Chelation equilibrium, calculation

Equilibrium calculations indicate that the Gd3+ chelates used as contrast agents in MRI may partly dissociate in the body fluids. However, the results of the kinetic studies point to the proton, Cu2+ and Zn2+ assisted dissociation of Gd3+ complexes being slow at pH > 5. The excretion of these contrast agents from the body is relatively fast so that the system is far from the equilibrium and the extent of in vivo dissociation is very low. 

We have seen how the pH and the presence of several complexing agents can be taken into account in equilibrium calculations. If, as happens often, we can identify one reaction in the array of reactions as the principal reaction, then all the others can be properly be called side reactions, and treated in a convenient manner. For example, in complexometric titrations of metal ions with EDTA or some other polydentate chelating titrant, the presence of auxiliary ligands like NHj, citrate anion, etc., can best be accounted for by the use of the conditional constant, first introduced by Schwarzenbach and widely applied by Ringbom. 

The product is equal to the equilibrium constant X for the reaction shown in equation 30. It is generally considered that a salt is soluble if > 1. Thus sequestration or solubilization of moderate amounts of metal ion usually becomes practical as X. approaches or exceeds one. For smaller values of X the cost of the requited amount of chelating agent may be prohibitive. However, the dilution effect may allow economical sequestration, or solubilization of small amounts of deposits, at X values considerably less than one. In practical appHcations, calculations based on concentration equihbrium constants can be used as a guide for experimental studies that are usually necessary to determine the actual behavior of particular systems. 

A major problem for cells is posed by the relative insolubility of ferric hydroxide and other compounds from which iron must be extracted by the organism. A consequence is that iron is often taken up in a chelated form and is transferred from one organic ligand, often a protein, to another with little or no existence as free Fe3+ or Fe2+. As can be calculated from the estimated solubility product of Fe(OH)3 (Eq. 16-1),7 the equilibrium concentration of Fe3+ at pH 7 is only 10-17 M. 

All data were collected on toluene solutions (28°C) which were 5 x 10 1 Min iridium the results of a representative experiment are shown in Figure 1. It proved necessary to use the simple substrates in considerable (20X) excess in order to detect the oxidative addition reaction using the measured equilibrium constant, we then calculated the [Ir(III)]/[Ir(I)] (Rjjx) which would exist if the initial concentrations of substrate and 1 were both 5 x 10-l> M. Those calculated values are shown in Table I, together with our results for the chelating substrates. 

While there is uncertainty in and for Fe2+(NTA)(NO) and Fe(II)(EDTA)(NO), the rate constants for reaction (1), Littlejohn and Chang (9), and Teramoto et al. (17) indicate that is on the order of 107 M-l sec l. The equilibrium constants Keq = k /k-i are fairly well established (10) as being about 10 at 25 °C, indicating that k i is about 10 sec . From this approximate value of k i and the consumption rate equation for NO + S032 , we can calculate a consumption rate for Fe2+(L)(NO). However, the calculated rate is considerably faster than the observed rate. The calculated rate was obtained assuming that the nitric oxide released by the ferrous chelate reacts at the rate for hydrated nitric oxide. 

For multistep complexation reactions and for ligands that are themselves weak acids, extremely involved calculations are necessary for the evaluation of the equilibrium expression from the individual species involved in the competing equilibria. These normally have to be solved by a graphical method or by computer techniques.26,27 Discussion of these calculations at this point is beyond the scope of this book. However, those who are interested will find adequate discussions in the many books on coordination chemistry, chelate chemistry, and the study and evaluation of the stability constants of complex ions.20,21,28-30 The general approach is the same as outlined here namely, that a titration curve is performed in which the concentration or activity of the substituent species is monitored by potentiometric measurement. 

Major advances and problems in the field of synthesis, properties, structure and applications of polymers containing metallochelate units are discussed. Included are terminology, classification and nomenclature of these compounds as well as major approaches to calculating the equilibrium constants of chelation with polymeric ligands and chelate effect in metallopolymeric systems. Special attention is paid to the production and structural features of polymers containing metallochelate units. The most important applications of such polymers are classified... 

Main Approaches to Calculating the Equilibrium Constants in Metal Ion — Chelating Macroligand Systems... 

A combination of chelators for divalent cations is suitable to buffer the free Ca " concentration from 0.1 -100 (iM under experimental conditions. Added Mg " and ATP as well as the pH of the medium must be considered, because they alter the equilibrium between Ca and the chelators present. The free Ca and Mg " concentrations are calculated by a computer program and controlled by Ca and Mg " specific electrodes (Fohr et al., 1993). Each Ca " buffer is prepared separately from stock solutions, with a final check of pH, pCa, or pMg. If no Ca electrode is available, the calculated total amount of Ca (as CaCy and Mg (as Mg(CH3COO)2) must be added before the pH adjustment. Buffers can be stored at -20 °C but should be thawed only once, mainly because of decomposition of ATP. 

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