Acid-base equilibrium constant for

One can write acid-base equilibrium constants for the species in the inner compact layer and ion pair association constants for the outer compact layer. In these constants, the concentration or activity of an ion is related to that in the bulk by a term e p(-erp/kT), where yp is the potential appropriate to the layer [25]. The charge density in both layers is given by the algebraic sum of the ions present per unit area, which is related to the number of ions removed from solution by, for example, a pH titration. If the capacity of the layers can be estimated, one has a relationship between the charge density and potential and thence to the experimentally measurable zeta potential [26]. 

The acid-base equilibrium constant for the Me residue can be determined by spectroscopic pH titration. An example for the titration is shown in Fig. 2. The electrostatic potential

difference between the apparent pK on the charged surface (pKobs) and that on an intrinsic neutral surface (pK1) by... 

The logarithmic form of the acid-base equilibrium constant for water will then be... 

With the usual type of dibasic acid the equilibrium constant for the second step is always smaller than that for the first (Kx>K2)y but the diazonium ion represents another kind of acid in which the second constant is greater than the first (K2 > Kf), Schwarzenbach (1943) was the first to discover analogous abnormal acid-base equilibria and he explained under what circumstances the phenomenon can occur (for a historical account of Schwarzenbach s work see Zollinger, 1992). 

An indication of the nature of the transition state in aromatic substitution is provided by the existence of some extrathermodynamic relationships among rate and acid-base equilibrium constants. Thus a simple linear relationship exists between the logarithms of the relative rates of halogenation of the methylbenzenes and the logarithms of the relative basicities of the hydrocarbons toward HF-BFS (or-complex equilibrium).288 270 A similar relationship with the basicities toward HC1 ( -complex equilibrium) is much less precise. The jr-complex is therefore a poorer model for the substitution transition state than is the [Pg.150]

With reference to a solvent, this term is usually restricted to Brpnsted acids. If the solvent is water, the pH value of the solution is a good measure of the proton-donating ability of the solvent, provided that the concentration of the solute is not too high. For concentrated solutions or for mixtures of solvents, the acidity of the solvent is best indicated by use of an acidity function. See Degree of Dissociation Henderson-Hasselbalch Equation Acid-Base Equilibrium Constants Bronsted Theory Lewis Acid Acidity Function Leveling Effect... 

Simplificaion of the acidity or basicity scale in low polar media consists of a series of equilibrium constants corresponding to the reaction above, with either B or HA being the reference compound. The acid-base association constants for the equilibria of substances interacting within these solvents, will be seen to fall mainly in the range logKBHA = 2 to 7 (Davies, 1968). 

These two solids will, of course, exhibit different catalytic effects. The equilibrium constant of reaction (XIV) is KAgCiIKAgBT which in water at 25°C equals 500. The tendency for reactions like (XIII) and (XIV) to occur can thus be calculated from known solubility products and stability constants. In the case of solids like oxides or carbonates, one must also take into account the relevant acid-base equilibrium constants because here dissolution can occur if the solution is made sufficiently acid or alkaline. 

Influence of Ionic Strength. It should be noted that eorrections to take aecount of ionic strength as discussed in Section 3.2.1.2 apply not only to the acid-base equilibrium constants but also to the stability constants for complex formation. 

Sections 3.3.1 and 4.2.1 dealt with Bronsted acid/base equilibria in which the solvent itself is involved in the chemical reaction as either an acid or a base. This Section describes some examples of solvent effects on proton-transfer (PT) reactions in which the solvent does not intervene directly as a reaction partner. New interest in the investigation of such acid/base equilibria in non-aqueous solvents has been generated by the pioneering work of Barrow et al. [164]. He studied the acid/base reactions between carboxylic acids and amines in tetra- and trichloromethane. A more recent compilation of Bronsted acid/base equilibrium constants, determined in up to twelve dipolar aprotic solvents, demonstrates the appreciable solvent influence on acid ionization constants [264]. For example, the p.Ka value of benzoic acid varies from 4.2 in water, 11.0 in dimethyl sulfoxide, 12.3 in A,A-dimethylformamide, up to 20.7 in acetonitrile, that is by about 16 powers of ten [264]. 

For HA a weak acid the products of the dissociation are hydrogen ion (H" ) and an anion (A ), which is the conjugate base. Equilibrium constants for acids can be written in the following form ... 

Obviously, all acid-base equilibrium constants depend on the pH scale used. It is possible to convert approximately an equilibrium constant determined in one scale to that of another scale. The problem of different definitions of equilibrium constants needs attention when applying an infinite dilution scale complex formation constant,—for example, for CuC03(aq)—in a seawater medium. 

The acid-base equilibrium constants of the beta-blockers atenolol, oxprenolol, timolol, and labetalol were determined by automated potentiometric titrations. The pKg values were obtained in water-rich or water methanol medium (20% MeOH) to obviate the solubility problems associated with the compounds. The initial estimates of pKa values were obtained from Gran s method and then, were refined by the NYTIT and ZETA versions of the LETAGROP computer program. The resultant values were 9.4 (/ = 0.1 M KCI, 20% methanol) for atenolol, 9.6 (/ = 0.1 M KCI) for oxprenolol, 9.4 (/ = 0.1 M KCI, 20% methanol) for timolol and 7.4 and 9.4 (/ = 0.1 M KCI) for labetalol. The potentiometric method was found to be accurate and easily applicable. The operational criteria for applying the methodology are indicated. 

