EDTA complexes Conditional constants

In equation (q) only the fully ionised form of EDTA, i.e. the ion Y4 , has been taken into account, but at low pH values the species HY3, H2Y2, H3 Y and even undissociated H4Y may well be present in other words, only a part of the EDTA uncombined with metal may be present as Y4. Further, in equation (q) the metal ion M"+ is assumed to be uncomplexed, i.e. in aqueous solution it is simply present as the hydrated ion. If, however, the solution also contains substances other than EDTA which can complex with the metal ion, then the whole of this ion uncombined with EDTA may no longer be present as the simple hydrated ion. Thus, in practice, the stability of metal-EDTA complexes may be altered (a) by variation in pH and (b) by the presence of other complexing agents. The stability constant of the EDTA complex will then be different from the value recorded for a specified pH in pure aqueous solution the value recorded for the new conditions is termed the apparent or conditional stability constant. It is clearly necessary to examine the effect of these two factors in some detail. 

In the case of a solution such as electroless Ni-P, Ni2+ is usually complexed by citrate, and the stability constants are ca. 104 and 2 x 108 (overall value) for the ML and ML2 complexes [67], Thus, pM will change relatively slowly with pH. On the other hand, the stability constant for the Pd-EDTA complex system (ML type only) is reported to be 1024 [67], i.e. Pd2+ is strongly complexed by EDTA. The Pd2+ pM value changes drastically, in a practical electroless deposition sense, over a rather narrow pH range. Consequently, in the case of an electroless Pd solution with EDTA as complexant, the solution may go from a condition of near spontaneous plating out to one where deposition virtually ceases. 

Calculate the conditional formation constant for the Zn-EDTA complex at pH 9.0 in a solution... 

The conditional formation constant allows us to look at EDTA complex formation as if the uncomplexed EDTA were all in one form ... 

You can see from the example that a metal-EDTA complex becomes less stable at lower pH. For a titration reaction to be effective, it must go to completion (say, 99.9%), which means that the equilibrium constant is large—the analyte and titrant are essentially completely reacted at the equivalence point. Figure 12-9 shows how pH affects the titration of Ca2+ with EDTA. Below pH 8, the end point is not sharp enough to allow accurate determination. The conditional formation constant for CaY2" is just too small for complete reaction at low pH. 

With the conditional formation constant we can treat EDTA complex formation as if all the free EDTA were in one form. 

Equation 12-18 states that the effective (conditional) formation constant for an EDTA complex is the product of the formation constant, Kf. times the fraction of metal in the form M" + times the fraction of EDTA in the form Y4- K" = aM, + aY4 h t. Table 12-1 told us that aY4- increases with pH until it levels off at 1 near pH 11. 

In a direct titration, analyte is titrated with standard EDTA. The analyte is buffered to a pH at which the conditional formation constant for the metal-EDTA complex is large and the color of the free indicator is distinctly different from that of the metal-indicator complex. 

The equilibrium between Mg2+ and the edta anion L4 can be compared to that between NH3 and H+, because the corresponding equilibrium constants are very dose in magnitude. In the latter case, it is possible to titrate NH3 with a solution of a strong acid in order to determine quantitatively its total concentration. It is therefore quite evident that, based on the values of the equilibrium constants of Scheme 3, the quantitative determination of the Mg2+ using edta should be possible. Because the other cations form more stable complexes than Mg2+, the complexometric titration should be of wide application. Some caution is necessary concerning the pH value at which the determination is done, because the ligand can be protonated, with consequent decrease of its chelating power. However, in the case of copper(II), its edta complex is already completely formed at pH 3 and therefore a titration is possible under these conditions. 

Calculate the conditional formation constant for the Zn-EDTA complex at pH 9.0. in a solution 0.100 M with respect to ammonia. Can zinc be titrated quantitatively with EDTA solution at this pH ... 

FIGURE 11-3 Conditional formation constants as a function of pH for metal-EDTA complexes. The dotted curve, Zn -I- NH3, represents zinc in the presence of[NHa] + [NH4 ] = 1 M. Adapted from Ringbom. )... 

Figure 11-7 depicts the relation between the titration error and the product Cm STm-l -Figure 11-3 indicates that conditional formation constants of at least 10 ° are possible for most metal ions in a selected pH region. According to Figure 11-7 a titration error of 0.1% should be easily attainable for 0.01 M solutions of such metals. For a 0.1% relative excess of reagent in the titration of 0.01 M metal ion, [Y ] of the reagent then equals 10" M. A value of = 10 then corresponds to a 99.9% conversion of metal ion to EDTA complex since [MY]/Cm = 1000. 

Curve A in Figure 17-6 is a plot of data for the titration in Example 17-4. Curve B is the titration curve for a solution of magnesium ion under identical conditions. The formation constant for the EDTA complex of magnesium is smaller than that of the calcium complex, which results in a smaller change in the p-function in the equivalence-point region. 

The major side effect of chelation therapy, particularly with EDTA, is hypocalcaemia, a condition caused by too rapid administration of the chelator. The result is a rapid drop in the ionized calcium in the blood plasma that causes muscle and abdominal cramps, convulsions, and even death. The condition is usually controlled by infusion of calcium gluconate, or prevented when the metal to be removed complexes with a much higher stability constant with EDTA than that of the Ca-EDTA, by administration as the Ca-EDTA complex. 

Thus at pH = 4, the conditional stability constants of some metal-are calculated as follows -EDTA complexes... 

Equation (5.2-2) is a useful form for the EDTA complexes of metal ions such as magnesium and calcium that are rather strong bases, and therefore have little tendency to form hydroxides. However, many other metal ions that can be titrated with EDTA will often form hydroxy complexes, and these are usually titrated in the presence of complexing agents that keep hydroxide formation at bay. In that case, the expression for the conditional formation constant must take such complex formation of the metal ion into account as well, and then reads Kf = KfaY0 aM0, where aM0 is the fraction of... 

Under stoichiometric conditions, fluorexon and its derivatives form 1 1 complexes with Ln ions. However when the ratio Ln Fx is increased, complexes with other stoichiometries are observed, the exact nature of which has not been determined. On the other hand, luminescence data of solutions with a ratio Yb Fx < 1 clearly indicate the presence of only one luminescent species, the 1 1 complex. Monoexponential luminescence decays are observed corresponding to a lifetime of 1.9 ps, whereas multi-exponential decays are measured when the Yb Fx ratio is increased. Further proof of the existence of 1 1 complexes has been brought by mass spectrometry. Competitive titration with edta has been followed by monitoring the Yb luminescence, since the edta complex is non-luminescent, contrary to the chelate formed with Fx. After addition of 5 equivalents of edta to a solution of [Yb(fx)] in Tris-HCl buffer, the Yb luminescence intensity decreases to 12% of its initial value. The thermodynamic stability of the fluorexon chelate is, therefore, comparable to [Yb(edta)] . In addition, the luminescence decay after addition of edta aliquots is relatively slow, the estimated rate constant being 7.1 x 10" s indicating a reasonably high kinetic stability of the fluorexon chelate. 

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. 

Figure 7 pH dependence of conditional stability constants, of some analytically important EDTA-metal complexes. The dashed line indicates the effect of 1 mol I ammonia/ammonium ion buffer on the conditional constants of the Zn-EDTA complex. 

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