Acid-ionization constants experimental determination

Ionization constants are determined by experimental measurements of equilibrium concentrations. For example, to determine Kx for acetic acid, we... 

The acid ionization constant for a weak acid represents a small value. To determine the numerical value of the ionization constant for acetic acid at a specific temperature, the equilibrium concentrations of H3O+ ions, CH3C00 ions, and CH3COOH molecules must be known. The ionization of a molecule of CH3COOH in water yields one H3O+ ion and one CH3C00 ion. These concentrations can therefore be found experimentally by measuring the pH of the solution. 

It can be seen from Table 2 that the intrinsic values of the pK s are close to the model compound value that we use for Cys(8.3), and that interactions with surrounding titratable residues are responsible for the final apparent values of the ionization constants. It can also be seen that the best agreement with the experimental value is obtained for the YPT structure suplemented with the 27 N-terminal amino acids, although both the original YPT structure and the one with the crystal water molecule give values close to the experimentally determined one. Minimization, however, makes the agreement worse, probably because it w s done without the presence of any solvent molecules, which are important for the residues on the surface of the protein. For the YTS structure, which refers to the protein crystallized with an SO4 ion, the results with and without the ion included in the calculations, arc far from the experimental value. This may indicate that con-... 

C). The value of a may be determined from its definition in terms of the ionization constant if the appropriate benzoic acid derivative has been measured. Such a values may then be used to determine p values for other reactions, and these p values in turn lead to the possibility of determining new o values. Equation 7.4.24 implies that a plot of log KA or log kA versus oA should be linear for a given series of reactions involving the same reactive groups. Extensive experimental evidence attests to this relation. Table 7.2 contains some values that have been reported in the literature. Extensive tabulations... 

A potentiometric method for determination of ionization constants for weak acids and bases in mixed solvents and for determination of solubility product constants in mixed solvents is described. The method utilizes glass electrodes, is rapid and convenient, and gives results in agreement with corresponding values from the literature. After describing the experimental details of the method, we present results of its application to three types of ionization equilibria. These results include a study of the thermodynamics of ionization of acetic acid, benzoic acid, phenol, water, and silver chloride in aqueous mixtures of acetone, tetrahydrofuran, and ethanol. The solvent compositions in these studies were varied from 0 to ca. 70 mass % nonaqueous component, and measurements were made at several temperatures between 10° and 40°C. 

In this equation, pK° is the pK of the unsubstituted molecule, o refers to the constant assigned to a respective substituent and p is a constant assigned to the particular ionizable group. The approach was later refined by Taft et al extending the theory and equation to aliphatic and other systems. The principles and methodologies of a classic experimental pK determination are well known. Ionization constants for an n-protic weak acid may be represented by the following generic definitions ... 

Ionization constants for weak acids (and bases) must be calculated from experimentally determined data. Measurements of pH, conductivity, or depression of freezing point provide data from which these constants can be calculated. 

It must be remarked, however, that the curves shown in Fig. 3 are hardly typical. Two definite points of inflection are observed for a dibasic acid only if the two ionization constants are of quite different orders of magnitude, Auerbach and Smolczyk7 have shown, theoretically, that the two points of inflection will not appear unless the ratio Ki/E2 is greater than 16. However, it is safe to say that the ratio must be considerably greater than that if the inflection points are to be obtained experimentally with any accuracy. Also, if the two ionization constants are not quite different the values of the constants determined by the potentiometric titration method described above may be somewhat in error, as is shown by the authors just mentioned and by Simms. ... 

The detailed treatment of pH effects is complicated, since various possibilities have to be taken into account. A very simple situation occurs when the reaction follows the original Michaelis-Menten mechanism, in which the enzyme and substrate form an addition complex which breaks down in a single stage. The enzyme-substrate complex may also exist in three states of ionization, and perhaps only the interrnediate form is capable of giving rise to products. This simple case is represented in Figure 10.8. Intuitively we can see that at low substrate concentrations, when the enzyme exists mainly in the free form, the pH behavior will be controlled by the ionization of the free enzyme. Analysis of the experimental pH dependence at low substrate concentrations therefore will allow us to determine the acid dissociation constants and for the free enzyme. On the other hand, if we saturate the enzyme with substrate, analysis of the pH behavior will now give the values of K and K, which relate to the ionization of the enzyme-substrate complex. 

In 1 the published dissociation constant of iminodiacetic acid is used since reasonable agreement between the two has already been demonstrated (5). The experimentally determined composition of the resin at equilibrium is used directly in the equation to represent the a vity of acid [H R] and metal chelate [MR] forms of the resin at equilibrium. Since ionization is negligible in the metal and acid forms these numbers are an accurate estimate of their concentration values. For this reason, also, there is little interaction between them to disturb their ideal behavior and the measured concentration ratio is not drastically different from their thermodynamic activity ratio. 

The concept of pH, however, is not applicable in such nonaqueous systems or in concentrated acid solutions. A new quantitative scale, therefore, was needed (5,9-11). The most useful and widely accepted method was proposed by Hammett and Deyrup in 1932 (12). They defined ho by equation 24, which can be determined experimentally by adding the neutral base B in low concentrations to an acid solution (BH+ is the acidic form of the indicator, Ksa+ is the thermodynamic equilibrium constant for BH+, and [BH+]/[B] is the ionization ratio generally determined spectrophotometrically). Equation 24 is usually written in the logarithmic form (eq. 25), where the quantity Hq is termed the Hammett acidity function. Since in dilute solutions of acids Abh+ is expressed as in equation 26, the Hammett acidity function becomes equal to pH. In concentrated solutions, however, Hq differs considerably from pH and this can be formally expressed by inserting activity coefficients in equation 24 (eq. 27). 

The ionization constant for a weak acid is usually determined experimentally by one of two methods. In one method, the electrical conductivity or some colligative property of a solution of the acid is measured to obtain its degree of ionization. The degree... 

Recall that the behavior of polyacids depends on macroscopic and microscopic ionization constants. The former can be determined experimentally. They globally quantify the more or less simultaneous ionization of several acid functions. Their value cannot be assigned in totality to only one acidity constant if the initial acid is dissymmetric. Let us consider malic acid (Fig. 11.1). It exhibits two different microscopic ionization methods (Fig. 11.3). 

The individual ionization constants for pairs of tautomers and the constant for the equilibrium between them can be deduced from the experimentally determined ionization constants in another way. It will be seen (p. 156) that the effect of substituents on the basic strength of the pyridine nitrogen atom can be represented by the Hammett equation. Accordingly, the ionization constant for the NM form (T b) can be estimated from the Hammett substituent constant of the group — A H. Combined with the experimentally determined Ki this leads to estimates of the remaining quantities. In this way Jaffe 2 deduced that nicotinic and isonicotinic acids exist predominantly as zwitterions in aqueous solution. Similarly, Bryson s obtained the values pi B 5-0 and Kt = 0 78 for 3-hydroxypyridine (cf. Tables 5,10 and 5.12), Hammett correlations have been used in evaluating the tautomeric equilibria in a number of 2-arylsulphonamidopyridines . ... 

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 equilibrium constants Ka, which are determined experimentally, are the macroscopic ones. They quantify together the two ionization ways. They are overall acid-dissociation equilibrium constants. They are defined as follows ... 

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