Figure 6.4 Solubility-pH profile of a weak acid, with salt precipitation taken into account. [Avdeef, A., Curr. Topics Med. Chem., 1, 277-351 (2001). Reproduced with permission from Bentham Science Publishers, Ltd.]

increases. Although precipitation of salts is not covered in detail in this chapter, it is nevertheless worthwhile to consider its formation in this limiting sense. As the pH change raises the solubility, at some value of pH the solubility product of the salt will be reached, causing the shape of the solubility-pH curve to change from that in Fig. 6.1a to that in Fig. 6.4, an example of a weak acid exhibiting salt precipitation.

As a new rule of thumb [473], in 0.15 M NaCl (or KCl) solutions titrated with NaOH (or KOH), acids start to precipitate as salts above log (S/S0) ~ 4 and bases above log (S/So) ~ 3. It is exactly analogous to the diff 3-4 rule; let us call the solubility equivalent the ''sdiff 3-4'' rule [473]. Consider the case of the monopro-tic acid HA, which forms the sodium salt (in saline solutions) when the solubility product Ksp is exceeded. In additions to Eqs. (3.1) and (6.1), one needs to add the following reaction/equation to treat the case:

Na+ A"(s) ^ Na+ + A" KSp = = [Na+][A-] (6.8)

Effective solubility is still defined by Eq. (6.3). However, Eq. (6.3) is now solved under three limiting conditions with reference to a special pH value:

1. If the solution pH is below the conditions leading to salt formation, the solubility-pH curve has the shape described by Eq. (6.4) (curve in Fig. 6.1a).

2. If pH is above the characteristic value where salt starts to form (given high enough a sample concentration), Eq. (6.3) is solved differently. Under this circumstance, [A ] becomes the constant term and [HA] becomes variable.

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