Salts ionic radii

The aluminium ion, charge -I- 3. ionic radius 0.045 nm, found in aluminium trifluoride, undergoes a similar reaction when a soluble aluminium salt is placed in water at room temperature. Initially the aluminium ion is surrounded by six water molecules and the complex ion has the predicted octahedral symmetry (see Table 2.5 ) ... 

Separation Processes. The product of ore digestion contains the rare earths in the same ratio as that in which they were originally present in the ore, with few exceptions, because of the similarity in chemical properties. The various processes for separating individual rare earth from naturally occurring rare-earth mixtures essentially utilize small differences in acidity resulting from the decrease in ionic radius from lanthanum to lutetium. The acidity differences influence the solubiUties of salts, the hydrolysis of cations, and the formation of complex species so as to allow separation by fractional crystallization, fractional precipitation, ion exchange, and solvent extraction. In addition, the existence of tetravalent and divalent species for cerium and europium, respectively, is useful because the chemical behavior of these ions is markedly different from that of the trivalent species. 

Figure 19-4 contrasts the effective sizes of the halide ions. Each of these dimensions is obtained from the examination of crystal structures of many salts involving the particular halide ion. The effective size found for a given halide ion is called its ionic radius. These radii are larger than the covalent radii but close to the van der Waals radii of neutral atoms. 

The agreement is satisfactory, except in the cases where there are deviations from additivity. This fact is a verification of our treatment and of the correctness of our screening constants, for the arbitrary selection of only one ionic radius in a series of salts showing additivity in inter-atomic distances is permitted, and our screening constants fixed four radii independently. 

A similar reasoning may explain the difference in reactivities of the lithium and sodium ion-pairs in THF. The larger ionic radius of the sodium than that of the lithium cation, favoring the formation of loose pairs, makes the sodium pair much more reactive than the lithium salt at lower temperatures. However, at higher temperatures the sodium salt becomes less reactive than the lithium salt as it looses its solvation more readily than the latter. 

Most lanthanide compounds are sparingly soluble. Among those that are analytically important are the hydroxides, oxides, fluorides, oxalates, phosphates, complex cyanides, 8-hydroxyquinolates, and cup-ferrates. The solubility of the lanthanide hydroxides, their solubility products, and the pH at which they precipitate, are given in Table 2. As the atomic number increases (and ionic radius decreases), the lanthanide hydroxides become progressively less soluble and precipitate from more acidic solutions. The most common water-soluble salts are the lanthanide chlorides, nitrates, acetates, and sulfates. The solubilities of some of the chlorides and sulfates are also given in Table 2. 

Because of the small ionic radius of lithium ion, most simple salts of lithium fail to meet the minimum solubility requirement in low dielectric media. Examples are halides, LiX (where X = Cl and F), or the oxides Li20. Although solubility in nonaqueous solvents would increase if the anion is replaced by a so-called soft Lewis base such as Br , I , S , or carboxylates (R—C02 ), the improvement is usually realized at the expense of the anodic stability of the salt because these anions are readily oxidized on the charged surfaces of cathode materials at <4.0 V vs Li. 

The ionophoric properties of 13 toward soft metal cations were evaluated by using the picrate extraction method. Metal ions like Ag" (91%), Tl" (38%) and Hg " (16%) were extracted efficiently with a maximum for Ag". Cu ", Co ", Zn ", Cd ", and ions were not extracted. The better ex-tractability of 13 toward silver(I) cation was attributed to the latter s high affinity for sulfur. A fast exchange process, on the NMR time scale, was observed by NMR between free and complexed 13. In presence of a twofold excess of Ag(I) salt, a 2 4 complex 132 (AgPic)4 was formed (see below) [70]. With this type of complex, the ionic radius of the metal ion is not concerned and the softness of the metal ion should be considered relatively to the soft... 

Bromine has an ionic radius of 1.96 A and thus easily substitutes for chlorine (1.81 A) in the halite crystal lattice as well as in the other chloride salts. The distribution coefficients for bromine in chloride salts deposited from seawater is less than 1 (Warren 2006). 

The salts KCrF3 and RbCrF3 have distorted perovskite structures, and from analysis of powder data the former contains tetragonally compressed CrFe octahedra in which two Cr—F bond distances are 2.00 A and four 2.14 A.244 The ionic radius of Cr23- in fluoride perovskites has been estimated at 0.73 A as a weighted average from lattice constants.245... 

J6 The ammonium km is about the same size (r+ = 151 pm) as the potassium ion ir. 152 pm) and this is a usef ul fact to remember when explaining the resemblance in properties between these two tuns. For example, (he solubilities of ammonium salts arc similar to those of potassium sails. Explain the relation between ionic radius and soloWiiy. On the other hand, all of the potassium halides crystallize in the NaClstrocture with C.N. = 6 (see Chapter 4). but none of the ammonium halides does so. The coordination numbers of the ammonium halides are either four or eight- Suggest an explanation. 

