Acid-base equilibria basic anions

To calculate the pH of a salt solution, we can use the equilibrium table procedure described in Toolboxes 10.1 and 10.2—an acidic cation is treated as a weak acid and a basic anion as a weak base. However, often we must first calculate the Ka or Kh for the acidic or basic ion. Examples 10.10 and 10.11 illustrate the procedure. 

If the conjugate-base-dissociation mechanism is the correct one, the lower substitution rates observed with nitrite, acetate, and azide may be considered to arise because these weakly basic anions are less effective than methoxide in generating conjugate base CB in the preliminary acid base equilibrium. 

The alkoxide doesn t have to be made first, though, because alcohols dissolved in basic solution are at least partly deprotonated to give alkoxide anions. How much alkoxide is present depends on the pH of the solution and therefore the pKa of the base (Chapter 8), but even a tiny amount is acceptable because once this has added it will be replaced by more alkoxide in acid-base equilibrium with the alcohol. In this example, allyl alcohol adds to pent-2-enal, catalysed by sodium-hydroxide in the presence of a buffer. 

By comparing the approximate pK values of the conjugate acids of the basic catalysts with those of the carbon acid of interest, it is possible to estimate the position of the acid-base equilibrium for a given reactant-base combination. If we consider the case of a simple alkyl ketone, for example, it can be seen that hydroxide ion, and primary alkoxide ions, will convert only a fraction of such a ketone to its anion ... 

The values of [HA] and [A ] in this expression are the equilibrium concentrations of acid and base in the solution, not the concentrations added initially. However, a weak acid HA typically loses only a tiny fraction of its protons, and so [HA] is negligibly different from the concentration of the acid used to prepare the buffer, [HA]initia. Likewise, only a tiny fraction of the weakly basic anions A- accept protons, and so [A-] is negligibly different from the initial concentration of the base used to prepare the buffer. With the approximations A ] [base]initia and [HA] [acid]initia, we obtain the Henderson-Hasselbalch equation ... 

An aqueous solution of a soluble salt contains cations and anions. These ions often have acid-base properties. Anions that are conjugate bases of weak acids make a solution basic. For example, sodium fluoride dissolves in water to give Na, F, and H2 O as major species. The fluoride anion is the conjugate base of the weak acid HF. This anion establishes a proton transfer equilibrium with water ... 

It is more common to hydrolyze esters under basic conditions because the equilibrium is favorable. The mechanism for this process, called saponification, is presented in Figure 19.4. The production of the conjugate base of the carboxylic acid, the car-boxylate anion, which is at the bottom of the reactivity scale, drives the equilibrium in the desired direction. To isolate the carboxylic acid the solution must be acidified after the hydrolysis is complete. Some examples are provided in the following equations. We saw another example of this hydrolysis reaction in Chapter 10, where it was... 

In case of a weak monoprotic acid-base analyte, two different species are retained the protonated acid form (HA, uncharged or cationic) and the deprotonated conjugate basic form (A, uncharged or anionic). In other words, if HA is uncharged, A is anionic if HA is charged, A is uncharged. The ionization of the analyte is described by the following equilibrium ... 

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. 

The ability to separate strong from weak acids depends on the acidity constants of the acids and the basicity constants of the bases as follows. In the first equation consider the ionization of benzoic acid, which has an equilibrium constant, K , of 6.8 x 10 The conversion of benzoic acid to the benzoate anion in Eq. 4 is governed by the equilibrium constant, K (Eq. 5), obtained by combining the third and fourth equations. 

Another nontraditional weak acid is the hydrogen on the a-carbon of a ketone such as 2-butanone. When 2-butanone (A) reacts with sodium ethoxide (B), the conjugate base is enolate anion 8 (CB), and ethanol is the conjugate acid (CA). Ethanol is a stronger acid than the ketone and can reprotonate the basic enolate. This shifts the equilibrium back to the left, and the net result is a solution that contains enolate 8, ethanol, unreacted 2-butanone, and NaOEt (sec. 9.2.E). It is important to note that HA does not function as an acid unless a sufficiently strong base (B) is present to remove and accept the acidic proton. If the base under consideration... 

From this follows that the equilibrium concentrations of B and A should be equal. This condition is met in the crossing point of the acid line of the acid corresponding to the anion and the base line of the cation, i.e., in Pj in Figs. 56, 57, and 58. The pH-coordinate of this point is the pH of the solutions. This graphical way of finding the pH is equivalent to the calculations derived for ampholytes which are weakly acidic and weakly basic (see Sect. 4.2.2.2) ... 

In general, the solubility of a compound containing a basic anion (that is, the anion of a weak acid) increases as the solution becomes more acidic. As we have seen, the solubility of Mg(OH)2 greatly increases as the acidity of the solution increases. The solubility of Pbp2 increases as the solution becomes more acidic, too, because F is a weak base (it is the conjugate base of the weak acid HF). As a result, the solubility equilibrium of Pbp2 is shifted to the right as the concentration of F is reduced by protonation to form HF. Thus, the solution process can be understood in terms of two consecutive reactions ... 

Solvent effects on acid-base equilibria are naturally most marked when the solvent itself enters into the equilibrium, as is the case for the conventional definition of acid strength by means of the equilibrium A-fSH B4-SH2 (where SH is the solvent). The existence of such an equilibrium implies that the solvent has some basic properties. Similarly, the occurrence of the reaction B + SH A-f S (where S is the anion derived by abstracting a proton from the solvent) implies that the solvent is acidic. The most important factor determining qualitative behaviour in a wide range of solvents is the acidic or basic nature of the solvent, as determined by its chemical nature. In a preliminary classification we can neglect other factors, notably the effect of dielectric constant on the association of ions or the forces between them. 

We recall that all sodium salts are soluble and that all soluble ionic salts are completely dissociated in H2O. We recognize that both NaCHsCOO and NaCN are salts of strong bases (which provide the cation) and weak acids (which provide the anion). The anions in such salts hydrolyze to give basic solutions. In the preceding text we determined that for CH3COO = 5.6 X 10" , and in Example 18-19 we determined that K, for CN" = 2.5 X 10 . As we have done before, we first write the appropriate chemical equation and equilibrium constant expression. Then we complete the reaction siunmary, substitute the algebraic representations of equilibriiun concentrations into the equilibrium constant expression, and solve for the desired concentration(s). 

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