Equivalence point EDTA titrations

The first task in calculating the titration curve is to determine the volume of EDTA needed to reach the equivalence point. At the equivalence point we know that... 

The equivalence point of a complexation titration occurs when stoichiometri-cally equivalent amounts of analyte and titrant have reacted. For titrations involving metal ions and EDTA, the equivalence point occurs when Cm and Cedxa are equal and may be located visually by looking for the titration curve s inflection point. 

Ethylenediaminetetra-acetic acid, largely as the disodium salt of EDTA, is a very important reagent for complex formation titrations and has become one of the most important reagents used in titrimetric analysis. Equivalence point detection by the use of metal-ion indicators has greatly enhanced its value in titrimetry. 

In acid-base titrations the end point is generally detected by a pH-sensitive indicator. In the EDTA titration a metal ion-sensitive indicator (abbreviated, to metal indicator or metal-ion indicator) is often employed to detect changes of pM. Such indicators (which contain types of chelate groupings and generally possess resonance systems typical of dyestuffs) form complexes with specific metal ions, which differ in colour from the free indicator and produce a sudden colour change at the equivalence point. The end point of the titration can also be evaluated by other methods including potentiometric, amperometric, and spectrophotometric techniques. 

A. Direct titration. The solution containing the metal ion to be determined is buffered to the desired pH (e.g. to PH = 10 with NH4-aq. NH3) and titrated directly with the standard EDTA solution. It may be necessary to prevent precipitation of the hydroxide of the metal (or a basic salt) by the addition of some auxiliary complexing agent, such as tartrate or citrate or triethanolamine. At the equivalence point the magnitude of the concentration of the metal ion being determined decreases abruptly. This is generally determined by the change in colour of a metal indicator or by amperometric, spectrophotometric, or potentiometric methods. 

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. 

The almost colourless leuco form of the base passes upon oxidation into the strongly coloured indamine. When titrating iron(III) at a pH of about 3 and the colourless hydrochloride of the leuco base is added, oxidation to the violet-blue indamine occurs with the formation of an equivalent amount of iron(II). At the end point of the EDTA titration, the small amount of iron(II) formed when the indicator was introduced is also transformed into the Fe(III)-EDTA complex FeY-, whereupon the blue indamine is reduced back to the leuco base. 

Pipette 25.0 mL of the 0.01 M calcium ion solution into a 250mL conical flask, dilute it with about 25 mL of distilled water, add 2mL buffer, solution, 1 mL 0.1M Mg-EDTA, and 30-40mg solochrome black/potassium nitrate mixture. Titrate with the EDTA solution until the colour changes from wine red to clear blue. No tinge of reddish hue should remain at the equivalence point. Titrate slowly near the end point. 

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. 

At the start of the experiment, a small amount of indicator (In) is added to the colorless solution of Mg2+ to form a red complex. As EDTA is added, it reacts first with free, colorless Mg2+. When free Mg2+ is used up, the last EDTA added before the equivalence point displaces indicator from the red Mgln complex. The change from the red Mgln to blue unbound In signals the end point of the titration (Demonstration 12-1). 

Formation constants for EDTA are expressed in terms of [Y4-], even though there are six protonated forms of EDTA. Because the fraction (aY4 1 of free EDTA in the form Y4 depends on pH, we define a conditional (or effective) formation constant as K = aYj Kf = MY" 4 /[M"+ [EDTA], This constant describes the hypothetical reaction Mn+ + EDTA MY 1-4, where EDTA refers to all forms of EDTA not bound to metal ion. Titration calculations fall into three categories. When excess unreacted M"+ is present, pM is calculated directly from pM = — log M l+]. When excess EDTA is present, we know both [MY"-4] and [EDTA], so IM"+] can be calculated from the conditional formation constant. At the equivalence point, the... 

Calcium ion was titrated with EDTA at pH 11. using Cal-magite as indicator (Table 12-3). Which is the principal species of Calmagite at pH 11 What color was observed before the equivalence point After the equivalence point ... 

Would tris(2,2 -bipyridine)iron be a useful indicator for the titration of Sn2+ with Mn(EDTA)- Hint The potential at the equivalence point must be between the potentials for each redox couple.)... 