The theoretical picture is expressed by the Debye and McAulay equation [50], which relates the salt effects on nonelectrolytes to the dielectric constant of the solution, e and ionic radius, rj of individual ions. In this theory, the ions make the solution a poorer solvent for a nonelectrolyte, so that its activity coefficient, f , is increased as follows ... 

For a large cation and large anion salt, an increase in either ionic radius would have little effect on the solubility. 

We saw in Fig. 6-30 the conversion of ethylene oxide to crown ethers upon reaction with appropriate metal salts, and demonstrated that the hole sizes of the products corresponded to the ionic radius of the template ion. However, lest we become over-confident, it should be pointed out that the major product from the reaction of ethylene oxide with caesium salts (r = 1.67 A) is not the expected 21-crown-7 with a hole size of about 1.7 A) but 18-crown-6 (hole size, 1.4 A) (Fig. 6-34). The reason for this lies in the structure of the complex formed. We have always assumed that the metal ion will try to lie in the middle of the bonding cavity of the macrocycle. There is no real reason why this should be. Caesium could form a complex with 21-crown-7 in which all of the oxygen atoms lie approximately planar with the metal in the centre of the cavity. It is also apparent that caesium could not occupy the middle of the cavity in 18-crown-6. However, a different type of complex can be formed with 18-crown-6, in which a caesium ion is sandwiched bet-... 

Each rubidium halide (Group VIIA element) crystallizing in the NaCl-type lattice has a unit cell length 30 pm greater than that for the corresponding potassium salt of the same halogen. What is the ionic radius of Rb+ computed from these data ... 

Alkali-earth metals (calcium, barium, and magnesium) complex with polysaccharides extensively (Reisenhofer et al., 1984). Calcium has a smaller atomic and ionic radius than does sodium and, because it has two valence electrons, it is endowed with greater polarizing and bonding ability than Na+. Ca and Ca2+ easily form insoluble complexes with oxygenated compounds. Polysaccharide salts of alkali-earth metals are generally insoluble. 

Partial molar entropies of ions can, for example, be calculated assuming S (H+) = 0. Alternatively, because K+ and Cl ions are isoelectronic and have similar radii, the ionic properties of these ions in solution can be equated, e.g. analysis of B-viscosity coefficients (Gurney, 1953). In other cases, a particular theoretical treatment which relates solvation parameters to ionic radii indicates how the subdivision could be made. For example, the Bom equation requires that AGf (ion) be proportional to the reciprocal of the ionic radius (Friedman and Krishnan, 1973b). However, this approach involves new problems associated with the definition of ionic radius (Stem and Amis, 1959). In another approach to this problem, the properties of a series of salts in solution are plotted in such a way that the value for a common ion is obtained as the intercept. For example, when the partial molar volumes of some alkylammonium iodides, V (R4N+I ) in water (Millero, 1971) are plotted against the relative molecular mass of the cation, M+, the intercept at M + = 0 is equated to Ve (I-) (Conway et al., 1966). This procedure has been used to... 

All the M2+ ions are smaller and considerably less polarizable than the isoelec-tronic M+ ions. Thus deviations from complete ionicity in their salts due to polarization of the cations are even less important. However, for Mg + and, to an exceptional degree for Be2+, polarization of anions by the cations does produce a degree of covalence for compounds of Mg and makes covalence characteristic for Be. Accordingly, only an estimated ionic radius can be given for Be2+ the charge/radius ratio... 

The dissociation constants of trityl and benzhydryl salts are KD 10 4 mol/L in CH2C12 at 20° C, which corresponds to 50% dissociation at 2-10-4 mol/L total concentration of carbocationic species (cf. Table 7) [34]. The dissociation constants are several orders of magnitude higher than those in analogous anionic systems, which are typically KD 10-7 mol/L [12]. As discussed in Section IV.C.l, this may be ascribed to the large size of counterions in cationic systems (e.g., ionic radius of SbCL- = 3.0 A) compared with those in anionic systems (e.g., ionic radius of Li+ 0.68 A), and to the stronger solvation of cations versus anions. However, the dissociation constants estimated by the common ion effect in cationic polymerizations of styrene with perchlorate and triflate anions are similar to those in anionic systems (Kd 10-7 mol/L) [16,17]. This may be because styryl cations are secondary rather than tertiary ions. For example, the dissociation constants of secondary ammonium ions are 100 times smaller than those of quaternary ammonium ions [39]. 

Another effect, called the cesium effect [843, 844], is also connected with the observation that salts of anions with large counterions are highly dissociated in dipolar aprotic solvents, and consequently more anion-reactive. For example, the higher solubility of cesium carbonate in dipolar aprotic solvents and the fact that this salt is far more dissociated than the corresponding Li+, Na" ", or K" " salts, makes this carbonate a superior base in organic synthesis. Amongst the alkali metal cations, the ionic radius of Cs+ r = 334 pm) is more than twice that of Li+ (r = 152 pm). 

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