Alkali metal ions solvation number

Ion solvation has been studied extensively by potentiometry [28, 31]. Among the potentiometric indicator electrodes used as sensors for ion solvation are metal and metal amalgam electrodes for the relevant metal ions, pH glass electrodes and pH-ISFETs for H+ (see Fig. 6.8), univalent cation-sensitive glass electrodes for alkali metal ions, a CuS solid-membrane electrode for Cu2+, an LaF3-based fluoride electrode for l , and some other ISEs. So far, method (2) has been employed most often. The advantage of potentiometry is that the number and the variety of target ions increase by the use of ISEs. 

Alkali-metal-ion-O-donor-solvent cages exhibit low-frequency bands in the i.r. spectrum characteristic of the cation-O polyhedra. Similar bands are seen with N-donors. The bands can be used to establish the nature of cation co-ordination and serve as probes to examine ion-solvent interactions.87 An analysis of the i.r. spectra of ternary mixtures LiClCX-Sj-Sz showed a preferential solvation of the Li+ ion by NH3 and methylamines (St) in MeCN or THF (S2) and by MeCN (S2) in MeN02 (SO. The appearance of wide bands in the ion-cage vibration region is related to the formation of different species [Li(Si) -i(S2)i]+. The Li+-solvent molecule interaction energy decreases when the number of S2 molecules in the first solvation shell increases. The mean composition of the first solvation shell was obtained from intensity measurements of the molecules not bonded to the ions in favourable cases (S2 = ND3 or MeCN), the solvation number of the Li+ ion in the pure solvent... 

Studies of the nuclear resonances of Cl, Br, and have also been carried out in various alkali metal halide solutions [17]. The magnitude of the chemical shift increases with electrolyte concentration and also with atomic number of the anion. In the case of the alkali metal ions the chemical shift becomes more positive in the series Na+ < K+ < Li" " < Rb+ < Cs+. The results were attributed to direct interaction between the cation and anion in solutions containing K, Rb, and Cs. With the smaller cations, interactions between the halide ion and the water molecules solvating the cation are more important. 

Self-organization as a result of the formation of amide-bridged dimers is not limited to rfiorganoamides of the alkali metals. Occasionally alkah metal derivatives of primary amines have been shown to be dimeric in the solid state. Structurally characterized examples are [LiNHPh(thf)2]2, 79a [77a], [LiNHC6F5(thf)2]2, 79b, and [K NHC6H2(CF3)3-2,4,6 (thf)3]2, 105 [77b]. These three examples nicely illustrate how the composition of the solvates depends on the size of the alkali metal ion. In 79a and 79b two additional THF ligands are sufficient to saturate the coordination sphere of lithium (coordination number 4), whereas the potassium ions in 105 are coordinated by three THF ligands and by two weak K- - F interactions. 

The half-wave potential of Cs+ is almost independent of the nature of the solvent as well as that postulated for Rb+. Na+ and K+ still have low solvation energies, but they vary with the nature of the solvent, showing a decrease with decreasing donor number. A stronger dependence is found for Li+, which is known to have stronger solvating tendencies than the other alkali metal ions. 

Mass spectrometry has been used to study the energetics of solvation and has shown that the enthalpies of attachment of successive water molecules to either alkali metal or halide ions become less exothermic as the number of water molecules increases (Kebarle, 1977). The Gibbs free energies of attachment for water molecules have also been found to be negative. 

The alkali metals share many common features, yet differences in size, atomic number, ionization potential, and solvation energy leads to each element maintaining individual chemical characteristics. Among K, Na, and Li compounds, potassium compounds are more ionic and more nucleophilic. Potassium ions form loose or solvent-separated ion pairs with counteranions in polar solvents. Large potassium cations tend to stabilize delocalized (soft) anions in transition states. In contrast, lithium compounds are more covalent, more soluble in nonpolar solvents, usually existing as aggregates (tetramers and hexamers) in the form of tight ion pairs. Small lithium cations stabilize localized (hard) counteranions (see Lithium and lithium compounds). Sodium chemistry is intermediate between that of potassium and lithium (see Sodium and sodium alloys). 

The number of solvents that have been used in SrnI reactions is somewhat limited in scope, but this causes no practical difficulties. Characteristics that are required of a solvent for use in SrnI reactions are that it should dissolve both the organic substrate and the ionic alkali metal salt (M+Nu ), not have hydrogen atoms that can be readily abstracted by aryl radicals (c/. equation 13), not have protons which can be ionized by the bases (e.g. Nth- or Bu O" ions), or the basic nucleophiles (Nu ) and radical ions (RX -or RNu- ) involved in the reaction, and not undergo electron transfer reactions with the various intermediates in the reaction. In addition to these characteristics, the solvent should not absorb significantly in the wavelength range normally used in photostimulated processes (300-400 nm), should not react with solvated electrons and/or alkali metals in reactions stimulated by these species, and should not undergo reduction at the potentials employed in electrochemically promoted reactions, but should be sufficiently polar to facilitate electron transfer processes. 

The formation of molecular radical ions by electron transfer reactions between alkali metals and a wide variety of aromatic and other organic compounds in polar solvents is well established. A very large number of radical anions have been prepared by this method and extensive studies of their e.s.r. and optical spectra have been made (Bowers, 1965 Gerson, 1967 Kaiser and Kevan, 1968). In solution the electron transfer reaction will be facilitated by the subsequent solvation of the two ions (or ion pair) by the polar solvent molecules. However, we have observed that similar electron transfer reactions occur readily when alkali metal atoms are deposited on a variety of relatively non polar substances at 77°K in the rotating cryostat. In most cases the parent compound acts as the matrix, though for some radical ions an inert matrix of a non-polar hydrocarbon has been used successfully. It is perhaps surprising that the reactions occur so readily as the energy of solvation of the ions must be quite small in most of these systems as compared with that in the polar liquids. 

According to the Lorentz-Lorenz equation (4.3.21) for the molar refraction at optical frequencies, Y is directly proportional to the molecular polarizability p. The Koppel-Palm equation has also been applied to the analysis of solvent effects on thermodynamic quantities related to the solvation of electrolytes [48, 49]. In the case of the systems considered in table 4.11, addition of the parameter X to the linear equation describing the solvent effect improves the quality of the fit to the experimental data, especially in the case of alkali metal halide electrolytes involving larger ions. The parameter Y is not important for these systems but does assist in the interpretation of other thermodynamic quantities which are solvent dependent [48, 49]. Addition of these parameters to the analysis is only possible when the solvent-dependent phenomenon has been studied in a large number of solvents. 

Many X-ray diffraction studies of electrolyte solutions have been carried out in aqueous solutions [Gl, 4, 5]. Values of the most probable distance, between the oxygen atom in water and a number of monoatomic ions are summarized in table 5.1. In the case of the cations, this distance reflects the radius of the cation plus the effective radius of the water molecule measured in the direction of the lone pairs on oxygen. In the case of alkali metals, the effective radius of water increases from 122 pm for Li" " to 131 pm for Cs when the Shannon and Prewitt radii are assumed for the cations (see section 3.2), the average value being 127 pm. This result can be attributed to the observation that the coordination number for water molecules around an alkali metal or alkaline metal earth cation changes with cation size and electrolyte concentration. In the case of the Li" " ion, this number decreases from six in very dilute solutions to four in concentrated solutions [5]. Because of the electrostatic character of the interaction between the cation and water molecules, these molecules exchange rapidly with other water molecules in their vicinity. For this reason, the solvation coordination number should be considered as an average. 

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