Stability constants concentration calculations with

Stability constants are calculated from the concentrations of the species present in equilibrium mixtures containing the metal ion and the ligand in a wide range of proportions. Activity coefficients are kept constant by appropriate additions of a salt, usually sodium perchlorate, whose ions do not compete with those of the cation and ligand. Concentrations at different ionic strengths are extrapolated to zero ionic strength. It may be necessary to find the number of water molecules displaced at each step the total of these is not necessarily the same as the co-ordination number of the cation in the solid compound. Particularly in a polar solvent such as water, the ligands may not displace all the solvent molecules. 

This stability constant deals only with the species of EDTA where all carboxylic acid groups are deprotonated. The pKa values for these groups are pKai 2.0, pKa2 2.67, pKaj 6.16, and pKa410.3, so the fully deprotonated EDTA dominates only in very alkaline solutions with pH above 12. In solutions of lower pH one has to calculate [Y ] at the actual pH and use this concentration in calculating a stability constant. Similarly, if the metal M participates its other complexes, one has to calculate the [M"+] actually available for EDTA complexation. The stability constants incorporating the effects of pH and other complex reaction are called apparent stability constants. The absolute stability constant for a number of cations has been determined and is tabulated in Table 6.6. 

Variamine blue (C.I. 37255). The end point in an EDTA titration may sometimes be detected by changes in redox potential, and hence by the use of appropriate redox indicators. An excellent example is variamine blue (4-methoxy-4 -aminodiphenylamine), which may be employed in the complexometric titration of iron(III). When a mixture of iron(II) and (III) is titrated with EDTA the latter disappears first. As soon as an amount of the complexing agent equivalent to the concentration of iron(III) has been added, pFe(III) increases abruptly and consequently there is a sudden decrease in the redox potential (compare Section 2.33) the end point can therefore be detected either potentiometrically or with a redox indicator (10.91). The stability constant of the iron(III) complex FeY- (EDTA = Na2H2Y) is about 1025 and that of the iron(II) complex FeY2 - is 1014 approximate calculations show that the change of redox potential is about 600 millivolts at pH = 2 and that this will be almost independent of the concentration of iron(II) present. The jump in redox potential will also be obtained if no iron(II) salt is actually added, since the extremely minute amount of iron(II) necessary is always present in any pure iron(III) salt. 

Speciation of Pb(II) in Glatt river. The concentrations given for CO2, Pb(II), Cu(II) and [Ca2+] as well as for the pollutants EDTA and NTA are representative of concentrations encountered in this river, The speciation is calculated from the surface complex formation constants determined with the particles of the river and the stability constants of the hydroxo-, carbonate-, NTA- and EDTA-complexes.The presence of [Ca2+] and [Cu2+] is considered. 

The examples in section 3.1.2 of calculations using stability constants involve concentrations of M, L, and ML . Rigorously, a stability constant, as any thermodynamic equilibrium constant should be defined in terms of standard state conditions (see section 2.4). When the system has the properties of the standard state conditions, the concentrations of the different species are equal to their activities. However, the standard state conditions relate to the ideal states described in Chapter 2, which can almost never be realized experimentally for solutions of electrolytes, particularly with water as the solvent. For any conditions other than those of the standard state, the activities and concentrations are related by the activity coefficients as described in Chapter 2, and especially... 

Figure 1. Plot of v/V ax versus the millimolar concentration of total substrate for a model enzyme displaying Michaelis-Menten kinetics with respect to its substrate MA (i.e., metal ion M complexed to otherwise inactive ligand A). The concentrations of free A and MA were calculated assuming a stability constant of 10,000 M k The Michaelis constant for MA and the inhibition constant for free A acting as a competitive inhibitor were both assumed to be 0.5 mM. The ratio v/Vmax was calculated from the Michaelis-Menten equation, taking into account the action of a competitive inhibitor (when present). The upper curve represents the case where the substrate is both A and MA. The middle curve deals with the case where MA is the substrate and where A is not inhibitory. The bottom curve describes the case where MA is the substrate and where A is inhibitory. In this example, [Mfotai = [Afotai at each concentration of A plotted on the abscissa. Note that the bottom two curves are reminiscent of allosteric enzymes, but this false cooperativity arises from changes in the fraction of total "substrate A" that has metal ion bound. For a real example of how brain hexokinase cooperatively was debunked, consult D. L. Purich H. J. Fromm (1972) Biochem. J. 130, 63.

Stability constants, from which AG° values are calculated, provide a direct measure of the extent of complexing in solution, and these values have been used to determine cation selectivity by macrocyclic compounds. Several of the methods commonly used to determine log K values cannot be used with many of these systems. Thus, procedures based on change in hydrogen ion concentration (pH titration, hydrogen electrode, etc.) cannot be used in those cases where the ligand is uncharged and its concentration is not pH dependent. Spectral methods generally have not been used because of the usual lack of favorable absorption characteristics by the compounds, cations or cation-complexes in the cases studied. 

An excess of OH (the common ion) should shift the reaction to the right, i.e., to more complete precipitation of the Zn(OH)2. This effect is a general one, but the conclusions are not always vahd the example (deliberately) given here is one where it is not valid. The reason is that OH can form a complex with Zn (Zn(OH)4 — the zincate ion), thus removing free Zn from solution and reducing the degree of precipitation. For a sufficiently high concentration of OH , which can be calculated from the stability constant of the zinc-hydroxide (zincate) complex, the Zn(OH)2 will completely redissolve. 

Menger et al. synthesized a Ci4H29-attached copper(II) complex 3 that possessed a remarkable catalytic activity in the hydrolysis of diphenyl 4-nitrophenyl phosphate (DNP) and the nerve gas Soman (see Scheme 2) [21], When 3 was used in great excess (ca. 1.5 mM, which is more than the critical micelle concentration of 0.18 mM), the hydrolysis of DNP (0.04 mM) was more than 200 times faster than with an equivalent concentration of the nonmicellar homo-logue, the Cu2+-tetramethylethylenediamine complex 9, at 25°C and pH 6 (Scheme 4). The DNP half-life is calculated to be 17 sec with excess 1.5 mM 3 at 25°C and pH 6. The possible reasons for the rate acceleration with 3 were the enhanced electrophilicity of the micellized copper(II) ion or the acidity of the Cu2+-bound water and an intramolecular type of reaction due to the micellar formation. On the basis of the pH(6-8.3)-insensitive rates, Cu2+-OH species 3b (generated with pK3 < 6) was postulated to be an active catalytic species. In this study, the stability constants for 3 and 9 and the thermodynamic pvalue of the Cu2+-bound water for 3a —> 3b + H+ were not measured, probably because of complexity and/or instability of the metal compounds. Therefore, the question remains as to whether or not 3b is the only active species in the reaction solution. Despite the lack of a detailed reaction mechanism, 3 seems to be the best detoxifying reagent documented in the literature. 

The thermodynamic stability constant of silver sulfadiazine in water has been measured by Boelema et al. (20). They made use of a microcomputer-controlled titrator and measured simultaneously the pH value with a Radiometer G2040C glass electrode and the silver ion concentration with an Orion research 941600 ion-selective silver sulfide electrode. The measured stability constant and standard deviation were logK=3.62i 0.05 n = 9, temperature 25°C, ionic strength 0.1 (sodium nitrate). The calculated conditional stability constant log K at pH 7.4 was 3.57. 

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