Bile acids, titration curves

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. 7. Potentiometric titration curves of free and glycine (G)-conjugated ursodeoxycholic (UDC) and chenodeoxycholic (CDC) acids. Inset in D represents typical titration curve for free and glycine-conjugated bile acids with notations as explained in the text. The arrows on the GCDC curves indicate the points of bile acid precipitation. (From ref. 35 with permission.)...

The effect of pH on the physical state of bile acids and their sodium salts is perhaps best explained by examining a titration curve of typical bile salt obtained by titrating an alkaline bile salt solution with hydrochloric acid. 

An ideal titration curve of a free bile salt at a concentration above its CMC is given in Fig. 21. The initial part of the curve represents the titration of the small excess NaOH. At the inflection point W, the convex slope of the curve suddenly becomes concave, indicating the first equivalence point (i.e., the point at which the reaction given in [Eq. (1)] starts). The curve then flattens slowly until at point X precipitation suddenly occurs and, due to the trapping of the H ions in the crystalline precipitate, the pH rises sharply to X without the further addition of acid. This precipitation is often heralded by the appearance of a Tyndall effect at about T. 

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. 

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.

Table V also gives the pH of precipitation and the number of bile acid anions needed to solubilize one molecule of bile acid (A /HA). The values in Table V were determined from the titration curves of the bile salts in water at 37°C(128).

Examples of the titration curves for a pure bile salt (NaC) and pure conjugated bile salt solutions (NaGC and NaTC) are shown in Fig. 24. The curves for sodium cholate and sodium glycocholate are similar to the ideal curve (Fig. 21), but since sodium taurocholate is the salt of a strong acid and remains completely ionized, the solution remains clear throughout the titration and therefore neither the supersaturation nor the precipitation characteristic of the free and glycine-conjugated bile salts takes place. 

Unlike the mixtures of NaC and NaTC in which the conjugate always remains in solution, both cholic and glycocholic acid form crystalline precipitates (Fig. 24) and therefore solutions containing mixtures of sodium salts of these bile acids exhibit a different type of titration curve (Fig. 27). The titration curves for pure NaC and NaGC are shown in broken lines together with the curve for a mixture containing 50% each of free and conjugated bile salt. 

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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