Acid-dissociation constant strength

The strength of a weak acid is measured by its acid dissociation constant, which IS the equilibrium constant for its ionization m aqueous solution... 

An inflection point in a pH-rate profile suggests a change in the nature of the reaction caused by a change in the pH of the medium. The usual reason for this behavior is an acid-base equilibrium of a reactant. Here we consider the simplest such system, in which the substrate is a monobasic acid (or monoacidic base). It is pertinent to consider the mathematical nature of the acid-base equilibrium. Let HS represent a weak acid. (The charge type is irrelevant.) The acid dissociation constant, = [H ][S ]/[HS], is taken to be appropriate to the conditions (temperature, ionic strength, solvent) of the kinetic experiments. The fractions of solute in the conjugate acid and base forms are given by... 

Consider a neutral base B of such strength that it can be protonated in dilute aqueous solution in the acidic range, say pH 1-2. In the conventional manner the acid dissociation constant /ibh + is defined. 

Apparent partition coefficient (logZ)) at an ionic strength of / = 0.02M, log P value of the neutral microspecies and the acidic dissociation constant of 5 was calculated (97ANC4143). The distribution coefficient of 5 was determined between 1-octanol and universal buffer in the pH range 3-10 at a... 

The strength of weak acids is expressed by pA, the negative log of the acid dissociation constant. Strong acids have low pA values and weak acids have high pAl, values. 

Metal-complex stability is also related to the basic strength of the ligand entity. For a series of 1 2 complexes of the bidentate naphthylazophenol ligand (5.64) with copper(II) ion, the acidic dissociation constants (pKa) are linearly related to the stability constants (log K1 2), the more acidic groups forming the less stable complexes. Thus where X = N02 in structure 5.64 then pKa = 8.1 and log K1 2 = 17.2, and where X = OCH3 then pKa = 8.5... 

Ka is known as the acid dissociation constant it is a measure of the strength of an acid in a particular solvent, which should be specified. 

The p/<, of a base is actually that of its conjugate acid. As the numeric value of the dissociation constant increases (i.e., pKa decreases), the acid strength increases. Conversely, as the acid dissociation constant of a base (that of its conjugate acid) increases, the strength of the base decreases. For a more accurate definition of dissociation constants, each concentration term must be replaced by thermodynamic activity. In dilute solutions, concentration of each species is taken to be equal to activity. Activity-based dissociation constants are true equilibrium constants and depend only on temperature. Dissociation constants measured by spectroscopy are concentration dissociation constants." Most piCa values in the pharmaceutical literature are measured by ignoring activity effects and therefore are actually concentration dissociation constants or apparent dissociation constants. It is customary to report dissociation constant values at 25°C. 

The earliest LFER, advanced by Bronsted, correlates the acid dissociation constant and base strength (1/A h) of species with its effectiveness as a catalyst in general acid (At h) and base (Atgl-catalyzed reactions respectively. The relationships take the form... 

The bipyridines are dibasic, and the two acid dissociation constants Ki and K2, for all the bipyridines have been determined. Typical values are recorded in Table I. There has been considerable interest in the first dissociation constants Ki of 2,2 -bipyridine and substituted 2,2 -bipyridines because of their use as metal complexing agents. In general, the order of relative basic strengths of derivatives of 2,2 -bipyridine is as expected. Electron-attracting substituents reduce the basicity, whereas electron-donating substituents increase the basicity of the molecule. " The dissociation constants of several substituted bipyridines correlate well with the Hammet equation. 2,2 -Bipyridines with an electron-donating substituent at position 4 are monoprotonated at N-1 and not at... 

The larger the acid dissociation constant, the stronger is the acid. Hydrochloric acid has an acid dissociation constant of 10, whereas acetic acid has an acid dissociation constant of only 1.74 x 10 . For convenience, the strength of an acid is generally indicated hy its pA a value rather than its A a value The of hydrochloric acid, strong acid, is —7, and the pA a of acetic acid, much weaker acid, is 4.76. 

Kg refers to the acid dissociation constant which is the measure of an acid s strength. Some references call the acid ionization constant. 

Bruice and Schmir (3) have shown that for a series of imidazole derivatives, klm depends on the base strength of the catalyst and since pKA is an approximate measure of base strength, the value of klm should increase with increase in pKA. Table I shows that this is indeed the case. Imidazole, pKA = 7.08, has a catalytic constant eight times larger than that of benzimidazole, pKA = 5.53. Bronsted and Guggenheim (2) have obtained a linear relationship between log k/ and pKA for a series of carboxylic acids in the pKA range of 2 to 5, where kB is the carboxvlate anion basic catalytic constant for the mutarotation of glucose and Ka is the acid dissociation constant of the acid. Our results for imidazole and benzimidazole fit fairly well into the Bronsted plot. 

It is possible to compare the strengths of weak acids by the values of their acid dissociation constants Ka. Figure 3.1 shows the titration curves for acids (HA or BH+) of different Ka values. The ordinate shows poH, which is defined by paH = -loga(SI I)). paH corresponds to the pH in aqueous solutions (see Section 3.2). The poH of non-aqueous solutions can be measured with a glass pH electrode or some other pH sensors (see Sections 3.2.1 and 6.2). For the mixture of a weak acid A and its conjugate base B, poH can be expressed by the Henderson-Hassel-balch equation ... 

