Acid-base equilibria in acetonitrile

FIGURE 4-5 Transfer activity co-efBcients for halide and alkaline-metal ions in acetonitrile as solvent. (From Coetzee and Campion. ) 

If other substances in solution can serve the same function, heteroconjugation may occur, as in the formation of BrHI . Owing to the poor solvating ability of acetonitrile, many possible chemical species may be present in an acetonitrile solution of a base B and an acid HA. Among them are SH2 (solvated proton), (HA), or (B). (self-association), SH2 A, A(HA), (polyhomoconjugation), BH A , BH AHA , and BHB A, as well as the simple ions A and BH.  

Activity coefficients for ions in acetonitrile may be calculated from 

For rigorous calculations of equilibrium concentrations of the various species in solution, values of the constants for all the equilibria involved are needed use of a computer may be indicated. For example, SeUers, Eller, and Caruso developed a theoretical model that assumes the existence of a wide variety of species such as can be found in acetonitrile solution. Eleven equations in eleven unknowns are solved with the aid of a computer to give information about conductance, potentiometric, and photometric titration curves in nonaqueous media. 

In the following discussions of expressions relating hydrogen ion concentration to various points in the titration of a weak base with an acid, several simplifying assumptions are made  

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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]. 

M.K. Chantooni, Jr. and I.M. Kolthoff, Acid-base equilibria in methanol, acetonitrile, and dimethyl sulfoxide in acids and salts of oxalic acid and homologs, fumaric and o-phthalic acids. Transfer activity coefficients of acids and ions, J. Phys. Chem. 79 (1975), pp. 1176-1182. 

With an analytical method in hand, kinetic studies were iiutiated. The self-decomposition kinetic studies were studied first by using isolated adduct 3 in a homogenous Me2NAc solution, which provided a simplified system by eliminating all the acid-base equilibria that exist in the one-pot process solution. As mentioned above, preparation of 5 is equally effective in either acetonitrile or Me2NAc. As shown in Figure 21.3, a first-order dependence in the acid form 3 (HA) was clearly observed, with a ko,j of 4.09 x Iti s. ... 

A mechanistic study of acetophenone keto-enol tautomerism has been reported, and intramolecular and external factors determining the enol-enol equilibria in the cw-enol forms of 1,3-dicarbonyl compounds have been analysed. The effects of substituents, solvents, concentration, and temperature on the tautomerization of ethyl 3-oxobutyrate and its 2-alkyl derivatives have been studied, and the keto-enol tautomerism of mono-substituted phenylpyruvic acids has been investigated. Equilibrium constants have been measured for the keto-enol tautomers of 2-, 3- and 4-phenylacetylpyridines in aqueous solution. A procedure has been developed for the acylation of phosphoryl- and thiophosphoryl-acetonitriles under phase-transfer catalysis conditions, and the keto-enol tautomerism of the resulting phosphoryl(thiophosphoryl)-substituted acylacetonitriles has been studied. The equilibrium (388) (389) has been catalysed by acid, base and by iron(III). Whereas... 

It should be noted that self-ionisation is not an essential prerequisite for a satisfactory polar solvent. Liquids such as acetonitrile CH3CN or dimethylsulphoxide SO(CH3)2 appear not to ionise but they make very useful solvents for electrolytes as well as for polar molecular substances. As with H20, NH3, H2S04 etc., they owe their solvent powers to their polarity, leading to dipole-dipole interaction in the case of polar molecules as solutes and ion-dipole attraction in the case of electrolytes. There may in addition be considerable covalent bonding, via coordinate bond formation, in the case of cations. In solvents which do undergo appreciable self-ionisation, coordination often needs to be considered explicitly in discussing acid/base and other reactions and equilibria. 

The pJCj values are now available for many hydride complexes. Extensive tables have been compiled recently by Bullock and by Tilset. The rate of proton transfer to and from transition metals is rather slow (see below), so it is often possible to detect separate NMR signals for M-H and M , and tiius to determine the position of proton transfer equilibria between hydride complexes (M-H) and bases (B), or metal bases (M") and organic acids (HA). The pX values in Table 3.1 have been obtained in acetonitrile, an excellent solvent for acid-base chemistry because it solvates cations well enough to minimize ion pair formation it is both a weak acid and a weak base, with a very low autoprotolysis constant (ion product). ... 

Organic solvents influence the ionization constants of weak acids or bases in several ways (note that they influence the analytes and the buffer as well). Concerning ionization equilibria, an important solvent property is the basicity (in comparison to water), which reflects the interaction with the proton. From the most common solvents, the lower alcohols and acetonitrile are less basic than water. Dimethyl sulfoxide is clearly more basic. However, stabilization of all particles involved in the acido-basic equilibrium is decisive for the pKa shift as well. For neutral acids of type HA, the particles are the free, molecular acid, and the anion, A . In the equilibrium of bases, B, stabilization of B and its conjugated acid, HB, takes place. As most solvents have a lower stabilization ability toward anions (compared to water), they shift the pK values of adds of type HA to higher values in general. No such clear direction of the change is found for the pK values of bases however, they undergo less pronounced shifts. 

Arnett (1963a) is using heats of solution to determine the position of protonation equilibria. The method is simple in principle. If a weak base does not protonate when added to various concentrations of acid, the heat of solution will be approximately constant as exemplified by the behavior of acetonitrile shown in Fig. 2. If a hypothetical base protonated at intermediate acidities, a curve such as the dashed one in Fig. 2 could occur. In such an idealized situation, cb/cbh" could be estimated at each acidity. 

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