Iodides ionic radii

The above speculation [21] may be extended to include the related quaternary ammonium compounds such as xylocholine (XXXIX). It is probable that the volumes of the guanidinium ion and the trimethylammonium group are similar. The ionic radius of the guanidinium ion (IX) is about 3A the ionic radius of the tetramethylammonium ion has been estimated [300] to be 3-4A, although rather smaller values have also been proposed [301-303]. Crystallographic analyses of muscarine iodide [304], choline chloride [305] and acetylcholine bromide [306] have revealed that the carbon to nitrogen distance is about l-SA, and that a hydrogen bond (C-H-0 distance 2-87-3 07A) exists in the crystals of these compounds. 

Thirdly, there is the purely structural argument from Relative Size if ions of one type are much the largest, they will effectively fix the structure since the others can pack between them. This argument, which makes no assumption whatever about electron-clouds, is often referred only to lithium iodide, but much more evidence is available. Such questions of crystal-form and isomorphism are in fact the most important applications of ionic-radius systems in chemistry and mineralogy (cp. the classical work of V. M. Goldschmidt (2)). 

Lithium iodide crystallizes in the NaCl lattice in spite of the fact that r+/r is less than 0.414. Its density is 3.49 g/cm3. Calculate from these data the ionic radius of the iodide ion. 

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... 

The values of AGfc calculated above are listed in Table IX and plotted against (l/rA+) in Fig. 3. The straight line is drawn with slope calculated from Eq. (75), using Ds = 78.5, D0 = 1.78, for water at 25°C. The points for iodide ion is clearly unacceptable, as already discussed, but it seems unlikely that any revision will bring AGFC(I) up from zero to the predicted value of ca. 70 kcal mole - h Clearly more data are needed before the continuum theory can be subjected to even an approximate quantitative test, but already we may forecast that the dependence of AGFC on ionic radius will be less sensitive than Eq. (75) implies. 

Explain the fact that the ionic radius of cesium ion is less than that of iodide ion. 

Ethanol has also received considerable attention as a solvent over a long period of time. Data on this solvent, however, are rather few compared to methanol and very few systematic studies exist. Several solubility studies have been made since the publication of Seidell and Linke. Thomas has reported solubilities for the alkali metal iodides at 20 and 25°C, and observed a decrease in solubility with an increase in ionic radius of the cation. Deno and Berkheimer have reported the solubilities of several tetraalkylammonium perchlorates. In every case the solid phase was the pure salt. Solubilities for several rare earth compounds have been reported.Since all of these salts form solvates in the solid phase, the results cannot be used in thermodynamic calculations without the corresponding thermodynamic values for the solid phases. Solubilities of silver chloride, caesium chloride, silver benzoate, silver salicylate and caesium nitrate have been measured in ethanol, using radioactive tracer techniques. Burgaud has measured the solubility of LiCl from 10.2 to 57.6°C and observed that there is a transition from the four-solvated solid phase to the non-solvated phase at 20.4°C. 

The polarizability of an ion is directly related to the ionic radius in the hydrated state it is one of the solute-specific properties that determines the affinity of an ion toward the stationary phase. In general, the retention time increases with increasing ionic radius in the hydrated state and thus, with stronger polarizabihty. Accordingly, halide ions elute in the orden fluoride < chloride < bromide < iodide. The retention time difference between bromide and iodide is already so large that the set of halide ions can be analyzed only in a single run by using special eluents or stationary phases. 

On crossing the lanthanide series from La to Lu, there is a 16% decrease in the ionic radius of the La + ion for a fixed coordination number. If coordination numbers were purely determined by the packing of spheres round a central ion, then a decrease in coordination number would be expected on crossing the series. This expectation is realized in the halides thus for the fluorides, the coordination number (CN) decreases from 11 to 9 in both the chlorides and bromides, the CN decreases from 9 to 6 and in the iodides, the CN decreases from 8 to 6. 

An excellent example of counterion influence is the quite different thermal behavior of double-chain l-methyl-3,5-bis(n-hexadecyloxycarbonyl)pyridinium ion in crystals with iodide or chloride as counterion [4]. The iodide salt revealed three phase transitions solid crystalline-solid crystalline at —326 K, solid crystalline-liquid crystalline at —358 K, and liquid crystalline-isotropic liquid at —378 K. The X-ray diffraction pattern of the liquid crystalline phase could be best rationalized in terms of a smectic-H phase. The chloride anion could be unfavorable for liquid crystalline behavior because of its smaller ionic radius relative to the iodide anion. Less shielding of the positive charges of the pyridinium rings by the chloride counterion leads to increased electrostatic repulsion between headgroups. 

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