Potentiometric titrations precipitation titration

Titrations that use the human eye as the primary detector are based on color change. Other types of titrations that rely on color change are oxidation-reduction, precipitation, and complexometric titrations. Other detectors indicate voltage or other types of changes such as potentiometric titrations, conductometric titrations and amperometric titrations—all of which require additional instrumentation—and may be quite colorless. 

Initial attempts at developing precipitation titration methods were limited by a poor end point signal. Finding the end point by looking for the first addition of titrant that does not yield additional precipitate is cumbersome at best. The feasibility of precipitation titrimetry improved with the development of visual indicators and potentiometric ion-selective electrodes. 

Potcntiomctric Titrations In Chapter 9 we noted that one method for determining the equivalence point of an acid-base titration is to follow the change in pH with a pH electrode. The potentiometric determination of equivalence points is feasible for acid-base, complexation, redox, and precipitation titrations, as well as for titrations in aqueous and nonaqueous solvents. Acid-base, complexation, and precipitation potentiometric titrations are usually monitored with an ion-selective electrode that is selective for the analyte, although an electrode that is selective for the titrant or a reaction product also can be used. A redox electrode, such as a Pt wire, and a reference electrode are used for potentiometric redox titrations. More details about potentiometric titrations are found in Chapter 9. 

The holistic thermodynamic approach based on material (charge, concentration and electron) balances is a firm and valuable tool for a choice of the best a priori conditions of chemical analyses performed in electrolytic systems. Such an approach has been already presented in a series of papers issued in recent years, see [1-4] and references cited therein. In this communication, the approach will be exemplified with electrolytic systems, with special emphasis put on the complex systems where all particular types (acid-base, redox, complexation and precipitation) of chemical equilibria occur in parallel and/or sequentially. All attainable physicochemical knowledge can be involved in calculations and none simplifying assumptions are needed. All analytical prescriptions can be followed. The approach enables all possible (from thermodynamic viewpoint) reactions to be included and all effects resulting from activation barrier(s) and incomplete set of equilibrium data presumed can be tested. The problems involved are presented on some examples of analytical systems considered lately, concerning potentiometric titrations in complex titrand + titrant systems. All calculations were done with use of iterative computer programs MATLAB and DELPHI. 

The indicator electrode employed in a potentiometric titration will, of course, be dependent upon the type of reaction which is under investigation. Thus, for an acid-base titration, the indicator electrode is usually a glass electrode (Section 15.6) for a precipitation titration (halide with silver nitrate, or silver with chloride) a silver electrode will be used, and for a redox titration [e.g. iron(II) with dichromate] a plain platinum wire is used as the redox electrode. 

The majority of potentiometric titrations involve chemical reactions which can be classified as (a) neutralisation reactions, (b) oxidation-reduction reactions, (c) precipitation reactions or (d) complexation reactions, and for each of these different types of reaction, certain general principles can be enunciated. 

Titrations can be carried out in cases in which the solubility relations are such that potentiometric or visual indicator methods are unsatisfactory for example, when the reaction product is markedly soluble (precipitation titration) or appreciably hydrolysed (acid-base titration). This is because the readings near the equivalence point have no special significance in amperometric titrations. Readings are recorded in regions where there is excess of titrant, or of reagent, at which points the solubility or hydrolysis is suppressed by the Mass Action effect the point of intersection of these lines gives the equivalence point. 

Mercury(II) chloranilate 700 Mercury(II) nitrate standard soln. of, 359 Mercury/mercury( II )-EDTA electrode (mercury electrode) 586 potentiometric titration of metallic ions with EDTA and, 588 prepn. of, 587 Mercury thiocyanate 700 Metaphosphoric acid in homogeneous precipitation, 426 Metal apparatus 93 Metal ion buffer 53... 

Potentiometric titrations - continued EDTA titrations, 586 neutralisation reactions, 578, 580 non-aqueous titrations, 589, (T) 590 oxidation-reduction reactions, 579, 581, 584 precipitation reactions, 579, 582 Potentiometry 548 direct, 548, 567 fluoride, D. of, 570 Potentiostats 510, 607 Precipitants organic, 437 Precipitate ageing of, 423 digestion of, 423... 

In fact, any type of titration can be carried out potentiometrically provided that an indicator electrode is applied whose potential changes markedly at the equivalence point. As the potential is a selective property of both reactants (titrand and titrant), notwithstanding an appreciable influence by the titration medium [aqueous or non-aqueous, with or without an ISA (ionic strength adjuster) or pH buffer, etc.] on that property, potentiometric titration is far more important than conductometric titration. Moreover, the potentiometric method has greater applicability because it is used not only for acid-base, precipitation, complex-formation and displacement titrations, but also for redox titrations. 

In the practice of potentiometric titration there are two aspects to be dealt with first the shape of the titration curve, i.e., its qualitative aspect, and second the titration end-point, i.e., its quantitative aspect. In relation to these aspects, an answer should also be given to the questions of analogy and/or mutual differences between the potentiometric curves of the acid-base, precipitation, complex-formation and redox reactions during titration. Excellent guidance is given by the Nernst equation, while the acid-base titration may serve as a basic model. Further, for convenience we start from the following fairly approximate assumptions (1) as titrations usually take place in dilute (0.1 M) solutions we use ion concentrations in the Nernst equation, etc., instead of ion activities and (2) during titration the volume of the reaction solution is considered to remain constant. 

