Complex ions approach

The results were presented in the form of isotherms, in which the properties are plotted versus the concentration. Nevertheless analysis of the isotherms was made based on available melting diagrams approach that the melts consist of TaFg3 and TaF7Cl3 complex ions. However, according to this general conception [312-314], the isotherm of the surface tension must, in such a case, have either a minimum or at least display prominence of the dependence in the direction of the concentration axis. 

M NH3(aq). Because of the large value of Kf = 1.6 x 107, we start by having the reagents form as much complex ion as possible, and approach equilibrium from that point. 

The molar conductance values of the complex Ln(DPSO)6 I3 in acetonitrile are slightly higher than those suggested for 1 1 electrolytes, due to the displacement of some coordinated iodide by the solvent (250). The conductance values observed for the complexes, however, approach more closely the values reported for 1 1 electrolytes as the ionic size of the lanthanide ion decreases. This may be due to the increasing strength of the metal-anion bond with decreasing cation size. 

If the constant X has a value much larger than unity, practically all ions of the complexing agent approaching the membrane react with AgX, liberating X in an amount corresponding to one half of their amount. If, making this assumption, the individual terms in the sums in the numerator and the denominator of (3.4.27) are compared, it can be seen that the denominator in this equation will always be a small quantity. Thus we can write... 

From Tables 6.3 and 6.4 it seems that the size and charge correlations can be extended to complex ions. This observation is very important because it indicates a possibility to estimate the ion interaction coefficients for complexes by using such correlations. It is, of course, always preferable to use experimental ion interaction coefficient data. However, the efforts needed to obtain these data for complexes will be so great that it is unlikely that they will be available for more than a few complex species. It is even less likely that one will have data for the Pitzer parameters for these species. Hence, the specific ion interaction approach may have a practical advantage over the inherently more precise Pitzer approach. 

Secondly, as illustrated in Table 1, the eneigy gained by solvation for ions is so great that a gas phase ion will accept essentially any species as a solvent, including the neutral molecule with which it is reacting. When an ion approaches a neutral reactant, an ion-miolecule complex is formed. This is typically about 7 to 35 kcal/mol more stable than the reactants. For any activation barrier to rise above the eneigy of the reactants (i.e., for the reaction to have a positive it must be... 

Anion effects have been observed especially in relation to dissolution of the cation complexes in media of low polarity. Soft organic and inorganic anions (phenates, thiocyanate, permanganate etc) generally allow ready dissolution the pier ate anion has been much used (62, 66). The interaction between the anion and the complexed cation may affect the stability of the complex. Ion pairing may occur when the anion can contact the complexed cation, as in the case of macrocyclic complexes, where approach of the anion from top and bottom is possible. This is observed in the RbNCS complex of IS, but not in its NaNCS complex, as shown by the crystal structure data (100). With bromide as anion both a complexed ion pair and a complexed sodium cation are found in the solid state for (15, NaBr) (118). 

These measurements have been carried out in collaboration with de Maeyek[4]). The rate constant was found to be (1 3 0-2)-10n litres/ mol-sec thus the neutralization reaction is the fastest known bi-molecular reaction in aqueous solution. Molecular-kinetic considerations show that the velocity of recombination is solely determined by the collision frequency of the ions. Furthermore, the effective cross section of the proton is so large that the reaction already proceeds spontaneously when ions approach each other within a distance of two to three H-bonds. This means that the motion of the proton within the hydration complex (the diameter of which corresponds to about two to three H-bonds) proceeds rapidly compared to the actual movement of the ions towards each other. 

Recent studies indicate that the adsorption of metal ions is controlled only in part by the concentration of the free (aquo) metal ion of considerable importance is the ability of hydroxo and other complex ions and molecules to adsorb. There have been two apparently divergent approaches to describe the role played by hydroxo metal complexes in adsorption at solid-aqueous electrolyte interfaces. Matijevic et al. (9) have proposed that specific hydrolysis products—e.g., Al8(OH)2o4+ in the A1(III)-H20 system, are responsible for extensive coagulation and charge reversal of hydrophobic colloids. It has also been demonstrated by Matijevic that the free (aquo) species of transition and other metal ions... 

Obviously one could measure the pH of a known concentration of a weak acid and obtain a value of its hydronium ion activity, which would permit a direct evaluation of its dissociation constant. However, this would be a one-point evaluation and subject to greater errors than by titrating the acid halfway to the equivalence point. The latter approach uses a well-buffered region where the pH measurement represents the average of a large number of data points. Similar arguments can be made for the evaluation of solubility products and stability constants of complex ions. The appropriate expression for the evaluation of solubility products again is based on the half-equivalence point of the titration curve for the particular precipitation reaction [AgI(OH2)2h represents the titrant] ... 

For multistep complexation reactions and for ligands that are themselves weak acids, extremely involved calculations are necessary for the evaluation of the equilibrium expression from the individual species involved in the competing equilibria. These normally have to be solved by a graphical method or by computer techniques.26,27 Discussion of these calculations at this point is beyond the scope of this book. However, those who are interested will find adequate discussions in the many books on coordination chemistry, chelate chemistry, and the study and evaluation of the stability constants of complex ions.20,21,28-30 The general approach is the same as outlined here namely, that a titration curve is performed in which the concentration or activity of the substituent species is monitored by potentiometric measurement. 

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