Complexometric Titrations

Complexometric Titrations EDTA After the Co2+ in your unknown is separated from Fe3+ it can be quantitatively measured by a complexometric titration. In such a titration, a metal ion is titrated with a ligand that readily forms a complex with the metal ion. Ethylenediaminetetraacetic acid (EDTA) is used to complex the Co2+. EDTA is a polydentate ligand. This means that it can donate more than one pair of nonbonding electrons when forming a complex with a metal ion. In fact, EDTA can form 6 bonds with a metal ion and thus occupy all the octahedral positions associated with a metal ion complex. The structure of the fully protonated form of EDTA (H6Y2+) is shown in Figure 18.1. Figure 18.1 Structure of fully protonated EDTA EDTA can lose the six highlighted hydrogens when it complexes with a metal ion. The nonbonding electron pairs on the nitrogen and on oxygen in the four carboxyl groups are used to form the metal ion - EDTA complex. Notice that as in ethylenediamine and phenanthroline there is a 2-carbon bridge between adjacent electrons pairs used by EDTA for complexing metal ions. The fully ionized form of EDTA (Y4-) will form a 1:1 complex with virtually all metal ions in solution. The formation of this complex can be represented by Mn+ + Y4- º MYn-4 with an equilibrium constant (Kf) given by The values of log Kf for various metal ions are given in the table below. Notice that Kf is often quite large, especially for highly-charged metal ions. Kf is greater than 1016 for the formation of the EDTA-Co2+ complex. This makes for a quantitative titration reaction, however, it assumes that all of the EDTA is in the fully ionized form. For any polyprotic acid like EDTA, the amount of any particular ionized form (such as Y4-) depends on the solution pH. The fraction of EDTA that is in the Y4- form (αY4-) is given by where [EDTA] is the total concentration of uncomplexed EDTA in solution. This fraction can be calculated for a given pH from the Ka values for the polyprotic EDTA. The values of αY4- in the table below indicate that, as expected, only in very basic solution (pH>12) is the EDTA fully ionized. However, it is difficult to titrate metal ions in basic solution because of competing reactions such as the formation of metal hydroxides. To account for the effect of pH on the ionization of the EDTA, it is convenient to rewrite the equilibrium expression for the complex formation as and to define a conditional formation constant (KfN) which is given by The value of KfN provides information about how quantitative the complex formation is at a given pH. Since αY4- is very small at acidic pH values, there is generally some minimum pH for which the complex formation reaction is sufficiently quantitative. Using an arbitrarily chosen value of KfN = 106 for an effective titration, a plot of the minimum pH to achieve this for various metal ions is given in Figure 18.2. For Co2+, this corresponds to a pH greater than 4. In our experiment, the titration solution is buffered at a pH = 5.8 to ensure a quantitative reaction between Co2+ and EDTA. It is useful to note that the fully ionized form of EDTA (Y4-) is not the only form of EDTA that reacts with metal ions. Figure 18.2 Minimum pH needed for a conditional formation constant of 106 The particular pH used to buffer the titration solution was also chosen because of the metal ion indicator used to signal the end point in the titration. A metal ion indicator is a substance that changes color when it binds to metal ions in solution. Metal ion indicators tend to be polyprotic complexing agents. Xylenol orange (Fig. 18.3) is used for the cobalt titration. This indicator is red when it forms a complex with metal ions such as Co2+. However, the color of the unbound (free) indicator depends on pH and which protonated form is predominantly present (Fig. 18.3). Above pH = 6.7, free xylenol orange is violet and the color change from red to violet is not sharp. However, below pH = 6.7 the free indicator color is yellow and the change from red to yellow is much easier to see. Figure 18.3 Structure of the metal ion indicator xylenol orange A small amount of xylenol orange is added initially to the cobalt solution and it forms a complex with some of the Co2+. As EDTA is added to this solution, the EDTA complexes the free Co2+. Just before the equivalence point, the free Co2+ is used up and the EDTA starts to remove cobalt ion from the indicator. As the xylenol orange loses the metal ion, its color changes. Obviously, the indicator must not bind the metal ion as strongly as the complexing agent (EDTA) used for the titration. Check for Understanding 18.1 Solutions 1. What is the concentration of uncomplexed Co2+ in solution at the equivalence point in an EDTA titration if 25.00 mL of 0.0100 M EDTA solution is needed to titrate the sample? Assume the unknown solution volume is 75.00 mL and that it is buffered to a pH = 6.0. 2. How many ppth of the original Co2+ in the unknown solution does this uncomplexed cobalt represent?