Stability constants potentiometric data

The sorption data of Cd2+ and Pb2+ by B. subtilis and E. coli were well described by a one-site complexation model (r2 > 0.9) with Cd2+ showing somewhat lower sorption affinities than Pb2+ (Kulczycki et al. 2002). A two-site sorption model yielded an improved fit but only for the E. coli data. The stability constants for the high- and low-affinity sorption sites differed by several orders of magnitude. The total metal sorption capacity of E. coli increased, and moved closer to the value of B. subtilis when the presence of low-affinity sorption sites was allowed. Ngwenya et al. (2003) used potentiometric titrations to assess the different types of sites present... 

Borrok et al. (2004a) used potentiometric titration to measure Cd sorption by different bacterial consortia, and a surface complexation approach to determine thermodynamic stability constants. When the data were modeled by adopting a single set of stability constants, a similar sorption behavior was shown by a wide range of bacterial species. Further, current models that rely on pure strains of laboratory-cultivated bacterial species appear to overestimate the extent of metal biosorption in natural systems. 

Complex stability constants are often determined by pH-potentiometric titration of the ligand in the presence and absence of the metal ion (129). This method works well when equilibrium is reached rapidly (within few minutes), which is usually the case for linear ligands. For macrocyclic compounds, such as DOTA and its derivatives, complex formation is slow, especially at pH-s where the formation is not yet complete, therefore a batch method is used instead of direct titration (130,131). A few representative examples of stability constant data mainly collected from Ref. (132), on MRI relevant Gdm complexes are presented in Table IV. 

Numerous investigations have shown the existence of the heptamolybdate, [Mo7024]6 , and octamolybdate, [Mo8026]4, ions in aqueous solution. Potentiometric measurements with computer treatment of the data proved to be one of the best methods to obtain information about these equilibria. Stability constants are calculated for all species in a particular reaction model, which is supposed to give the best fit between calculated and experimental points. In the calculations the species are identified in terms of their stoichiometric coefficients as described by the following general equation for the various equilibria... 

Not mentioned in Table 2 (and often not in the original papers ) is the optical form (chirality) of the amino acids used. All the amino acids, except for glycine (R = H), contain an asymmetric carbon atom (the C atom). In the majority of cases the optical form used, whether l, d or racemic dl, makes little difference to the stability constants, but there are some notable exceptions (vide infra). Examination of the data in Table 2 reveals (i) that the order of stability constants for the divalent transition metal ions follows the Irving-Williams series (ii) that for the divalent transition metal ions, with excess amino acid present at neutral pH, the predominant spedes is the neutral chelated M(aa)2 complex (iii) that the species formed reflect the stereochemical preferences of the metal ions, e.g. for Cu 1 a 2 1 complex readily forms but not a 3 1 ligand metal complex (see Volume 5, Chapter 53). Confirmation of the species proposed from analysis of potentiometric data and information on the mode of bonding in solution has involved the use of an impressive array of spectroscopic techniques, e.g. UV/visible, IR, ESR, NMR, CD and MCD (magnetic circular dichroism). 

NMR titrations (of anion into ligand at fixed pH) and pH-potentiometric titrations (of pH at fixed anion ligand ratios) provide comparable values of the stability constants for binding of mononegative oxoanions by protonated R3Bm, R3F, and R3P hosts [15,20,21] Table 2. The weak complexation at hexaprotonated levels for tetrahedral monoanionic oxoanions makes it difficult to obtain reliable data for protonation levels below 5. This has however been achieved for nitrate with the cleft binding host R3P as well as for Re O4 with the most basic cryptand R3Bm. 

The results of potentiometric titration, the number of edge sites, and intrinsic stability constants of the protonation and deprotonation reactions of calcium-, copper-, zinc-, manganese(II)-montmorillonites, and KSF montmorillonite are shown in Table 2.4. As a comparison, some similar data for other montmoril-lonites are also listed. 

