Potentiometry redox titrations

A major branch of analytical chemistry uses electrical measurements of chemical processes at the surface of an electrode for analytical purposes. For example, hormones released from a single cell can be measured in this manner. Principles developed in this chapter provide a foundation for potentiometry, redox titrations, electrogravimetric and coulometric analysis, voltammetry, and amperometry in the following chapters.1-2... 

Finding the End Point Potentiometrically Another method for locating the end point of a redox titration is to use an appropriate electrode to monitor the change in electrochemical potential as titrant is added to a solution of analyte. The end point can then be found from a visual inspection of the titration curve. The simplest experimental design (Figure 9.38) consists of a Pt indicator electrode whose potential is governed by the analyte s or titrant s redox half-reaction, and a reference electrode that has a fixed potential. A further discussion of potentiometry is found in Chapter 11. 

Redox titrations are based on the transfer of electrons between the titrant and the analyte. These types of titrations are usually followed by potentiometry, although dyes which change colour when oxidised by excess titrant may be used. 

The characteristics of redox reactions in non-aqueous solutions were discussed in Chapter 4. Potentiometry is a powerful tool for studying redox reactions, although polarography and voltammetry are more popular. The indicator electrode is a platinum wire or other inert electrode. We can accurately determine the standard potential of a redox couple by measuring the electrode potential in the solution containing both the reduced and the oxidized forms of known concentrations. Poten-tiometric redox titrations are also useful to elucidate redox reaction mechanisms and to obtain standard redox potentials. In some solvents, the measurable potential range is much wider than in aqueous solutions and various redox reactions that are impossible in aqueous solutions are possible. 

Construct a coulometric titration curve of 100.0 mL of a 1 M H2SO4 solution containing Fe(ll) titrated with Ce(lV) generated from 0.075 M Ce(lll). The titration is monitored by potentiometry. The initial amount of Fe(II) present is 0.05182 mmol. A constant current of 20.0 mA is used. Find the time corresponding to the equivalence point. Then, for about 10 values of time before the equivalence point, use the stoichiometry of the reaction to calculate the amount of Fe produced and the amount of Fe + remaining. Use the Nemst equation to find the system potential. Find the equivalence point potential in the usual manner for a redox titration. For about 10 times after the equivalence point, calculate the amount of Ce " produced from the electrolysis and the amount of Ce + remaining. Plot the curve of system potential versus electrolysis time. 

Redox Titrations Electrochemistry Chapter 19 Standard Electrode Potentials Chapter 20 Oxidation/ReductionTitrations Chapter 21 Potentiometry Chapter 17 Using Electrode Potentials Chapter 18 Oxidation/Reduction Titrations Chapter 19 Potentiometry... 

Oxidation-Reduction Titrations. Potentiometry can be used to follow reduction-oxidation (redox) titrations. For example, the oxidation of stannous ions by ceric ions follows the chemical reaction... 

With a low constant current -1 (see Fig. 3.71) one obtains the same type of curve but its position is slightly higher and the potential falls just beyond the equivalence point (see Fig. 3.72, anodic curve -1). In order to minimize the aforementioned deviations from the equivalence point, I should be taken as low as possible. Now, it will be clear that the zero current line (abscissa) in Fig. 3.71 yields the well known non-faradaic potentiometric titration curve (B B in Fig. 2.22) with the correct equivalence point at 1.107 V this means that, when two electroactive redox systems are involved, there is no real need for constant-current potentiometry, whereas this technique becomes of major advantage... 

S, spectroelectrochemistry P, photochemical reduction T, reductive titration CV, cyclic voltammetry R, redox potentiometry. 

Lead tetraacetate consumption is measured conveniently by iodometry.4 The reaction mixture is added to excess potassium iodide solution, usually in the presence of sodium acetate,6 and the iodine liberated is then titrated with standard thiosulfate. Oxidation may also be measured potentiometri-cally,78 210 211 a procedure especially useful for fast glycol groups,78 or with redox indicators.211... 

Use of the potential of a galvanic cell to measure the concentration of an electroactive species developed later than a number of other electrochemical methods. In part this was because a rational relation between the electrode potential and the concentration of an electroactive species required the development of thermodynamics, and in particular its application to electrochemical phenomena. The work of J. Willard Gibbs1 in the 1870s provided the foundation for the Nemst equation.2 The latter provides a quantitative relationship between potential and the ratio of concentrations for a redox couple [ox l[red ), and is the basis for potentiometry and potentiometric titrations.3 The utility of potentiometric measurements for the characterization of ionic solutions was established with the invention of the glass electrode in 1909 for a selective potentiometric response to hydronium ion concentrations.4 Another milestone in the development of potentiometric measurements was the introduction of the hydrogen electrode for the measurement of hydronium ion concentrations 5 one of many important contributions by Professor Joel Hildebrand. Subsequent development of special glass formulations has made possible electrodes that are selective to different monovalent cations.6"8 The idea is so attractive that intense effort has led to the development of electrodes that are selective for many cations and anions, as well as several gas- and bioselective electrodes.9 The use of these electrodes and the potentiometric measurement of pH continue to be among the most important applications of electrochemistry. 

It is an attractive feature of potentiometry that the equipment is rather inexpensive and simple one needs a reference electrode, an indicator electrode and a voltagemeasuring instrument with high input impedance. The potential measurement has to be accomplished with as low a current as possible because otherwise the potential of both electrodes would change and falsify the result. In the past, a widespread method was the use of the so-called Poggendorf compensation circuit. In most cases today, amplifier circuits with an input impedance up to 10 2 are used. The key element for potentiometry is the indicator electrode. Currently, ion-selective electrodes are commercially available for more than 20 different ions and almost all kinds of titrations (acid-base, redox, precipitation and complex titrations) can be indicated. In the following, some indicator electrodes and the origin of the electrode potentials will be described. 

In both, direct potentiometry and potentiometric titration redox reactions as well as equilibrium adjustments cause the function of the electrochemical cells. 

Titrimetric methods with potentiometric end point location can be applied when an electrode with the needed selectivity is not available. The precision and accuracy of potentiometric titrations are superior comparing it with the properties of direct potentiometry. However, the concentration range where potentiometric titration can be used effectively is narrower. A solution with analyte concentration below 1 mM seldom is determined by potentiometric titrations. Potentiometric end point location is most often employed in the case of acid-, base-, precipitate-, redox-, or complexometric titrations. 

Surprisingly, at first sight, redox indioators may also be used in some cases to detect the endpoint of a complexometric titration with EDTA. In fact, the endpoint of an EDTA titration may be accompanied by a ehange in the redox potential of the solution. When a mixture of Fe + and Fe + is titrated with EDTA, Fe + disappears before Fe + since Fe gives more stable eomplexes with EDTA than Fe + does. A simple inspection of Nernst s equation shows that in these conditions, the solution s redox potential decreases markedly, in particular at the equivalence point. The sharp change may be detected by potentiometry with a platinum electrode or with a redox indicator such as Variamine blue. 

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