The strength of an acid is usually described by the acid dissociation constant Ka... 

The acid dissociation constants for tri- and tetranuclear aqua chromium(III) species are summarized in Table XXIII. The structures of these species are not known. If, however, it is assumed that they have linear structures, such as structures 4a, 7b, and 7c shown in Fig. 1, then the observed acid strengths can be rationalized in terms of a -and /1-type hydrogen bond interactions, as discussed recently (118). 

The inductive effect of one carboxyl group is expected to enhance the acidity of the other. In Table 18-4 we see that the acid strength of the dicarboxylic acids, as measured by the first acid-dissociation constant, K1, is higher than that of ethanoic acid (Ka = 1.5 X 10-5) and decreases with increasing number of bonds between the two carboxyl groups. The second acid-dissociation constant, K2, is smaller than Ka for ethanoic acid (with the exception of oxalic acid) because it is more difficult to remove a proton under the electrostatic attraction of the nearby carboxylate anion (see Section 18-2C). 

In many reference works, it is customary to express the strengths of organic bases not as Kb values but as the acid-dissociation constants, Ka (or pica s) for the corresponding conjugate acids. These Ka values are then the acid constants of the corresponding ammonium ions in aqueous solution (Equation 23-4) ... 

As discussed in Section 3.10.3, in the gas phase the basicity of simple amines follows the order NMe3 > NHMe2 > NH2Me > NH3 because of the electron donating effect of the methyl (Me) groups. In solution, however, we can define a basicity constant as the equilibrium constant for the reaction shown in Equation 3.4. Note it is important to specify temperature, solvent (usually water) and solution ionic strength, 1 Basicity constants are related to the acid dissociation constants (/Q of the base s conjugate acid via the dissociation constant of water, K = 10 14 at 25 °C. Thus Kbx K = Kw. 

The acid dissociation constant Ka is independent of ionic strength, but the acid dissociation constant Kc depends on the ionic strength, as indicated by equation 1.2-7. The equilibrium constant expression in equation 1.2-7 will be used in the rest of the book, but the subscript c will be omitted. This will make it possible for us to deal with concentrations of species, rather than activities. 

Since the acid dissociation constants are known, the value of Krcf can be calculated from the value of K at a pH in the neutral region in the absence of metal ions by using equation 1.4-3. Values of Kief at zero ionic strength are given in Table 1.2 for six reference reactions. 

In this chapter we have seen that acid dissociation constants are needed to calculate the dependence of apparent equilibrium constants on pH. In Chapter 3 we will discuss the calculation of the effects of ionic strength and temperature on acid dissociation constants. The database described later can be used to calculate pKs of reactants at 298.15 K at desired ionic strengths. Because of the importance of pKs of weak acids, Table 1.3 is provided here. More experimental measurements of acid dissociation constants and dissociation constants of complex ions with metal ions are needed because they are essential for the interpretation of experimental equilibrium constants and heats of reactions. A major database of acid dissociation constants and dissociation constants of metal ion complexes is provided by Martell, Smith, and Motekaitis (2001). 

The equilibrium constants involved in the reaction Fe3+ + 3 cat2" Fe(cat)33- were determined as follows. An aqueous solution of Fe3+ (5.5 X 10-3M) and catechol (1.48 X 10-2M), initially made basic with the addition of KOH, was titrated with 1.24M HC1 under an oxygen-free atmosphere at 22° and ionic strength (KC1) 0.16-0.22M (Figure 12). The acid dissociation constants for catechol were determined independently (under similar experimental conditions) to be pKai = 9.38 and... 

This table can be helpful in estimating the pATs of other weak acids from their structures. In using this table it is important to remember that -log [H ] is used in the expression for the acid dissociation constant in terms of pH. To obtain pKs based on -log y (H ) [H ], add 0,0.08, 0.11,0.12, and 0.14 at ionic strengths of 0, 0.05, 0.10,0.15, and 0.25 M, respectively, at 298.15 K as indicated by Table 1.3. PaddedForm rounds the output to two figures to the right of the decimal point. There is a list of full names of reactants in the Appendix of this book. The reactants bpg, nmn, pep, and prpp are bisphosphoglycerate, nicotinamidemononucleotide, phosphoenolpyruvate, and 5-phosphoribosyl-alpha-pyrophosphate, respectively. 

Calorimetric measurements on a reaction like the hydrolysis of ATP yields Ar // ° at the experimental T, pH, and ionic strength. The calorimetric heat of reaction Aj/Zc must be corrected for the heat effect of the hydrogen ions produced by the enzyme-catalyzed reaction on the acid dissociation of the buffer, as described in Chapter 15. If Zf is measured at several temperatures and the acid dissociation constants of all the reactants are known at these temperature, the equilibrium constant K for the reference reaction can be calculated at each temperature. Plotting InAT versus 1/Tyields, which is given by... 

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