A similar result was obtained for the corresponding silver(T) trihelicates using both spectrophotometric and potentiometric (silver electrode) data. Thus, the stability constant [corresponding to reaction (6.2)] is greater than [for reaction (6.1)]. In fact, for positive co-operativity, it is sufficient that > (AT /3). 

Stepwise stability constants for complex formation between Al111 of In111 and I-asparagine or l -glutamine have been evaluated from potentiometric data.388... 

Surface charging of ZnS was interpreted in terms of the following surface species =SZn, =SH2, =SZnOH", =ZnS and —ZnSH" [79]. Stability constants of the surface species were fitted to potentiometric titration data combined with the readings of the S " ion selective electrode. 

One of the major difficulties in evaluating Zr complexation constants is the unavoidable coexistence with hydroxo species over all pH regions of interest. Therefore, any determination of stability constants for complex formation with ligands other than OH critically depends on the quality and precision of the stability constants assigned to the hydrolysis. This is particularly true in the carbonate system where OH and CO3 concentrations will co-vary with pH. In the course of the review, it became evident that all Zr-carbonate constant determinations found in the literature relied on a hydrolysis model that differs from that selected in this review. Consequently, all constant determinations had to be re-evaluated. Due to the limited information provided by most references (missing raw data, insufficient declaration of experimental conditions), or because of the inherent unsuitability of the data, a meaningful reinterpretation was possible only in a few cases (mainly for potentiometric titrations). Table V-37 is a compilation of the Zr-carbonate complexation constants reported in the... 

The interpretation of the titration data given by the authors starts from the premise that Zr(OH)2 is the main hydrolysis species. Moreover, as in [99VEY], the implicit assumption is made that all Zr is bound to carbonate complexes, i.e. a very high formation constant is assumed a priori. Therefore, the stability constant ( "4 = 8.0 x 10 °) derived by [80MAL/CHU] for the formation reaction Zr(C03)j" + CO " = Zr(C03)4, had to be rejected. We re-interpreted the potentiometric data on the base of full equilibrium calculations and using the Zr hydroxo, chloride and sulphate stability constants selected in this review (see below). 

In order to obtain more detailed information from the potentiometric data, we calculated the initial pH-values (j.e. before addition of any titrant) obtained assuming formation of specific complexes. The calculations were carried out for four titrations (2 and 3 in Fig. 1, 1 and 2 in Fig. 2 of the paper) with the GEMS-PSI code, using the extended Debye-Hiickel approximation. In each computation (see Table A-26) we assumed the formation of a single complex with either two values of the stability constant. The first value was adjusted to achieve about 20% binding of the available Zr to the selected complex. The second value was selected sufficiently high to ensure that... 

From titration of solutions b) and c), the author deduced the existence of the species Zr2(OH)) and Zr(OH)4 and their corresponding stability constants from the potentiometric equivalence points corresponding to the neutralisation of 1.5 and 2 protons per Zr atom. However, the stoichiometry of these species can only be deduced from these data if the speciation of Zr in the initial solution is known. Uncertainties in the speciation of the initial solution result in equivalent uncertainties in the stoichiometries of the observed hydrolysis reactions. For example, in case of tetramer formation, titration end point data could equally well be explained by formation of and Zr4(OH)i6(aq). 

In the second step, with the hydrolysis constants and the specific interaction parameter for ZrOH" and for Zr3(OH) fixed to the values optimised as detailed above, the equilibrium constants and interaction parameter for all other species in the overall hydrolysis model were obtained by a global fit of the potentiometric, solubility, solvent extraction and ion exchange data mentioned above. The fit was extended to the determination of equilibrium constants for heterogeneous reactions ion exchange constants, solubility constants and liquid/liquid distribution coefficients. The fit was based on a preselection of the stoichiometries of dominant species which included invariably the species Zr(OH)4(aq), Zr ) ), Zr (OH)Jj and Zr4(OH)i6(aq) and various other mono-, di-, tri- and tetravalent species to improve the fit. The potential formation of chloride complexes of Zr was considered for chloride containing solutions, using the stability constants determined in Section V-4. If all fitted results were found insensitive to the equilibrium constants of a given species, the respective species was removed from the list of species. 

