Bile acids, equilibrium precipitation

With each successive addition of HCl, an equivalent amount of bile acid is precipitated, which appears to act as a buffer producing a plateau in the curve, in which there is very little change in the pH. Toward the end of the bile salt titration, this plateau portion changes to a convex slope, and finally when the reaction shown in [Eq. (1)] is complete, the curve shows a second inflection point Z, which is the final equivalence point. Extrapolation of the plateau section of the curve horizontally from point X to point Y indicates the point at which precipitation of bile acid crystals would have occurred, had there been no supersaturation. Between points Y and X, therefore, the solution is supersaturated and is not in physical equilibrium. 

On the basis of surface and bulk interaction with water. Small [85] classified bile acids as insoluble amphiphiles and bile salts as soluble amphiphiles. On account of the undissociated carboxylic acid group, the aqueous solubility of bile acids is limited [35] in contrast, many bile salts have high aqueous solubilities as monomers [33] and, in addition, their aqueous solubilities are greatly enhanced by the formation of micelles [5,6]. Because many bile salts are weak electrolytes, their ionization and solubility properties are more complicated than those of simple inorganic or organic electrolytes [5,35]. For example, the p/Tj, values of bile acids in water vary markedly as functions of bile salt concentration and, because micelles formed by the A (anionic) species can solubilize the HA (acid) species [5,35], the equilibrium precipitation pH values of bile acids also vary as functions of bile salt concentration. Finally, certain bile salts are characterized by insolubility at ambient temperatures [2,5,6,86,87], only becoming soluble as micelles at elevated temperatures (the critical micellar temperature) [6]. 

A simple method for estimating the pH-solubility relationship of bile acids and salts is to carry out aqueous acidometric titration of a bile salt in water with a stronger mineral acid [5,35], Once the molarities of bile salt and mineral acid are known, the titration curves provide a direct measurement of equivalence, equilibrium and metastable pH values, the pH at which precipitation of the HA species occurs (pHpp,), an estimate of the solubilities of the HA species in water (if the system is < CMC) or in water plus micelles (if the system is > CMC) and a calculation of the apparent pK (pATg). The methods, results and interpretation of such titration curves for the common bile salts, titrated with HCl, have been described in detail elsewhere [5,6]. 

Fig. 21. Hypothetical titration curve for solutions of free bile salts or for glycine conjugates. IV — first equivalence point where titration of bile salt with hydrochloric acid commences, Y = last point where bile salt solution is in thermodynamic equilibrium as a single aqueous phase. T = Tyndall effect noted in this region of titration curve. X = point where precipitation of bile acid crystals commences. X — equilibrium pH at point of bile acid precipitation. Z = second equivalence point where titration of bile salt with hydrochloric acid is complete. TOT = total amount of acid required to complete the titration. HA = the amount of acid added from the first equivalence point (IV), to point V, which represents the maximum solubility of the bile acid (HA) in the bile salt solution (A"). For further explanation of the symbols, see text.

From the start of precipitation (X) onward, there are two phases present (liquid plus crystal), and since solutions between Y and X are unstable, the only portion of the curve that is in thermodynamic equilibrium for the single liquid phase is portion WY. Point 7, therefore, represents the maximum solubility of bile acid (HA) in the bile salt solution (A ) before supersaturation occurs. 

Fig. 26. Equilibrium pH levels at point of precipitation of cholic acid (O) and ratios of the number of molecules of bile salt necessary to solubilize one molecule of bile acid (J) from titration of varying mixtures of 1 % solutions of NaC and NaTC. The broken circle was taken from curve 9 in Fig. 25 and does not represent a true equilibrium pH.

Fig. 27. Hypothetical titration curve for a mixture containing two bile acids (cholic and glycocholic acid), both of which precipitate from solutions of their sodium salts (solid line). The broken lines represent the titration curve for the free bile salt alone, NaC (above) and the pure glycine conjugate alone, NaGC (below). W—Z = amount of hydrochloric acid required to titrate the NaC in the mixture. Z—R = amount of hydrochloric acid required to titrate the NaGC in the mixture. X — equilibrium pH at the point of precipitation of cholic acid. Q = equilibrium pH at the point of precipitation of glycocholic acid.

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