EDTA complexes stability constants

Even plutonyl(V) is strongly complexed by EDTA. The stability constant of Pu02Y has been determined by three methods ... 

A comparison of the formation constants for metal-fulvic acid complexes shown in Table 5-13 with those values for the complexes of the same metals with NTA (Table 5-10) would seem to indicate that the metal-fulvic acid complexes are much weaker than the metal-NTA complexes. Indeed some of the values of the constants for metal-fulvic acid complexes approach the values for the very weak inorganic ion pairs (e.g., MgCOs , log = 2.9 CaS04 logff = 3.4). Morgan has pointed out that the values for the metal-fulvic acid complexes were measured at much lower pH values than the K values for other complexing agents such as NTA and EDTA. The stability constants for NTA complexes reported in Table 5-10 are for the totally deprotonated species,... 

The formation constants of EDTA complexes are gathered in Table 11.34. Based on their stability, the EDTA complexes of the most common metal ions may be roughly divided into three groups ... 

EDTA Must Compete with Other Ligands To maintain a constant pH, we must add a buffering agent. If one of the buffer s components forms a metal-ligand complex with Cd +, then EDTA must compete with the ligand for Cd +. For example, an NH4+/NH3 buffer includes the ligand NH3, which forms several stable Cd +-NH3 complexes. EDTA forms a stronger complex with Cd + and will displace NH3. The presence of NH3, however, decreases the stability of the Cd +-EDTA complex. 

Compared to later elements in their respective transition series, scandium, yttrium and lanthanum have rather poorly developed coordination chemistries and form weaker coordinate bonds, lanthanum generally being even less inclined to form strong coordinate bonds than scandium. This is reflected in the stability constants of a number of relevant 1 1 metal-edta complexes ... 

Some values for the stability constants (expressed as logX) of metal-EDTA complexes are collected in Table 2.4 these apply to a medium of ionic strength 7 = 0.1 at 20 °C. 

Table 2.4 Stability constants (as log K) of metal-EDTA complexes...

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. 

The factor at can be calculated from the known dissociation constants of EDTA, and since the proportions of the various ionic species derived from EDTA will be dependent upon the pH of the solution, a will also vary with pH a plot of log a against pH shows a variation of logoc = 18 at pH = 1 to loga = 0 at pH = 12 such a curve is very useful for dealing with calculations of apparent stability constants. Thus, for example, from Table 2.4, log K of the EDTA complex of the Pb2+ ion is 18.0 and from a graph of log a against pH, it is found that at a pH of 5.0, log a = 7. Hence from equation (30), at a pH of 5.0 the lead-EDTA complex has an apparent stability constant given by ... 

The extent of hydrolysis of (MY)(n 4)+ depends upon the characteristics of the metal ion, and is largely controlled by the solubility product of the metallic hydroxide and, of course, the stability constant of the complex. Thus iron(III) is precipitated as hydroxide (Ksal = 1 x 10 36) in basic solution, but nickel(II), for which the relevant solubility product is 6.5 x 10 l8, remains complexed. Clearly the use of excess EDTA will tend to reduce the effect of hydrolysis in basic solutions. It follows that for each metal ion there exists an optimum pH which will give rise to a maximum value for the apparent stability constant. 

EDTA is a very unselective reagent because it complexes with numerous doubly, triply and quadruply charged cations. When a solution containing two cations which complex with EDTA is titrated without the addition of a complex-forming indicator, and if a titration error of 0.1 per cent is permissible, then the ratio of the stability constants of the EDTA complexes of the two metals M and N must be such that KM/KN 106 if N is not to interfere with the titration of M. Strictly, of course, the constants KM and KN considered in the above expression should be the apparent stability constants of the complexes. If complex-forming indicators are used, then for a similar titration error KM/KN z 108. 

This reaction will proceed if the metal indicator complex M-In is less stable than the metal-EDTA complex M EDTA. The former dissociates to a limited extent, and during the titration the free metal ions are progressively complexed by the EDTA until ultimately the metal is displaced from the complex M-In to leave the free indicator (In). The stability of the metal-indicator complex may be expressed in terms of the formation constant (or indicator constant) Ku ... 

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. 

There is an appreciable difference between the stability constants of the CDTA complexes of barium (log K = 7.99) and calcium (log K = 12.50), with the result that calcium may be titrated with CDTA in the presence of barium the stability constants of the EDTA complexes of these two metals are too close together to permit independent titration of calcium in the present of barium. 

In a similar manner, in a solution containing the species Hg2+, HgY2-, MY,n 4)+ and M"+, where Y is the complexing agent EDTA and M"+ is a metallic ion which forms complexes with it, the concentration of the mercury ion is controlled by the stability constants of the complex ions MYhigh stability constant), and the concentration of the metal ions M"+. Hence, a mercury electrode placed in this solution will acquire a potential which is determined by the concentration of the ion M"+. 

Discussion. Salicylic acid and iron(III) ions form a deep-coloured complex with a maximum absorption at about 525 nm this complex is used as the basis for the photometric titration of iron(III) ion with standard EDTA solution. At a pH of ca 2.4 the EDTA-iron complex is much more stable (higher stability constant) than the iron-salicylic acid complex. In the titration of an iron-salicylic acid solution with EDTA the iron-salicylic acid colour will therefore gradually disappear as the end point is approached. The spectrophotometric end point at 525 nm is very sharp. 

The size of the EDTA ring structure complex (not too big or too small), plus the multiple points of attachment, provides a very stable complex (high stability constant) that overcomes reaction tendencies to reach equilibrium via the lowest possible energy state or maximum state of disorder (entropy). 

The transport of EDTA into a bacterial strain capable of its degradation has been examined (Witschel et al. 1997). Inhibition was observed with DCCD (ATPase inhibitor), nigericin (dissipates ApH), but not valinomycin (dissipates Av /), and was dependent on the stability constant of metal-EDTA complexes. 

For a complex-forming metal ion detectable by its own metal electrode, e.g. in the titration of Cu2+ with EDTA by means of a Cu electeode and a double junction calomel electrode, p/ curves are obtained of a nature comparable to those in Fig. 2.20 and with a Cu range of about 20 (cf., stability constant... 

Majer65 in 1936 proposed measuring, instead of the entire polarographic curve, only the limiting current at a potential sufficiently high for that purpose if under these conditions one titrates metal ions such as Zn2+, Cd2+, Pb2+, Ni2+, Fe3+ and Bi3+ with EDTA66, one obtains a titration as depicted in Fig. 3.55 i, decreases to a very low value, in agreement with the stability constant of the EDTA-metal complex and the titration end-point is established by the intersection of the ij curves before and after that point correction of the i values for alteration of the solution volume by the titrant increments as in conductometric titration is recommended. 

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. 

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