Dissociation constants of weak acids

Table 1.7 Dissociation constants of weak acids and bases at 25°C. (From CRC Handbook of Chemistry and Physics)... 

Alternatively one could suggest that a micellized sulfuric or sulfonic acid is not strong. For example, apparent acid dissociation constants of weak acids decrease when the acids are bound to anionic micelles (Hartley, 1948), and the rapid hydrolysis of micellized alkyl sulfates at low pH is consistent with... 

Dissociation constants of weak acids and bases can be determined in the laboratory. However, it is easy to find the or K values of acids or bases by using a simple mathematical expression between and K[ of conjugate acid - base. The multiplication of and K of conjugate acid - base is Kw. 

The dissociation constants of weak acids are determined, in the described manner, according to conductivity measurements. From the diagram of the dissociation of acetic acid it follows that the intercept on the axis of ordinates log K c = log Kc = —4.756, from which the true dissociation constant Kt = = 1.75.1(T5 at 25° C. 

Dissociation constants of weak acids can also be derived from measurements on unbuffered cells of the type... 

The dissociation constants of weak acids or weak bases are often determined by monitoring the pH of the solution while the acid or base is being titrated. A pH meter with a glass pH electrode (see Section 21D-3) is used for the measurements. For an acid, the measured pH when the acid is exactly half neutralized is numerically equal to pA a- For a weak base, the pH at half titration must be converted to pOH, which is then equal to pf<(b-... 

On the other hand, the dissociation constants of weak acids are readily available. A linear correlation between stability constants of surface complexes of 6... 

DAS/NAR] Dasgupta, P. K., Nara, O., Measurement of acid dissociation constants of weak acids by cation exchange and conductometry. Anal. Chem., 62, (1990), 1117-1122. Cited on pages 141, 142. 

An important conclusion has been arrived at by McBam and Kam (loc at) in a paper with the following self-explanatory title The effect of salts on the vapour pressure and degree of dissociation of acetic acid in solution An expenmental refutation of the hypothesis that neutral salts increase the dissociation constants of weak acids and bases ... 

If complexes and chelates involve intimate chemical interactions, the extent of association should reflect the chemical nature of the ions involved. Equilibrium constants should be different, and possibly even very grossly different, for equilibria which superficially seem very similar and alike, for instance, association of one species with ions of similar size and charge. The situation is reminiscent of that found for the dissociation constants of weak acids and bases where the magnitude of the equihbrium constants depends on the chemical nature of the species involved. It is in stark contrast to that expected for the formation of ion pairs, where the magnitude of the association constant is expected to be independent of the chemical nature of the ions involved. 

The range of values for the acid-dissociation constants of weak acids extends over many orders of magnitude. Listed below are some benchmark values for typical weak acids to give you a general idea of the fraction of HA molecules that dissociate into ions ... 

S.2.2 Weak Bases In the discussion of dissociation constants of weak acids and bases (see Physical Chemical Properties, Chapter 2) it was shown that the pK of the conjugate acid of a base defined the following equilibrium ... 

The conductance method is satisfactory only if the solvent can be rigorously purified. Through failure to appreciate this, the first values of pKa of picric acid in acetonitrile proved to be much too small, 5.6 and 8.9 as compared with 11.0 from electromotive force measurements on buffered solutions. D Aprano and Fuoss found that acetonitrile having a satisfactory specific conductance of about 10 cm still contained a trace of ammonia. This was converted to ammonium pic-rate when acid was added to the solvent giving a spurious contribution to the conductance of picric acid solutions. This discovery moved them to make the flat assertion that dissociation constants of weak acids cannot be determined in aprotic solvents conductimetrically . This may be an overly pessimistic view, conductance values of pKa for acids in di-methylsulphoxide and dimethylformamide agree well with those from spectrophotometric and electromotive force measurements. Approximate values of pKa and pKf can be obtained from conductometric titrations of a weak acid with a weak base. ... 

It may be expected, that also introduction of other so-called negative groups (which in general increase also the dissociation constant of weak acids) have a similar influence. 

Dissociation constants of weak acids increase with increasing applied pressure and the pressure effect for a particular step of dissociation is related by the thermodynamic formula (see Eq. (3.21)) [20, 30]... 

Equilibrium constants at 25" C Dissociation constants of weak acids... 

Experimental determinations of the conducting properties of electrolyte solutions are important essentially in two respects. Firstly, it is possible to study quantitatively the effects of interionic forces, degrees of dissociation and the extent of ion-pairing. Secondly, conductance values may be used to determine quantities such as solubilities of sparingly soluble salts, ionic products of self-ionizing solvents, dissociation constants of weak acids and to form the basis for conductimetric titration methods. 

Voltammetric methods also provide a convenient approach to establish the thermodynamic reversibility of an electrode reaction and for the evaluation of the electron stoichiometry for the electrode reaction. As outlined in earlier sections, the standard electrode potential, the dissociation constants of weak acids and bases, solubility products, and the formation constants of complex ions can be evaluated from polarographic half-wave potentials, if the electrode process is reversible. Furthermore, studies of half-wave potentials as a function of ligand concentration provide the means to determine the formula of a metal complex. 

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