Cation desolvation

Fig. 8.11 The Uthium cation desolvation free energy profile for EC/LiPF and EC DMC(3 7)/ LiPFj 1 M electrolytes next to the graphite at 298 K. Z=0 corresponds to position of the hydrogen atoms of the graphite. A snapshot of the simnlation box used for calculating the free energy barrier for charge transfer is shown in the right...

Troxler and Wipff determined conformational preferences of free ligands and solvation patterns of the host-guest complexes for cryptand 222 and their metal cation complexes in acetonitrile. The calculations were carried out with the AMBER force field and also with Aqvist s ion parameters the two sets of results were then compared. When AMBER was used, the relative order of cryptate stabilities and the recognition ability of the cryptand 222 to select K from among Li, Na, Rb, and Cs cations were in qualitative agreement with experimental data. The results obtained with Aqvist s parameters were comparable to or better than those obtained with the AMBER force field. In general, complexes of alkali cations and cryptand 222 are more stable in nonaqueous solvents such as acetonitrile than in water, in part because of the reduced energy cost for cation desolvation upon complexation.i -i i... 

Fig. 7.24 The free-energy profile for the lithium cation desolvation from EC DMC(3 7)A iPFg at 1 M electrolyte at 298 K calculated using the Li probability profile from equilibrium MD simulations and the integration of the constrained force method. Z = 0 at the position of hydrogen atoms of graphite...

Fig. 7.25 The lithium cation desolvation free-energy profile for EC/LiPFe and EC DMC(3 7)/LiPF6 1 M electrolytes next to graphite at 298 K...

Ultrasonic studies have been reported on solutions of copper(ii) nitrate and perchlorate in ethylene glycol. In each case a single, concentration-independent relaxation effect is observed which is probably due to cation desolvation coupled with the diffusive approach of the two solvated ions. Tanaka has proposed a modification to the Bennetto-Caldin scheme for solvent exchange at bivalent metal ions. 

Rates and equilibria within these cation-anion recombination reactions are not correlated. Ritchie considers that extensive desolvation of the reactant ions... 

Even greater disruption is encountered in the case of trivalent cations (Figures 4.9,4.10). They completely penetrate both hydration regions and destroy the structure of water around the polyion. This amounts to complete desolvation. The same is true of bound hydrogen ions which are localized. 

The ratio of the two forms depends on the cation as well as on a. Ba has a greater tendency to make linkages of the COO-Me-OOC type than Mg and this difference is accentuated when the density of COO" in the polyanion is low. Thus, at a = 025 more Ba ions are in the COO-Ba-OOC form than in the COO-Ba form, while the reverse is true for Mg ions. Moreover, the structure COO-Mg is more stable and soluble than COO-Ba because Mg is more hydrophilic than Ba. For these reasons, Ba is precipitated at a = 0-25 while Mg is not. This interpretation is supported by titration experiments in the presence of divalent cations (Jacobsen, 1962). Magnesium forms very stable hydrates and would be expected to be more difficult to desolvate. 

It is well known that lyophilic sols are coagulated by the removal of a stabilizing hydration region. In this case, conversion of a sol to a gel occurs when bound cations destroy the hydration regions about the polyanion, and solvated ion-pairs are converted into contact ion-pairs. Desolvation depends on the degree of ionization, a, of the polyacid, and the nature of the cation. Ba ions form contact ion-pairs and precipitate PAA when a is low (0-25), whereas the strongly hydrated Mg + ion disrupts the hydration region only when a > 0-60. 

As the poly(alkenoic acid) ionizes, polymer chains unwind as the negative charge on them increases, and the viscosity of the cement paste increases. The concentration of cations increases until they condense on the polyadd chain. Desolvation occurs and insoluble salts precipitate, first as a sol which then converts to a gel. This represents the initial set. 

Whenever the concentration of a species at the interface is greater than can be accounted for by electrostatic interactions, we speak of specific adsorption. It is usually caused by chemical interactions between the adsorbate and the electrode, and is then denoted as chemisorption. In some cases adsorption is caused by weaker interactions such as van der Waals forces we then speak of physisorption. Of course, the solvent is always present at the interface so the interaction of a species with the electrode has to be greater than that of the solvent if it is to be adsorbed on the electrode surface. Adsorption involves a partial desolvation. Cations tend to have a firmer solvation sheath than anions, and are therefore less likely to be adsorbed. 

In previous sections, numerous examples of anion activation by cationic micelles and polysoaps were presented. The extent of rate augmentation— 102—lO -fold—cannot be rationalized in terms of concentration effects alone. We believe that these observations are explained most reasonably by the concept of the hydrophobic ion pair (Kunitake et al., 1976a). According to this concept, anionic reagents are activated probably due to desolvation when they form ion pairs with an ammonium moiety in a hydrophobic microenvironment. The activation of anionic species in the cationic micellar phase... 

As mentioned repeatedly, a variety of anionic reagents are highly activated in the hydrophobic microenvironment of cationic micelles and polysoaps. The range of anionic reagents studied in the past includes imidazole, hydroxide, thiolates, oximates, hydroxamates, carboxylates and carbanions. Polyanionic coenzymes are similarly activated. These results can be interpreted in a unified way by the concept of hydrophobic ion pairs, and the major source of activation seems to be concentration and desolvation of the anionic reagent in the... 

The precursor model of FAB applies well to ionic analytes and samples that are easily converted to ionic species within the liquid matrix, e.g., by protonation or deprotonation or due to cationization. Those preformed ions would simply have to be desorbed into the gas phase (Fig. 9.6). The promoting effect of decreasing pH (added acid) on [M+H] ion yield of porphyrins and other analytes supports the precursor ion model. [55,56] The relative intensities of [Mh-H] ions in FAB spectra of aliphatic amine mixtures also do not depend on the partial pressure of the amines in the gas phase, but are sensitive on the acidity of the matrix. [57] Furthermore, incomplete desolvation of preformed ions nicely explains the observation of matrix (Ma) adducts such as [M+Ma+H] ions. The precursor model bears some similarities to ion evaporation in field desorption (Chap. 8.5.1). 

Figure 14, Schematic drawings of conformational changes upon cation binding by glymes, crown ethers, and cryptands (D denotes donor atom) also shown are the slopes (a) and intercepts TASo oi AH-TAS plots as measures of conformational change and desolvation.

Figure 15. Schematic drawing of complexation-induced desolvation from cation and ligand (solvation to the latter is not shown) the complex formation itself reduces the entropy of the system to some extent but the accompanying desolvation leads to much larger increase in entropy ascribed to the liberation of solvent molecules.

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