Dissociation ionogens

It was found in later work that it is precisely the idea of ionic hydration that is able to explain the physical nature of electrolytic dissociation. The energy of interaction between the solvent molecules and the ions that are formed is high enough to break up the lattices of ionophors or the chemical bonds in ionogens (for more details, see Section 7.2). The significance of ionic hydration for the dissociation of electrolytes had first been pointed out by Ivan A. Kablukov in 1891. 

Although cellulose acetate is not inherently a polyelectrolyte there are reports which indicate that it contains a low concentration of weak acid, presumably carboxylic, groups (1). Water absorbed by cellulose acetate membranes might be preferentially located, to some extent, in the region of these ionogenic groups and so assist in their dissociation. 

Non-ionogenic They are those whose molecule cannot undergo dissociation e.g, when alcohol having a high molar mass reacts with several molecules of ethylene oxide, a non-ionogenic surfactant is produced ... 

There are some cases where a reaction, that is, the formation or dissolution of a chemical bond, is involved along with ion exchange phenomena (Helfferich, 1983). Examples of this are acid-base neutralization, dissociation of weak electrolytes in solution or weak ionogenic groups in ion exchangers, complex formation, or combinations of these (Table 5.2). With some of these, very low apparent D in ion exchangers have been noted. 

The charge on colloidal particles may arise from several sources (i) dissociation of ionogenic groups, e.g., proteins or... 

Figure 2.11. Percent of ionogenic (ionizable) species present for weak acids and bases when solution pH is 2 units above or below the acid dissociation constant.

The fact that many compounds are dissociated into ions in aqueous solution is incorrectly put forward as sufficient proof that these substances must also be constructed from ions in the solid state. It is just as incorrect, however, to conclude from non-solubility that there is a bonding other than ionogenic. 

The benzhydryl chlorides and BC13 react with formation of ion pairs (ionization constant, Ki) which dissociate to give the free ions (dissociation constant, KD). Because paired and free diarylcarbenium ions show only slightly different UV-visible spectra, [41], spectrophotometric measurements allow the determination of the total carbocation concentration. On the other hand, only free ions are detected by conductometric analysis, and a combination of both methods allows the determination of Ki and Kd using the theory of binary ionogenic equilibria [42,43]. 

Solutions of non-electrolytes contain neutral molecules or atoms and are nonconductors. Solutions of electrolytes are good conductors due to the presence of anions and cations. The study of electrolytic solutions has shown that electrolytes may be divided into two classes ionophores and ionogens [134]. lonophores (like alkali halides) are ionic in the crystalline state and they exist only as ions in the fused state as well as in dilute solutions. Ionogens (like hydrogen halides) are substances with molecular crystal lattices which form ions in solution only if a suitable reaction occurs with the solvent. Therefore, according to Eq. (2-13), a clear distinction must be made between the ionization step, which produces ion pairs by heterolysis of a covalent bond in ionogens, and the dissociation process, which produces free ions from associated ions [137, 397, 398]. 

The ionization of an ionogen and its subsequent dissociation according to Eq. (2-13) can be further elaborated. Between the ion pair immediately formed on heterolysis of the covalent bond and the independently solvated free ions, there are several steps of progressive loosening of the ion pair by penetration of solvent molecules between the ions. At least four varieties of ion interactions representing different stages of dissociation have been postulated [96, 134, 138, 141] cf. Eq. (2-19) and Fig. 2-14. 

The result of the proton transfer is that two ions have been produced (1) an acetate ion and (2) a hydrated proton. Thus, potential electrolytes (organic acids and most bases) dissociate into ions by ionogenic, or ion-forming, chemical reactions with solvent molecules, in contrast to true electrolytes, which often give rise to ionic solutions by physical interactions between ions present in the ionic crystal and solvent molecules (Fig. 3.1). 

