Complexometric calculations

Karcher and KruII [30] used complexometric calculations to determine the mobile phase concentrations of HIBA and tartaric acid needed to fine tune their separation of eleven metal cations on C and Ci silica-based reverse-phase columns which had been dynamically modified with n-octanesulfonate. Isocratic elution was used to separate the metals into three distinct windows each window corresponding to one of the three dififerent valence states spanned by the eleven metal cations. The metals eluted in the order of increasing valence, with the sole exception of La(III) which did not elute within the trivalent ion window. Figure 6.15 illustrates the separation of ZrdV), Ga(III), Sc(III), Y(III), Al(III), In(III), Zn(II), La(III), Cd(II), Ca(II), and Mg(II) using a Cis column and an eluent comprised of 2.27 mM n-octanesulfonate, 8.18 mM tartaric acid, 52.9 mM HIBA, and 10.7% (v/v) methanol. 

The authors indicate that the elution orders for samples of metal ions can be accurately predicted using complexometric calculations. The algorithm used requires inputs of mobile phase ligand concentrations, mobile phase pH, and the formation constants for reactions between metal ions and the mobile phase ligands. In addition to their own experimental results, the authors have successfully calculated elution orders for many of the literature reported in situ complexation-based separations which have been achieved on either cation-exchange or d3mamically modified reverse-phase columns [31]. 

Now that we know something about EDTA s chemical properties, we are ready to evaluate its utility as a titrant for the analysis of metal ions. To do so we need to know the shape of a complexometric EDTA titration curve. In Section 9B we saw that an acid-base titration curve shows the change in pH following the addition of titrant. The analogous result for a titration with EDTA shows the change in pM, where M is the metal ion, as a function of the volume of EDTA. In this section we learn how to calculate the titration curve. We then show how to quickly sketch the titration curve using a minimum number of calculations. 

Sketching the Titration Curve As we have done for acid-base, complexometric titrations, and redox titrations, we now show how to quickly sketch a precipitation titration curve using a minimum number of calculations. 

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. 

Common chemical titrations include acid-base, oxidation-reduction, precipitation, and complexometric analysis. The basic concepts underlying all titration are illustrated by classic acid-base titrations. A known amount of acid is placed in a flask and an indicator added. The indicator is a compound whose color depends on the pH of its environment. A solution of base of precisely known concentration (referred to as the titrant) is then added to the acid until all of the acid has just been reacted, causing the pH of the solution to increase and the color of the indicator to change. The volume of the base required to get to this point in the titration is known as the end point of the titration. The concentration of the acid present in the original solution can be calculated from the volume of base needed to reach the end point and the known concentration of the base. 

The above calculations apply to all types of reactions, including acid-base, redox, precipitation, and complexometric reactions. The primary requirement before making calculations is to know the ratio in which the substances react, that is, start with a balanced reaction. 

Following is a hst of typical precipitation and complexometric titration reactions and the factors for calculating the milligrams of analyte from millimoles of titrant. ... 

These formulas are useful calculations involving the precipitation and complexometric titrations described in Chapters 8 and 11. 

In the first mode, known portions of the polymer were equilibrated with solutions of CaCl2 and HCl, as well as with their mixtures of known concentrations. The final composition of the bulk solutions in equilibrium with the polymeric phase was determined by titrating the excess HCl acid with NaOH and hy complexometric titration of the Ca ions with ethylenediamine tetraacetate (EDTA). From these data the concentrations of the electrolytes within the porous space of the polymeric material were calculated and then the apparent phase distribution coefficients k of HCl and CaCl2, defined as the ratio between the equihbrium concentrations of the corresponding electrolytes within and outside the polymeric beads. These calculations are strongly facifitated by the outstanding property of the neutral hypercrossfinked polystyrene sorbents, namely that their swelling does not depend on the electrolyte concentration, so that the volume of the porous space remains constant in all experiments. Thus,... 

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. 

The titration curve developed in Example 9.1 was hypothetical. In practice, complexometric titrations are conducted at pH values at which the titrant is protonated to some extent. Often, an auxiliary complexing agent such as ammonia is present so that there are various metal-containing species in addition to just the simple hydrated metal ions. Calculations of such titration curves can be readily accomplished within the framework of the material presented in Chapter 5, particularly that dealing with p, the conditional constant. In the treatment that follows, side reactions of the titration ligand and metal ion will be considered using the familiar Ul and... 

As shown in Chapter 8, selected individual points in complexometric titration curves can be calculated without recourse to the titration curve equation by (a) performing the reaction stoichiometry at the desired point, (b) recognizing the type of equilibrium calculation represented by the... 

This approach is effective not only for calculations involving individual solutions. The PBE and other balance equations apply to titration curve calculations as well. A single, easily derived, equation suffices to describe all points on a titration curve, be it for acid-base, complexometric, redox, or precipitation titrations. 

Weigh out accurately about 0.2 g of dried cement and treat as above for the determination of the alkalis. Pipette 25.0 cm of the solution into a conical flask and use the same procedure as in Sec.3.4.1 for the complexometric determination of calcium. Calculate from an average of two concordant titres, the concentration of calcium in the solution and hence the % Ca in the dried cement expressed as CaO%. 

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