The thermodynamic stability constant of the NiSCN" complex was determined by spec-trophotometric t = (25 1)°C) and potentiometric (/ = 35°C) methods. The Davies equation was used to calculate the activity coefficients, and this is not compatible with the SIT. Therefore, the potentiometric data were re-evaluated, using the SIT approach. The resulting log o Pi - Im plot is rather scattered and the ionic strength range was relatively narrow (see Figure A-8). Therefore the reported data were not considered further in this review. 

The equilibria in the nickel(II)-imidazole-chloride ternary system have been investigated by means of pH-metric and spectrophotometric titrations in 3 M Na(C104, Cl) solution. The latter method was used to determine the stability constant of the binary NiCr complex in the nickel(II)-chloride system. The concentration range studied was [NP" ] = 0.096 to 0.300 M and [Cl ] = 1.25 to 3.00 M. The log, yff value obtained (- (0.47 0.10)) is in good agreement with that determined from potentiometric data in... 

The following entry defines the commonly used stability constants (stepwise, overall, conditional, association, dissociation, and pK) and relates the values to a rigorous thermodynamic definition of equilibrium constants. In addition, the article briefly outlines experimental techniques (potentiometric titration, spectroscopic methods involving ultraviolet/visible, infrared, Raman, fluorescence. and nuclear magnetic resonance spectroscopy), together with the numerical methods and computer programs that can be used to derive stability constants from such experimental data. 

The computer program Hyperquad may be used to determine stability constants from potentiometric data. This program employs the general procedure outlined above with the following specifics for implementation. 

The incorrect nature of this concept has been convincingly demonstrated by the equilibrium studies of Ahrland [Ah 76] in aqueous and in dimethyl sulphoxide solutions. In order to illustrate the effect of the solvation of the anions on the stabilities of their metal complexes, Ahrland compared the stepwise stability constants, determined potentiometrically in water and in dimethyl sulphoxide, of the zinc and cadmium complexes of halide ions, which form hydrogen bonds with strengths decreasing in the sequence Cl >Br >I . As can be seen from the data relating to the cadmium complexes in Table 6.1, the absolute values of the equilibrium constants in dimethyl sulphoxide solution were considerably higher than those in water. In addition, the-stability sequences for the various halide... 

It is possible to apply a simultaneous regression estimation of stoichiometric coefficients and stability constants (i,e, ESI) [30,64], in which both stoichiometric coefficients and extraction constants are given as adjustable parameters and the program searches for the best model changing also stoichiometric indices as real numbers. This approach introduced by Havel et al, [30,64] has been implemented in the program POLET [65] and used for the treatment of potentiometric [66], spec-trophotometric [67], and kinetic data [68], The method has also been applied to... 

To consider the reasons responsible for such behavior, the distribution of species at the electrode surface should be analyzed. Procedures appropriate for this purpose are described in Section 3.3. They require the stability constants of metal complexes and those of protonated ligands to be employed. According to the potentiometric data [92] that were improved in Ref [93], at least two Sn(II) citrate complexes, namely SnL and SnLH", should be taken into account at 3.5 < pH5.5 with log equal to 15.35 and 19.5, respectively. Stability constants of protonated ligand species are given in Table 8.4. Simulations have shown that the surface concentrations of Sn(II) citrate complexes vary differently with cathodic current density (Figure 8.28). Despite the fact that the total surface concentration of Sn(II) decreases with cathodic i, the surface concentration of... 

Table 3 Negative logarithms of the acidity constants (eq. 3) of some monoprotonated purine nucleobase (NB) derivatives and logarithms of the stability constants (eq. 2) of the corresponding Cd(NB) and Cd(NB - H) complexes as determined by potentiometric pH titration in aqueous solution at 25°C and / = 0.5 M (entries 1 - 7) or 0.1 M (entries 8 - 10) (NaNOs), together with data for 1-methylimidazole (IMIm) or pyridine (Py) derivatives (see Figure 5) .

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