Acid dissociation constant of weak acids

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... 

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. 

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 ... 

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

Conductance measurements also have been used for the estimation of dissociation constants of weak electrolytes. If we use acetic acid as an example, we find that the equivalent conductance A shows a strong dependence on concentration, as illustrated in Figure 20.2. The rapid decline in A with increasing concentration is largely from a decrease in the fraction of dissociated molecules. 

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. 

Table 2.10 Dissociation constants of weak arsenic acids. pKa = —logw Ka. 

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... 

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. 

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. 

This section contains a discussion of pH scales and the electrometric measurement of pH, tables of assigned pH values for various buffered solutions, and tables (about 600 entries) of thermodynamic acid dissociation constants of weak organic acids (some as a function of temperature). 

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. 

The conductance of weak electrolytes, e.g. weak acids, is also increased under the influence of high fields. This dissociation field effect, or second Wien effect, is caused quite differently to that described for strong electrolytes. The high field in this case changes the values of the dissociation constants of weak electrolytes. For an acid dissociation... 

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. 

Determination of the dissociation constants of acids and bases from the change of absorption spectra with pH. The spectrochemical method is particularly valuable for very weak bases, such as aromatic hydrocarbons and carbonyl compounds which require high concentrations of strong mineral acid in order to be converted into the conjugate acid to a measurable extent. 

Partanen, J. I. Karki, M. H. Determination of the Thermodynamic Dissociation Constant of a Weak Acid by Potentiometric Acid-Base Titration, /. Chem. Educ. 1994,... 

Ethylenediaminetetraacetic acid, H4Y, is a tetraprotic weak acid with successive acid dissociation constants of 0.010,... 

In the discussion of the relative acidity of carboxylic acids in Chapter 1, the thermodynamic acidity, expressed as the acid dissociation constant, was taken as the measure of acidity. It is straightforward to determine dissociation constants of such adds in aqueous solution by measurement of the titration curve with a pH-sensitive electrode (pH meter). Determination of the acidity of carbon acids is more difficult. Because most are very weak acids, very strong bases are required to cause deprotonation. Water and alcohols are far more acidic than most hydrocarbons and are unsuitable solvents for generation of hydrocarbon anions. Any strong base will deprotonate the solvent rather than the hydrocarbon. For synthetic purposes, aprotic solvents such as ether, tetrahydrofuran (THF), and dimethoxyethane (DME) are used, but for equilibrium measurements solvents that promote dissociation of ion pairs and ion clusters are preferred. Weakly acidic solvents such as DMSO and cyclohexylamine are used in the preparation of strongly basic carbanions. The high polarity and cation-solvating ability of DMSO facilitate dissociation... 

The shapes of the titration curves of weak electrolytes are identical, as Figure 2.13 reveals. Note, however, that the midpoints of the different curves vary in a way that characterizes the particular electrolytes. The pV, for acetic acid is 4.76, the pV, for imidazole is 6.99, and that for ammonium is 9.25. These pV, values are directly related to the dissociation constants of these substances, or, viewed the other way, to the relative affinities of the conjugate bases for protons. NH3 has a high affinity for protons compared to Ac NH4 is a poor acid compared to HAc. 

Potentiometric titration curves The procedure involves the addition of a salt of a weak acid to the resin and the determination of the pH of the equilibrated solution. Table 9 shows the pK values of the OH groups and dissociation constants of the studied resin. The first ionization occurs at a pH slightly higher than that of sul-... 

Different Types of Proton Transfers. Molecular Ions. The Electrostatic Energy. The ZwiUertons of Amino Acids. Aviopro-tolysis of the Solvent. The Dissociation Constant of a Weak Acid. Variation of the Equilibrium Constant with Temperature. Proton Transfers of Class I. Proton Transfers of Classes II, III, and IV. The Temperature at Which In Kx Passes through Its Maximum. Comparison between Theory and Experiment. A Chart of Occupied and Vacant Proton Levels. 

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