Cation solvation sphere

Presumably, the function of M+ is to "cushion the repulsion of the two negative ions. The larger, softer Cs+ can do this more effectively than the smaller, harder ions such as Li+ or Na+. Also, to form these bridged transition states, solvent molecules must be displaced from the solvation sphere of the cations. That process, because of their smaller sizes, would require more energy for the more strongly solvated Li+ and Na+. For the Cs+ ion, which forms effective bridges, the rate of electron exchange has been found to be linearly related to Cs+ concentration. 

Adsorption of Ag on the surface of PdO is also an interesting option offered by colloidal oxide synthesis. Silver is a well-known promoter for the improvement of catalytic properties, primarily selectivity, in various reactions such as hydrogenation of polyunsaturated compounds." The more stable oxidation state of silver is -F1 Aquo soluble precursors are silver nitrate (halide precursors are aU insoluble), and some organics such as acetate or oxalate with limited solubility may also be used." Ag" " is a d ° ion and can easily form linear AgL2 type complexes according to crystal field theory. Nevertheless, even for a concentrated solution of AgNOs, Ag+ does not form aquo complexes." Although a solvation sphere surrounds the cation, no metal-water chemical bonds have been observed. 

The importance of solvent in determining the stability of crown ether complexes can be seen in Table 2 in which log Ks values for Na+ and K+ complexes of crown ethers in different solvents are compared. The greater values of log Ks in methanol compared to water suggest that the stability of crown ether complexes is mainly enthalpic in origin. Cations are more strongly solvated in water than methanol and the solvation sphere must be removed before complexation can take place. 

The number of donor atoms can be influential in complex formation. Ideally, as the incoming ligand is comparable to an inner solvation sphere, the number of donor atoms should match the preferred coordination number of the cation. An example of this factor comes from a comparison of two similarly sized ligands, [2.2.2]cryptand and [2.2.C8]cryptand, the latter having one dioxoether chain replaced by a C8 unit.478 A reversal of the Ba2+/K+ selectivity of the order of 106 is found, the ratio being 104 for [2.2.2]cryptand and <10 2 for [2.2.C8]cryptand the preferred coordination numbers are six for potassium and eight for barium. 

The two center structures show the complex that forms between iV,iV -bis(2-phenylethyl)-4,13-diaza-18-crown-6 and KI <2000PNAS6271>. The K+ ion is bound in the center of the macroring, as expected for any 18-crown-6 macrocycle. The twin sidearms of the bibracchial lariat ether turn inward in this complex and the arenes serve as apical donors. The top center structure shows the symmetrically bound cation and illustrates that the iodide anion is excluded from the cation s solvation sphere. The bottom center structure shows the superimposition of the two benzene rings upon each other and upon the K+ ion. Note in the bottom center structure that the iodide anion is not illustrated. The ideal sandwich of arene-cation-arene confirms the cation-pi interaction between benzene and K+. 

With the exception of lithium, which has a comparatively strong tendency to form covalent bonds, the chemistry of the alkali metals is essentially ionic. In spite of this common property there are many dissimilarities which result from differences in the ionic radii and the concomitant different solvation abilities. While in aqueous solution at least Li+, and to a lesser extent also Na, has a kinetically comparatively inert solvation sphere, i.e. exchange between solvated and bulk solvent is slow, this is not at all the case for the larger cations K, Rb, and Cs where, owing to the- larger radius of the ion. 

A complex in general is any species formed by specific association of molecules or ions by donor-acceptor interactions (see Topic C9). In aqueous solution the most important complexes are those formed between a metal cation and ligands, which may be ions (e.g. halides, cyanide, oxalate) or neutral molecules (e.g. ammonia, pyridine). The ligand acts as a donor and replaces one or more water molecules from the primary solvation sphere, and thus a complex is distinct from an ion pair, which forms through purely electrostatic interactions in solvents of low polarity (see Topic El). Although complex formation is especially characteristic of transition metal ions it is by no means confined to them. 

Lanthanide halides, nitrates and triflates are not only common reagents in organic synthesis (Fig. 1) but also represent, in dehydrated form, key precursor compounds for the more reactive organometallics (Scheme 2). As a rule, in compounds of strong monobasic acids or even superacids, cation solvation competes with anion complexation, which is revealed by fully or partially separated anions and solvated cations in their solid state structures. The tendency to form outer sphere complexation in coordinating solvents [47] is used as a criterion of the reactivity of inorganic salt precursors in organometallic transformations. 

To associate with a central ion, ligands must compete with the water molecules in (lie central ion s solvation sphere and must lose some of the water molecules in their own solvation sphere. In addition, since many ligands are the anions of weak acids, H+ competes with the central cation for the ligand. Forming a complex ion or ion pair involves competition between the cation and H+ for the ligands, and between the water, OH-, and ligands for the central cation. 

Azamacrobicyclic ethers such as 5 form stable inclusion complexes (cryptates) in which the metal cation is fully sequestered inside a 10 A spherical cavity taking the place of the solvation sphere. The resulting solvent separated ion pairs are particularly reactive in low-polarity media due to their scarce stabilization by the solvent and the cation. 

Solubility data have been used to determine solvation enthalpies, U [defined as in equation (6)] for 0, H , and D in Li and for H in Na and K. The values of 17, are collected in Table 1. Those for H and D in Li are lower than those for 0 and N by factors of ca. 2 and 3, respectively, corresponding to increasing (7, with increasing charge of solute. Those for H" in Li, Na, and K are very similar, that in Li being the greatest. Solvation enthalpies have been derived in at initio M.O. calculations of solvation clusters in Li and Na. By comparison with experimental data, the best model was deduced to be that of a tetrahedral solvation sphere of cations supplemented by a further metal tetrahedron positioned on the three-fold axes of the first solvation sphere. Other incidental results to emerge from the calculations are the effective radii for Li (0.1675 nm), Na (0.1715 nm), and H (0.0525 nm in Li and... 

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