Metal ion/buffer interactions. Stability of alkali and alkaline earth ion complexes with triethanolamine (tea), 2-amino-2(hydroxymethyl)-1,3-propanediol (tris)and 2-[bis(2-hydroxyethyl)-amino] 2(hydroxymethyl)-1,3-propanediol (Bistris) in aqueous and mixed solvents☆

Add time:08/20/2019    Source:sciencedirect.com

The acidity constants of the protonated buffers given in the title, i.e. of H(Tea)+, H(tris)+ and H(Bistris)+, have been measured at 25 °C in water, 50% aqueous dioxane or methanol, and in 75 or 90% dimethylsulfoxide (Dmso) with tetramethylammonium nitrate as background electrolyte. The interaction of Tea, Tris, or Bistris (L) with the alkali or alkaline earth ions (Mn+ was studied by potentiometric pH titrations in the same solvents and the stability constants of the MLn+ complexes were determined. The stability constants, log KMML, of the Na+ complexes with the several buffer-ligands in the given solvents vary from −1.05 [Na(Tea)+ in H2O; I = 1.0] to 0.54 log units [Na(Bistris)+ in 90% Dmso; I = 0.25]; the corresponding values for the Mg2+ complexes range from 0.24 [Mg(Tea)2+ in H2O; I = 1.0] to 0.91 log units [Mg(Bistris)2+ in 90% Dmso; I = 0.25]. Unexpectedly, Ca(Bistris)2+ is the most stable among the alkaline earth ion complexes in aqueous solution (log KCaCa(Bistris) = 2.25; the corresponding values for the Mg2+, Sr2+ and Ba2+ complexes are 0.34, 1.44 and 0.85, respectively; I = 1.0), while in 90% Dmso Sr(Bistris)2+ is most stable (log KSrSr(Bistris) = 1.87; the corresponding values for the Mg2+, Ca2+ and Ba2+ complexes are 0.91, 1.64 and 1.14, respectively; I = 0.25). A similar, but less pronounced pattern is observed for the M(Tea)n+ complexes. obviously, the stabilities of the alkaline earth ion complexes with Bistris and Tea follow neither the order of the ionic radii nor that of the hydrated radii of the cations. In contrast, in all solvents the stability of the alkali ion complexes increases with decreasing ionic radii; this being also true for the alkaline earth iron complexes of Tris in aqueous solution. The possible reasons for these observations, the structures of the complexes in solution, and some biological implications are discussed. Calculations of the extent of complex formation show that in the physiological pH range the concentration of certain complexes may be quite pronounced; hence reservations should be exercised in employing these buffers in systems which also contain metal ions.

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