During the course of the titration the concentration of HgY2- remains essentially constant because it is so much more stable than the MY2- complex. The is determined mainly by the ratio [Afu(OH2)2+]/[MY2 ], which changes slowly in the middle of the titration but rapidly near the equivalence point as the concentration of Mn(OH2)2+ drops to a small value. This gives a sharp potential change that signals the endpoint. The method is general and can be applied to most cations that form soluble EDTA complexes that are appreciably less stable than the mercury(II)-EDTA complex. 

From (11-35), and as illustrated schematically by Figure 11-9A, at constant [CaY ] the electrode potential decreases as calcium ion concentration increases. During an EDTA titration, pCa increases sharply near the equivalence point (Figure 11-9.B). If we consider now the entire titration and take into account the sharply increasing value of the ratio [CaY ]/[Ca ] during the early parts of the titration, the potential... 

Calculation of the Cation Concentration in EDTA Solutions In an EDTA titration, we are interested in finding the cation concentration as a function of the amount of titrant (EDTA) added. Prior to the equivalence point, the cation is in excess, and its concentration can be found from the reaction stoichiometry. At the equivalence point and in the postequivalence-point region, however, the... 

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. 

Figure 1 7-6 EDTA titration curves for 50.0 mL of 0,00500 M Ca-+ (i cav = 1.75 X 10 ) and Mg2+ = 1.72 X 10 ) at pH 10.0, Note that because of the larger formation constant, the reaction of calcium ion with EDTA is more complete, and a larger change occurs in the equivalence-point region. The shaded areas show the transition range for the indicator Eriochrome Black T.

The metal complexes of Eriochrome Black T are generally red, as is H2ln . Thus, for metal ion detection, it is necessary to adjust the pH to 7 or above so that the blue form of the species, HIn , predominates in the absence of a metal ion. Until the equivalence point in a titration, the indicator complexes the excess metal ion so that the solution is red. With the first slight excess of EDTA, the solution turns blue as a consequence of the reaction... 

Eriochrome Black T forms red complexes with more than two dozen metal ions, but the formation constants of only a few are appropriate for end point detection. As shown in Example 17-5, the applicability of a given indicator for an EDTA titration can be determined from the change in pM in the equivalence-point region, provided that the formation constant for the metal indicator complex is known. ... 

The Pb + and Mg are then titrated with standard EDTA. After the equivalence point has been reached, a solution of the complexing agent BAL (2-3-dimercapto-1-propanol, CH2SHCHSHCH2OH), which we will write as R(SH)2, is added to the solution. This bidentate ligand reacts selectively to form a complex with Pb- that is much more stable than PbY- ... 

See Problem 21 for a spreadsheet calculation of the Ca-EDTA titration curve in Figure 9.3 at pH 10. As with calculated acid-base titration curves, the calculations here break down very near the equivalence point due to simplifying assumptions we have made. 

CaEDTA "] = i.93 X 10" M. . 5% of Cj.ca is not complexed at the equivalence point for calcium. Note that in the standard procedure for calcium titration, the pH is such that OH" will complex some of the Mg ". Thus, less EDTA than calculated above will be associated with the Mg " " and more will be associated with the calcium. 

The presence of Mg (of Mg-EDTA which is transformed by Czi to Mg ) renders Eriochrome Black T a useful indicator for the titration of Ca at pH 10, since the indicator complex which it forms is not dissociated until after all of the Ca is titrated. At pH 10, pCa at the equivalence point is 6.1. The value of log P Mgin much closer to this pCa than the... 

The use of Eriochrome Black T as an indicator in the Zn -EDTA titration illustrates a case in which the indicator metal-complex is so stable that the color change occurs after the equivalence point, (log P nin equivalence point). Since log znin changes more rapidly with pH than does pZn (Why ), it is possible to adjust the pH to reduce the difference between them to a reasonably small value. (Table 9.2). 

C. (a) Explain why the change from red to blue in Reaction 13-2 occurs suddenly at the equivalence point instead of gradually throughout the entire titration, (b) EDTA buffered to pH 5 was titrated with standard Pb, with xylenol orange as indicator (Table 13-2). (i) Which is the principal species of the indicator at pH 5 (ii) What color was observed before the equivalence point (iii) What color was observed after the equivalence point (iv) What would the color change be if the titration were conducted at pH 8 instead of pH 5 ... 

Figure 13-9 Three regions in an EDTA titration of 50.0 mL of 0.050 0 M Mg or Ca with 0.050 0 M EDTA at pH 10. Region 2 is the equivalence point. The concentration of free M decreases as the titration proceeds.

---