Acidity constants of simple carbonyl compounds

When they are sufficiently low, the pAa (pA p values of carbonyl compounds (corresponding to enolate formation) in water can be obtained directly from pH measurements of partly neutralised compounds or by spectroscopic determination of the ratio of the ionised and unionised forms in weakly basic media. Indeed, most of the acidity constants of di- or tri-carbonyl compounds have been determined by these methods. However, such direct determinations are not valid for monocarbonyl compounds for which the pA values are known to be usually higher than that of water. For these very weak acids the competitive and acidity function approaches have been used, although both have serious limitations. 

Scales for the relative acidities of C—H compounds, including some carbonyl compounds, have been proposed since the pioneer work of Conant and Wheland (1932) and of McEwen (1936). All of these have been established from polarimetric and spectrophotometric measurements in aprotic solvents of low dielectric constant of the ratio [HIn]/[In ] in equilibria of type 

In contrast, the acidity function approach, which involves a study of the dependence of the ionisation ratio ([S ]/[HS]) on an appropriate acidity function, was recently used for substituted acetophenones and some aliphatic ketones (Cockerill et al., 1974 Earls et al., 1975 Kankaanpera et al., 1978). All these studies used the property that the addition of DMSO increases the basicity of aqueous hydroxide ion solutions p/ s values were determined from (46) by plotting the ionisation ratio versus the H acidity function for 

OH- solutions in DMSO-water mixtures. For acetone, the ionisation ratios were measured by spectrophotometry, but in other cases an indirect kinetic method was used. This latter is based on monitoring the rates of detritiation of a standard labelled carbon acid (HS1) both in the presence and absence of a second acid HS2 (ketone). The dissociation of HS2 brings about a decrease in hydroxide ion concentration, and, since the rate of detritiation of HS1 is proportional to [OH ], the consequent decrease in rate can be related to [(S2) ]/[HS2]. Data listed in Table 7 exhibit large variations with structure, far larger than those expected from ionisation rate constants if an enolate-like transition state were assumed (see p. 34). 

Although the definition of acidity functions implies that the p.Ka-values determined in this manner refer to a standard state of infinite dilution in water, the acidity function approach has serious limitations since the necessary assumptions concerning activity coefficients may not always be valid. Moreover, it can be asserted that the pKf values obtained are about 2 units too high. Indeed, if the data from Fig. 4 were associated with the pK value of 21.7 obtained by Kankaanpera et al., this would give a rate constant for iodine addition to the enolate equal to about 5 x 10n dm3 mol-1 s-1, i.e. 100 times higher than that expected for a diffusion-controlled process. 

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