The importance of water as a medium for inorganic reactions stems not only from the fact that it is far more readily available than any other solvent, but also because of the abundance of accurate physicochemical data for aqueous solutions compared with the relative scarcity of such data for solutions in non-aqueous solvents. This chapter is concerned mainly with equilibria and we begin by reviewing calculations involving acid-base equilibrium constants. 

The reactivity sequence stated by Staudinger and Gaule was the very early qualitative precursor for the quantitative determination of the acid-base equilibrium constant of diazomethane by McGarrity and Smyth (1980). The major difficulty of... 

For TIOA with hydrochloric acid the concentration-based equilibrium constant for salt formation" according to reaction (8.2-6) is 1.51 x 10 and the equilibrium constant for amine-hydrochloride salt dimerization" is 8.0 M Combination of these parameters and the ion-complex stability constants with experimental metal-distribution data allows determination of the equilibrium constants for reactions (8.2-5) or (8.2-7). This completes the description of the amine-metal extraction-phase equilibria. For cobalt(II) in acidic sodium chloride solutions the equilibrium constant" for reaction (8.2-7) with TIOA is 2.0 X 10 and that for coppeifll) is 370 The corresponding value for zinc" is 7.5 x 10 Af -In spile of these relative values, the order of selectivity of TIOA for extraction of the metals is Zn > Cu > Co because of the relative extent of chloride complex formation. For the same reason, zinc stripping is difficult in this system, and copper has a tendency to be reduced to cuprous, which also complexes and extracts extensively. 

Binding constants are a special case of equilibrium constants such as acid-base equilibrium constants (best known as pATa and pAfb) and solubility equilibrium constants (Ks). Binding constants measure the bonding affinity between two or more molecules at equilibrium. In supramolecular chemistry, binding constants for host-guest complexation or host-host aggregation (e.g., dimerization) are usually... 

According to the Arrhenius definitions an acid ionizes m water to pro duce protons (H" ) and a base produces hydroxide ions (HO ) The strength of an acid is given by its equilibrium constant for ionization m aqueous solution... 

The carbon-metal bonds of organolithium and organomagnesium compounds have appreciable carbamomc character Carbanions rank among the strongest bases that we 11 see m this text Their conjugate acids are hydrocarbons—very weak acids indeed The equilibrium constants for ionization of hydrocarbons are much smaller than the s for water and alcohols thus hydrocarbons have much larger pA s... 

A species that can serve as both a proton donor and a proton acceptor is called amphiprotic. Whether an amphiprotic species behaves as an acid or as a base depends on the equilibrium constants for the two competing reactions. For bicarbonate, the acid dissociation constant for reaction 6.8... 

The equilibrium constant for equation 6.13 is K. Since equation 6.13 is obtained by adding together reactions 6.11 and 6.12, may also be expressed as the product of Ka for CH3COOH and Kb for CH3COO-. Thus, for a weak acid, HA, and its conjugate weak base, A-,... 

In a simple liquid-liquid extraction the solute is partitioned between two immiscible phases. In most cases one of the phases is aqueous, and the other phase is an organic solvent such as diethyl ether or chloroform. Because the phases are immiscible, they form two layers, with the denser phase on the bottom. The solute is initially present in one phase, but after extraction it is present in both phases. The efficiency of a liquid-liquid extraction is determined by the equilibrium constant for the solute s partitioning between the two phases. Extraction efficiency is also influenced by any secondary reactions involving the solute. Examples of secondary reactions include acid-base and complexation equilibria. 

Since the position of an acid-base equilibrium depends on the pH, the distribution ratio must also be pH-dependent. To derive an equation for D showing this dependency, we begin with the acid dissociation constant for HA. 

Although not commonly used, thermometric titrations have one distinct advantage over methods based on the direct or indirect monitoring of plT. As discussed earlier, visual indicators and potentiometric titration curves are limited by the magnitude of the relevant equilibrium constants. For example, the titration of boric acid, ITaBOa, for which is 5.8 X 10 °, yields a poorly defined equivalence point (Figure 9.15a). The enthalpy of neutralization for boric acid with NaOlT, however, is only 23% less than that for a strong acid (-42.7 kj/mol... 

Although this experiment is written as a dry-lab, it can be adapted to the laboratory. Details are given for the determination of the equilibrium constant for the binding of the Lewis base 1-methylimidazole to the Lewis acid cobalt(II)4-trifluoromethyl-o-phenylene-4,6-methoxysalicylideniminate in toluene. The equilibrium constant is found by a linear regression analysis of the absorbance data to a theoretical equilibrium model. 

The equilibrium constant for an acid-base indicator is determined by preparing three solutions, each of which has a total indicator concentration of 1.35 X lQ-5 M. The pH of the first solution is adjusted until it is acidic enough to ensure that only the acid form of the indicator is present, yielding an absorbance of 0.673. The absorbance of the second solution, whose pH was adjusted to give only the base form of the indicator, was measured at 0.118. The pH of the third solution was adjusted to 4.17 and had an absorbance of 0.439. What is the acidity constant for the acid-base indicator ... 

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