Because the [R2AI] cation can be stabilized by neutral Lewis bases, we assume that this is also true of ionogenic Lewis base alkali metal alkyls R-M. Thus, 1 2 complexes of alkali metal halides with aluminum trialkyls, especially fluorides, may dissociate (3a) into M and [R3AI-F-AIR3] , as shown by the moleeular structure in the crystalline state. Yet there is another possibility, shown in Scheme 3b, i. e., dissociation into aluminum-containing cations and anions. 

In order to illustrate the hybrid ion theory, we shall compare ammonium acetate with an amino acid of which the acid and basic dissociation constants are respectively the same as those of acetic acid and ammonium hydroxide. We know that in 0.1 molar solution, ammonium acetate is 0.5% hydrolyzed and the salt is 99.5% ionized. If we employ the same equation (40 in Chapter One) to calculate the degree of hydrolysis of the amino acid, we find that it too is 0.5% hydrolyzed. Thus it appears logical to assume that the remainder of the ampholyte is present in the ionogenic form. 

The picture is different for weak electrolytes which include most organic acids. A. Hantzsch has demonstrated spectroscopically that an organic acid in water exists in two forms. One has the same structure as the ester of the acid (undissociated form), while the other has the salt structure (ionogenic form). The latter probably dissociates completely into ions, as do strong electrolytes, and it is really unnecessary even to speak of an ionogenic form. 

The simpler equation (4) embodies the Ostwald conception with the difference that Km is not the true dissociation constant but the apparerd constant of the indicator since it represents the product of the true dissociation constant and the equilibrium constant for the normal and aci-forms. The latter equilibrium favors the normal compound in the case of p-nitrophenol so that this substance appears to be a very weak acid. With o-nitro-phenol, however, the existence of the aci-form is favored so that this compound behaves as a stronger acid. The ratio of aci to normal is so large in the case of picric acid that relatively much of the aci- or ionogen form, as compared with the pseudo-compound, is present in aqueous solution. Consequently this substance is a rather strong acid. As the apparent dissociation constant increases, the intensity of the yellow color of aqueous solutions must likewise grow because more of the aci-form will be found in solution. This statement can be confirmed easily. Picric acid in water solutions is yellow, but colorless in organic solvents due to the predominance of the pseudo-form. 

The assumption that all undissociated acetic acid in aqueous solutions is present only in a single form is erroneous because the equilibrium between the ionogen- and pseudo-forms is disregarded. Hence the true dissociation constant of acetic acid is much larger than that usually employed. 

The definition furthermore explains why the solid salt of phenolphthalein is red and that of p-nitrophenol is yellow. The salts have the constitution and, therefore, also the color of the ionogenic form, whether it be completely or only partially dissociated. This is in accord with the findings of Hantzsch (l.c.) and Hantzsch and Robertson (l.c.) mentioned early in this chapter, namely, that Beer s law holds for colored salts as long as the concentration is not too great. This is to be expected if the undissociated salt molecules are ionogenic and possess the same structure and color as do the ions. 

It is important to note that ionophores are not always completely dissociated. For example, when NaCl is dissolved in a solvent of lower relative permittivity, such as methanol, it is ion paired to some extent. The thermodynamics of systems with ion pairing is considered separately in section 3.10. Under these circumstances the ionophore behaves in the same way as a weak electrolyte. On the other hand, all ionogenes are not weak electrolytes. For example, HCl, which is a molecule in the gas phase, is completely dissociated in water and therefore is a strong electrolyte. Acetic acid is completely dissociated in liquid ammonia, which is a much stronger base than water. Thus, the solvent plays an important role in determining the extent of electrolyte dissociation in solution. In the following discussion the traditional terms, strong and weak electrolytes, are used. 

If one or more of the ions in an ion-exchange column is the conjugate acid or base of a weak base or weak acid, the pH of the solution will have an important effect on the degree of dissociation of the weak acid or base. Since it is ions that are retained on ion-exchange columns and since the pH of the solution can affect the relative number of ions in the column because of the dependence of the dissociation constant (or, in other words, pK values) of the ionogenic group on pH, the pH can have a large effect on the retention of the species on the column. 

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