Solvents, acidic ionizing power

The ionization eonstant should be a function of the intrinsic heterolytic ability (e.g., intrinsic acidity if the solute is an acid HX) and the ionizing power of the solvents, whereas the dissoeiation constant should be primarily determined by the dissociating power of the solvent. Therefore, Ad is expeeted to be under the eontrol of e, the dieleetrie eonstant. As a consequenee, ion pairs are not deteetable in high-e solvents like water, which is why the terms ionization constant and dissociation constant are often used interchangeably. In low-e solvents, however, dissociation constants are very small and ion pairs (and higher aggregates) become important species. For example, in ethylene chloride (e = 10.23), the dissociation constants of substituted phenyltrimethylammonium perchlorate salts are of the order 10 . Overall dissociation constants, expressed as pArx = — log Arx, for some substanees in aeetie acid (e = 6.19) are perchloric acid, 4.87 sulfuric acid, 7.24 sodium acetate, 6.68 sodium perchlorate, 5.48. Aeid-base equilibria in aeetie acid have been earefully studied beeause of the analytical importance of this solvent in titrimetry. 

In this contribution, we describe and illustrate the latest generalizations and developments[1]-[3] of a theory of recent formulation[4]-[6] for the study of chemical reactions in solution. This theory combines the powerful interpretive framework of Valence Bond (VB) theory [7] — so well known to chemists — with a dielectric continuum description of the solvent. The latter includes the quantization of the solvent electronic polarization[5, 6] and also accounts for nonequilibrium solvation effects. Compared to earlier, related efforts[4]-[6], [8]-[10], the theory [l]-[3] includes the boundary conditions on the solute cavity in a fashion related to that of Tomasi[ll] for equilibrium problems, and can be applied to reaction systems which require more than two VB states for their description, namely bimolecular Sjy2 reactions ],[8](b),[12],[13] X + RY XR + Y, acid ionizations[8](a),[14] HA +B —> A + HB+, and Menschutkin reactions[7](b), among other reactions. Compared to the various reaction field theories in use[ll],[15]-[21] (some of which are discussed in the present volume), the theory is distinguished by its quantization of the solvent electronic polarization (which in general leads to deviations from a Self-consistent limiting behavior), the inclusion of nonequilibrium solvation — so important for chemical reactions, and the VB perspective. Further historical perspective and discussion of connections to other work may be found in Ref.[l],... 

Gropp measured the conductivity of liquid and frozen soln. of lithium, sodium, and potassium chlorides. F. Ratig studied the electro-chemical action—vide alkali chlorates. The electrical conductivity of soln. of lithium chloride in several non-aqueous solvents has been investigated. Formic acid as a solvent exerts an ionizing power of the same order of magnitude as water in acetic acid, the lithium chloride seems to be partially associated to double molecules, > and in some solvent,... 

Final cleavage of the peptide from the support requires HF or other acids of high ionizing power (TfOH, HBr/AcOH [18]). HF is a good solvent for amino acids and proteins, and does not usually lead to the cleavage of peptide bonds [23], Because of its toxicity and capacity to dissolve glass, HF must be handled with great care in special HF-resistant containers. 

Solutions of alkali metal and ammonium iodides in liquid iodine are good conductors of electricity, comparable to fused salts and aqueous solutions of strong acids. The liquid is therefore a polar solvent of considerable ionizing power, whereas its own electrical conductivity suggests that it is appreciably ionized, probably into I+ and I"3 (triodide). Iodine resembles water in this respect. The metal iodides and polyiodides are bases, whereas the iodine halides are acids. 

Ionization of liquid ammonia and water solutions.—Solutions of certain salts in liquid ammonia are good conductors of electricity so that liquid ammonia approaches water in its ionizing power. The effect, however, is largely due to the high speed at which the ions are supposed to travel in the solvent. For example, E. C. Franklin and H. P. Cady1 find that univalent ions travel, at —33°, nearly three lames as fast as in aq. soln. at 18°. Just as the solvent water, in the ionization theory of hydrolysis, is supposed to be ionized H20=0H -f-H, so in ammonolysis, the solvent ammonia is supposed to be ionized NH3==NH2-j-H . Sodamide, NaNH, furnishes sodium ions Na and amide ions NH 2 when dissolved in liquid ammonia, and it is to be considered as a base. It reddens phenolphthalein. The neutralization of this solution results in the union of H ions with NH2 ions to form ammonia molecules, just as the neutralization of bases is regarded as an effect of the union of H and OH ions. Acetamide, CH3.CO.NH2, ionizes in liquid ammonia in an analogous manner CH3.CO.NH2 CH3.CO.NH -f-H, and it thus behaves as an acid. 

Consequently there appears to be a sound empirical basis for the use of the OTs scale of solvent ionizing power. Its use should be restricted to sulphonates, however, because of the differential effects of electrophilic solvation in acidic solvents (see Section 4). The importance of these effects can be seen by comparing the Y and Iqxs values for carboxylic acids (Table 5) it appears that, relative to 80% ethanol/water, a carboxylic acid ionizes a tosylate about ten times more rapidly than a chloride. 

The difference in mechanisms of addition of trifluoroacetic acid (see above) and HC1 (in acetio acid) to 3-hexyne has been tentatively explained as due to the different nucleophilic and ionizing power of the two solvents. 

Sections 3.3.1 and 4.2.1 dealt with Bronsted acid/base equilibria in which the solvent itself is involved in the chemical reaction as either an acid or a base. This Section describes some examples of solvent effects on proton-transfer (PT) reactions in which the solvent does not intervene directly as a reaction partner. New interest in the investigation of such acid/base equilibria in non-aqueous solvents has been generated by the pioneering work of Barrow et al. [164]. He studied the acid/base reactions between carboxylic acids and amines in tetra- and trichloromethane. A more recent compilation of Bronsted acid/base equilibrium constants, determined in up to twelve dipolar aprotic solvents, demonstrates the appreciable solvent influence on acid ionization constants [264]. For example, the p.Ka value of benzoic acid varies from 4.2 in water, 11.0 in dimethyl sulfoxide, 12.3 in A,A-dimethylformamide, up to 20.7 in acetonitrile, that is by about 16 powers of ten [264]. 

Another instructive example of electrophilic or H-bonding assistance of protic solvents (or co-solvents) in SnI reactions is the accelerated acetolysis rate of 2-bromo-2-methylpropane upon the addition of phenols to a tetrachloromethane/acetic acid solution of the reactant [582] see reference [582] for further examples. The usefulness of phenol as a solvent for SnI solvolysis reactions, in particular phenolysis of 1-halo-l-phenylethanes, has been stressed by Okamoto [582], In spite of its low relative permittivity (fir = 9.78 at 60 °C), its low dipolarity fi = 4.8 10 Cm = 1.45 D), and its low nucleophihcity, it represents a solvent of high ionizing power due to its electrophilic driving force. 

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