Whereas in many instances potentiometric non-aqueous titrations of acids can show anomalies24 depending on the type of solvents and/or electrodes (owing to preferential adsorption of ions, ion pairs or complexes on the highly polar surface of the indicator electrode, or even adherence of precipitates on the latter), conductometric non-aqueous titrations, in contrast, although often accompanied by precipitate formation30, are not hindered by such phenomena sometimes, just as in aqueous titrations, the conductometric end-point can even be based on precipitate formation34. 

Conductometric titrations. Van Meurs and Dahmen25-30,31 showed that these titrations are theoretically of great value in understanding the ionics in non-aqueous solutions (see pp. 250-251) in practice they are of limited application compared with the more selective potentiometric titrations, as a consequence of the low mobilities and the mutually less different equivalent conductivities of the ions in the media concerned. The latter statement is illustrated by Table 4.7108, giving the equivalent conductivities at infinite dilution at 25° C of the H ion and of the other ions (see also Table 2.2 for aqueous solutions). However, in practice conductometric titrations can still be useful, e.g., (i) when a Lewis acid-base titration does not foresee a well defined potential jump at an indicator electrode, or (ii) when precipitations on the indicator electrode hamper its potentiometric functioning. 

Streng, W. H. Zoglio, M. A., Determination of the ionization constants of compounds which precipitate during potentiometric titration using extrapolation techniques, J. Pharm. Sci. 73, 1410-1414 (1984). 

The potentiometric titrations of [Cu1(MeCN)4](CIO4), AgIC104, and AuIC104 with (Bu4N)0H(in MeOH) are illustrated in Figure 4, and demonstrate that each process has one-to-one stoichiometry. The three systems form precipitates such that all of the metal is removed from solution at the equivalence point. Addition of excess "OH causes some dissolution of the CuOH and AgOH precipitates, and appears as a second step for the titration curve of Ag(I) (Figure 4b). 

Figure 4.6 Schematic representation of the apparatus required when monitoring a precipitation process via a potentiometric titration. The salt bridge is impregnated with a saturated solution of KNO3.

During a precipitation reaction, a potentiometric titration can also be employed, but here we generally determine an activity since the emf is related to the Nemst equation. For this reason, an absolute value of freference electrode should be known in this case. 

Potentiometric titration is actually a form of the multiple known subtraction method. The main advantage of titration procedures, similar to multiple addition techniques in general, is the improved precision, especially at high determinand concentrations. ISEs are suitable for end-point indication in all combination titrations (acid-base, precipitation, complexometric), provided that either the titrand or the titrant is sensed by an ISE. If both the titrant and the titrand are electro-inactive, an electrometric indicator must be added (for example Fe ion can be titrated with EDTA using the fluoride ISE when a small amount of fluoride is added to the sample solution [126]). 

Selig reported a potentiometric titration method for the analysis of procaine and some other organic cations precipitated by tetraphenylborate [67]. The development of ion selective coated-wire electrodes, and their application in the titration of procaine and other pharmaceutically important substances, was reported [68]. 

Acid-base, redox, precipitation and chelometric titrations are usually dealt with in textbooks on analytical chemistry. The titration curves in these titrations can be obtained potentiometrically by use of appropriate indicator electrodes, i.e. a pH-glass electrode or pH-ISFET for acid-base titrations, a platinum electrode for redox titrations, a silver electrode or ISEs for precipitation titrations, and ISEs for... 

By running a potentiometric precipitation titration, we can determine both the compositions of the precipitate and its solubility product. Various cation- and anion-selective electrodes as well as metal (or metal amalgam) electrodes work as indicator electrodes. For example, Coetzee and Martin [23] determined the solubility products of metal fluorides in AN, using a fluoride ion-selective LaF3 single-crystal membrane electrode. Nakamura et al. [2] also determined the solubility product of sodium fluoride in AN and PC, using a fluoride ion-sensitive polymer membrane electrode, which was prepared by chemically bonding the phthalocyanin cobalt complex to polyacrylamide (PAA). The polymer membrane electrode was durable and responded in Nernstian ways to F and CN in solvents like AN and PC. 

The extremely low solubility of lead phosphate in water (about 6 x 10 15m) again suggests potentiometric analysis. Selig57,59 determined micro amounts of phosphate by precipitation with lead perchlorate in aqueous medium. The sample was buffered at pH 8.25-8.75 and a lead-selective electrode was used to establish the end-point. The detection limit is about 10 pg of phosphorus. Anions which form insoluble lead salts, such as molybdate, tungstate or chromate, interfere with the procedure. Similar direct potentiometric titrations of phosphate by precipitation as insoluble salts of lanthanum(III), copper(II) or cadmium(II) are suggested, the corresponding ion-selective electrodes being used to detect the end-point. 

Tartaric Acid. Quantitative measures of total tartrate are useful in determining the amount of acid reduction required for high acid musts and in predicting the tartrate stability of finished wines. Three procedures may be used. Precipitation as calcium racemate is accurate (85), but the cost and unavailability of L-tartaric acid are prohibitive. Precipitation of tartaric acid as potassium bitartrate is the oldest procedure but is somewhat empirical because of the appreciable solubility of potassium bi-tartrate. Nevertheless, it is still an official AO AC method (3). The colorimetric metavanadate procedure is widely used (4, 6, 86, 87). Tanner and Sandoz (88) reported good correlation between their bitartrate procedure and Rebeleins rapid colorimetric method (87). Potentiometric titration in Me2CO after ion exchange was specific for tartaric acid (89). 

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