ion-solvent interactions made it possible to evaluate equilibrium constants describing the composition in the solvation shell of an ion in mixed solvents. Therefore, the estimation of single ion free ener es of transfer from spectroscopic measurements is the subject of several recent efforts and is theme of Part III. 

The solvation and correlation are taken into account, thereby departing from the RPM. The dependence of solvation on distance was recently investigated [67], The expression of the ion-pairing equilibrium constant is analogous to that developed by Gilkerson or by Fuoss. Again, the solvation parameter avoids the linearity of In /Teh with Me. 

Therefore, knowledge of the solvent properties and ion solvation in such media is essential. Although an exact determination of the solvation energy of individual ions is not possible, extrathermodynamic assumptions have been introduced in order to estimate this parameter. These ideas will be presented before the influence of solvents on equilibrium and kinetic parameters of electrode reactions is discussed. As a basis for this discussion, a brief presentation of the properties of solvents frequently used in electrochemical experiments will be given. 

In order to use Eq. (6) in electrochemical studies of ion solvation, the problems related to the liquid junction potential have been presented in Sec. 2.2.2. Equation (11) may also be used in such studies, but the measured potentials should be expressed versus the same solvent-independent reference electrode. Such electrodes, which give a basis for the formation of a uniform scale of electrode potentials in different solvents, are available. A scale of this kind is also needed for a correlation of equilibrium potentials (E° 1/2) of electrode systems in various solvents. 

The second equilibrium involves hydrated 1 ions in equilibrium with E" complexes. In the forward step, an iodide ion donates an electron to the working electrode and is hence oxidized to a I atom. We may speculate that the 1 ion remains outside the double layer and that an electron tunnels through the double layer. However, from many experimental results and molecular dynamic simulations (see Section 4.7.2), it became clear that this is not the case. Instead, a solvated ion penetrates the double layer and becomes chemisorbed as a 1" ion (<5 < 1) on the metal surface, losing about half of its hydration shell [18, 19]. Moreover, there is a local restructuring of the double layer. Here also, the electrochemical reaction does not involve tunneling of an electron through the double layer. 

Very little is known about the effect of this interaction and how important this equilibrium is for the cationic polymerization, especially in solid/liquid interface reactions. Triethyloxonium fluoroborate, an excellent initiator for formaldehyde polymerization, can be visualized as an ethylcarbonium ion solvated by one mole of diethylether. 

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