EDTA solutions constants

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. 

Onken and Matheson (1982) studied kinetics of phosphorus dissolution in EDTA (ethylenediamine tetraacetic acid) solution for several soils. They examined eight kinetic models (Table 2.2) and found that phosphorus dissolution in EDTA solution was best described using the two-constant rate, Elovich, and differential rate equations as indicated by high r2 and low SE values. None of the models best described the dissolution for all soils. 

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 7.12 shows a chromatogram of the same sample in which EDTA is added to complex the iron(lll). The additional peak is from an iron(ll) impurity in the iron(II) solution used. Work thus far indicated that any metal ion that has an EDTA formation constant of about 10 - or higher should be masked effectively by adding EDTA to the sample. 

Any metal ion that has an EDTA formation constant higher than calcium or magnesium will interfere. Cyanide complexes strongly with copper, cobalt, nickel, zinc, and ferrous iron. Hydroxylamine or ascorbic acid is added to reduce iron to the ferrous state. If the solution is buffered to pH 10 before the indicator is added, then iron will not interfere because it precipitates as the hydroxide before it can react with the indicator or the EDTA. 

Cellular Viability and Concentration Determined by counting cells in nine squares of a Neubauer chamber after dying them with methylene blue solution. Viable cells were not colored, and dead cells were blue free yeast concentration was also obtained by filtering a known volume of cell suspension and drying the wet mass until the constant weight was achieved viability and concentration of immobilized yeast were determined as already described, after dissolution of the pectin gel [1.0 g of cured pellets was dissolved in 20 ml of 5% ethylenediaminetetraacetic acid (EDTA) solution at constant agitation]. The cell concentration was calculated as ... 

To determine the sum of calcium and magnesium in the second sample, adjust to pH 10.5 and a constant current of 20 pA and titrate in the way described above, but using EDTA solution, until the curve bends sharply. Evaluation is carried out graphically by extending straight lines from both curve... 

From the acid dissociation constants of EDTA (1,2) it can be calculated that between pH 7,4 and 6,6 an EDTA solution is represented to an extent greater than 99,8<7oby the forms HY and HjY . If the EDTA is in sufficient excess with respect to the cation studied, the pH of the solution, on complexing, will remain between the above limits and the experimental reaction will be described unambigously by the equation ... 

Aqueous molybdate-edta solutions have been studied by temperature jump in the pH range 7.25—8.25 at 25 °C and ionic strength 0.1 mol 1. The rate constant for... 

Exercise 5 The value of the conditional stability constant of the Mg-EDTA complex is 2.4 X 10 at pH = 9 in a solution buffered with the couple NH4VNH3. Calculate the fraction of complexed magnesium with EDTA when we mix 100 ml of a buffered solution of 10 mol/L magnesium ions with 100 ml of a 10 mol/L EDTA solution. 

Amperometry involves the measurements of currents at constant voltage apphed at the dropping mercury electrode. The value of electrode potential is chosen in such a way that only the metal ion is reduced. This method is generally used for the determination of metal ion present in aqueous solution. An aqueous solution of Zn" may be titrated with an EDTA solution at an applied electrode potential of —1.4 V. 

In the now familiar pattern discussed above, the titration involves the buret addition of EDTA solution to the metal ion solution, which generates a titration curve with an abrupt change in — log[M " ] (pM). This is governed by the equilibrium constant for the formation of the metal-EDTA complex ... 

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 colour change can be observed with the ions of Mg, Mn, Zn, Cd, Hg, Pb, Cu, Al, Fe, Ti, Co, Ni, and the Pt metals. To maintain the pH constant (ca 10) a buffer mixture is added, and most of the above metals must be kept in solution with the aid of a weak complexing reagent such as ammonia or tartrate. The cations of Cu, Co, Ni, Al, Fe(III), Ti(IV), and certain of the Pt metals form such stable indicator complexes that the dyestuff can no longer be liberated by adding EDTA direct titration of these ions using solochrome black as indicator is therefore impracticable, and the metallic ions are said to block the indicator. However, with Cu, Co, Ni, and Al a back-titration can be carried out, for the rate of reaction of their EDTA complexes with the indicator is extremely slow and it is possible to titrate the excess of EDTA with standard zinc or magnesium ion solution. 

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"+. 

Procedure. Charge the titration cell (Fig. 17.24) with 10.00 mL of the copper ion solution, 20 mL of the acetate buffer (pH = 2.2), and about 120mL of water. Position the cell in the spectrophotometer and set the wavelength scale at 745 nm. Adjust the slit width so that the reading on the absorbance scale is zero. Stir the solution and titrate with the standard EDTA record the absorbance every 0.50 mL until the value is about 0.20 and subsequently every 0.20 mL. Continue the titration until about 1.0 mL after the end point the latter occurs when the absorbance readings become fairly constant. Plot absorbance against mL of titrant added the intersection of the two straight lines (see Fig. 17.23 C) is the end point. 

The interaction between aequorin and a chelator must be carefully considered when estimating Ca2+ concentrations with aequorin in a calcium buffer containing EDTA or EGTA. This is particularly crucial when using a common calcium buffer system that contains a constant total concentration of a chelator in the buffer solutions of various Ca2+ concentrations in such a buffer system, a buffer of lower Ca2+ concentration contains a higher concentration of the free form of the chelator, resulting in an increased inhibition. 

Fig. 4.1.8 Influence of various calcium chelators on the relationship between Ca2 " concentration and the luminescence intensity of aequorin, at 23-25°C (panel A) in low-ionic strength buffers (I < 0.005) and (panel B) with 150 mM KC1 added. Buffer solutions (3 ml) of various Ca2+ concentrations, pH 7.05, made with or without a calcium buffer was added to 2 pi of 10 pM aequorin solution containing 10 pM EDTA. The calcium buffer was composed of the free form of a chelator (1 or 2mM) and various concentrations of the Ca2+-chelator (1 1) complex to set the Ca2+ concentrations (the concentration of free chelator was constant at all Ca2+ concentrations). The curves shown are obtained with 1 mM MOPS (A), 1 mM gly-cylglycine ( + ), 1 mM citrate (o), 1 mM EDTA plus 2mM MOPS ( ), 1 mM EGTA plus 2 mM MOPS ( ), 2 mM NTA plus 2 mM MOPS (V), and 2 mM ADA plus 2 mM MOPS (A). In the chelator-free buffers, MOPS and glycylglycine, Ca2+ concentrations were set by the concentration of calcium acetate. Reproduced with permission, from Shimomura and Shimomura, 1984. the Biochemical Society.

From this material, samples are cut and swelled to constant weight in a buffered saline solution prepared from 8.43 g sodium chloride (NaCl), 9.26 g boric acid (H3BO3), 1.0 g sodium borate (Na3B03), and 0.1 g of the disodium salt of the dihydrate of ethylenediaminetetraacetic acid [Na2 EDTA -(/ 0)21 ini L of distilled water. 

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