Acid-base equilibria relative strengths

The extent of the ionization step depends on the relative strength of the conjugate acid-conjugate base pairs. The amphiprotic properties of the solvent have an essential effect on the equilibrium constant of this reaction step. The extent of the dissociation step is influenced by the polarity of the solvent, increasing with the dielectric constant of the solvent. In water, all products of acid-base reactions of moderate to low concentrations are essentially completely dissociated into solvated ions (Pecsok et al., 1976). The dissociation step is suppressed by addition or substitution with cosolvents of lower polarity, e.g., alcohols in aqueous formulations. The ion-pair aggregates may have absorption spectra different from the dissociated species. Thus, the amphiprotic properties and polarity (expressed as the dielectric constant) of the solvent are essential for the acid-base equilibrium of the drug and thus the absorption spectrum of the compound. This subject is further discussed in Section 14.2.3. 

To express the relative strengths of an acid and its conjugate base (a conjugate acid-base pair ), we consider the special case of the ammonia proton transfer equilibrium, reaction C, for which the basicity constant was given earlier (Kb = [NH4+l[OH ]/ NH3]). Now let s consider the proton transfer equilibrium of ammonia s conjugate acid, NH4+, in water ... 

This is always the case for any two acids, and by measuring the positions of the equilibrium the relative strengths of acids and bases can be determined. Of course, if the two acids involved are close to each other in strength, a measurable reaction will occur from both sides, though the position of equilibrium will still be over to the... 

The above equation has an equilibrium constant that is less than 1. What are the relative strengths of the acids and bases ... 

These equations represent a transfer of a proton from A, (Acid,) to B2 (Base2). Reactions between acids and bases are hence termed protolytic reactions. All these reactions lead to equilibrium, in some cases the equilibrium may be shifted almost completely in one or another direction. The overall direction of these reactions depends on the relative strengths of acids and bases involved in these systems. 

Boric acid is a relatively weak acid compared to other conunon acids, as illustrated by the acid equilibrium constants given in Table 4. Boric acid has a similar acid strength to sihcic acid. Calculated pH values based on the boric acid equihbrium constant are significantly higher than those observed experimentally. This anomaly has been attributed to secondary equilibria between B(OH)3, B(OH)4, and polyborate species. Interestingly, the aqueous solubihty of boric acid can be increased by the addition of salts such as potassium chloride and sodium sulfate, but decreased by the addition of others salts, such as the chlorides of lithium and sodium. Basic anions and other nucleophiles such as fluorides and borates significantly increase boric acid solubility. 

Recently, the C-NMR method was also applied to these studies and the following order of the relative strengths of Lewis acids and bases, rdated to equilibrium (47) was given ... 

A proton transfer reaction represents an equilibrium. Because an acid donates a proton to a base, thus forming a conjugate acid and conjugate base, there are always two acids and two bases in the reaction mixture. Which pair of acids and bases is favored at equilibrium The position of the equilibrium depends on the relative strengths of the acids and bases. 

Any consideration of sovent effects on rates or equilibria must start from solvent activity coefficients, VI for reactants, transition states and products (Wiberg, 1964 Laidler, 1950 Parker, 1966). Once solvent activity coefficients are available, or can be predicted, it is highly probable, as indicated at the end of this article, that an enormous amoimt of information on the kinetics of reactions in solution and on equilibrium properties such as solubility, acid-base strength, ion-association, complexing, redox potentials and kinetics of reactions in different solvents (Parker, 1962, 1965a, 1966) can be reduced to a relatively small number of constants which can then be used in appropriate linear free energy relationships. 

If we know the relative strengths of two acids, we can predict the position of equilibrium between one of the acids and the conjugate base of the other, as illustrated in Example 15.7. 

The relative strengths of weakly basic solvents are evaluated from the extent of protonation of hexamethylbenzene by trifluoro-methanesulfonic acid (TFMSA) in those solvents or from the effect of added base on the same protonation in solution in trifluoroacetic acid (TFA), the weakest base investigated. The basicity TFA < di-fluoroacetic acid < dichloroacetic acid (DCA) < chloroacetic acid < acetic acid parallels the nucleophilicity. 2-Nitropropane appears to be a significantly stronger base than DC A by the first approach, although in the second type of measurement, the two have essentially equal basicity. The discrepancy is due to an interaction, possible for hydroxylic solvents such as DC A, with the anion of TFMSA. This anion stabilization is a determining factor of carbocationic reactivity in chemical reactions, including solvolysis. A distinction is made between carbocation stability, determined by structure, and persistence (existence at equilibrium, e.g., in superacids), determined by environment, that is, by anion stabilization. 

Equilibrium 6.1 illustrates that water can function as both a Bronsted acid and a Bronsted base. In the presence of other Bronsted acids or bases, the role of water depends on the relative strengths of the various species in solution. When HCl is bubbled into water, the gas dissolves and equilibrium 6.3 is established. 

When comparing two acids, several factors can be used to evaluate relative acid strength. In an acid-base reaction where acids are on both sides of the equilibrium (A and CA from above), this information can be useful to estimate the position of the equilibrium, which is essentially an estimation of Ka. 

Intermolecular interactions and complexes Dielectric measurements on interacting solutes in inert solvents provide information about molecule complex formation. Some such dipoles induced by intermolecular interactions and molecular complexes in benzene solution are listed in Table 1.3. The dipole moment of the complex is a function of the relative strengths of the acid and base and the intramolecular equilibrium is described by Eq. (49) ... 

Ka and Kb are particular types of equilibrium constants that give us an idea of the relative strengths of acids and bases, respectively. The acid-dissociation constant, Ka, is the equilibrium constant for the ionization of a weak acid to a hydrogen ion and its conjugate base. Likewise, the base-dissociation constant, Kb, is the equilibrium constant for the addition of a proton to a weak base by water to form its conjugate acid and an OH ion. 

In these reactions, there is a competition for protons between the two bases. Therefore, the relative amounts of the conjugate acid-base pairs that exist at equilibrium will be a measure of the strengths of the acids and bases. This is equivalent to saying that the dissociation constants measure strengths of acids and bases. 

The position of the equilibrium in these acid-base reactions wUl depend on the relative acidity of the carbon acid and of the species BH (or, conversely, the basic strengths of and the carbanion). Some approximate pK values for typical carbon acids and the conjugate acids of some species commonly employed as bases are shown in Table 1.1. The numerical values recorded are approximate, since there is no method of accurately establishing absolute acidity in a single solvent medium for... 

As we have seen in several examples in this chapter, HCN acts as an acid in aqueous solutions. We introduced a few fundamental concepts of acids and bases in Chapter 3, but the context of equilibrium allows us to explore them further. Recall that we distinguished between strong acids (or bases), which dissociate completely in solution, and weak acids (or bases), which dissociate only partially. At this point in our study of chemistry, we should realize that this partial dissociation of weak electrolytes was an example of a system reaching equilibrium. So we can use equilibrium constants to characterize the relative strengths of weak acids or bases. One common way to do this is to use the pH scale, which we will define in this section. 

Identify the two acids in the equilibrium and their relative strengths and the two bases and their relative strengths.The position of the equilibrium lies toward the weaker acid and the weaker base. 

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