Reduction titration curve

Because most redox indicators respond to changes in electrode potential, the vertical axis in oxidation/reduction titration curves is generally an electrode potential instead of the logarithmic p-functions that were used for complex formation and... 

In earlier chapters, we considered the effects of reactant concentrations and completeness of the reaction on titration curves. Here, we describe the effects of these variables on oxidation/reduction titration curves. 

How is an oxidation/reduction titration curve generated through the use of standard electrode potentials for the analyte species and the volumetric titrant ... 

Shown in Fig. 11.12 is the computed reduction titration curve of a model groundwater, assuming the titrant is dissolved organic carbon (DOC). Note that the pH remains in a narrow range between 6.5 and 7.5 as the water drops in pE and Eh. 

Thus in 2bN sulfuric acid the semiquinone ion III is so stable that the oxidation-reduction titration curve shows a definite step at the 50 per cent point and the color of the semiquinone is quite clearly apparent throughout one half of the titration. 

Figure 14. Oxidation-reduction titration curves of three reducing agents, showing midpoint potential Eq.

In Sections 10.11-10.16 it is shown how the change in pH during acid-base titrations may be calculated, and how the titration curves thus obtained can be used (a) to ascertain the most suitable indicator to be used in a given titration, and (b) to determine the titration error. Similar procedures may be carried out for oxidation-reduction titrations. Consider first a simple case which involves only change in ionic charge, and is theoretically independent of the hydrogen-ion concentration. A suitable example, for purposes of illustration, is the titration of 100 mL of 0.1M iron(II) with 0.1M cerium(IV) in the presence of dilute sulphuric acid ... 

It is possible to titrate two substances by the same titrant provided that the standard potentials of the substances being titrated, and their oxidation or reduction products, differ by about 0.2 V. Stepwise titration curves are obtained in the titration of mixtures or of substances having several oxidation states. Thus the titration of a solution containing Cr(VI), Fe(III) and V(V) by an acid titanium(III) chloride solution is an example of such a mixture in the first step Cr(VI) is reduced to Cr(III) and V(V) to V(IV) in the second step Fe(III) is reduced to Fe(II) in the third step V(IV) is reduced to V(III) chromium is evaluated by difference of the volumes of titrant used in the first and third steps. Another example is the titration of a mixture of Fe(II) and V(IV) sulphates with Ce(IV) sulphate in dilute sulphuric acid in the first step Fe(II) is oxidised to Fe(III) and in the second jump V(IV) is oxidised to V(V) the latter change is accelerated by heating the solution after oxidation of the Fe(II) ion is complete. The titration of a substance having several oxidation states is exemplified by the stepwise reduction by acid chromium(II) chloride of Cu(II) ion to the Cu(I) state and then to the metal. 

This concept allows the shape of the titration curves to be explained by postulating that the chloroform droplet size decreases as the interfacial tension (ift) between the aqueous and chloroform phases is decreased by the presence of active surfactant. As the endpoint in a titration is approached the amount of active SDBS decreases as it complexes with the injected hyamine. The reduction in the amount of active surfactant material results in an increase... 

Goldaman, J.A., Oxidation—Reduction Equilibria and Titration Curves , In Treatise on Analytical Chemistry, ed. by I.M. Kolthoff and PJ. Elving, 2nd ed., Vol. 3, New York, Interscience Publications, John Wiley Sons Inc., 1983. 

Schwertmann, 1993). Such soils are characterized by a hydraulic conductivity somewhere in the profile which is too low to cope with the high rainfall, so that all pores will be filled with water for certain periods of time (see above). In this case, the oxygen supply is limited by the low level of O2 dissolved in the soil water (46 mg O2 at 25 °C) and reduction of Mn-oxides, nitrate and Fe oxides sets in. Soils containing Fe oxides are, therefore, redox-buffered (poised). The redox titration curve (Fig. 16.14) of a soil with 23 g kg Fe as Fe oxides shows buffering at two different pe -1- pH levels, one at ca. 11 and another at ca. 9, which indicate the presence of a more reducible (e. g. ferrihydrite) and a less reducible (e. g. goethite) Fe oxide, respectively, in accordance with their different solubilities (see Chap. 9). 

We now turn our attention to details of precipitation titrations as an illustration of principles that underlie all titrations. We first study how concentrations of analyte and titrant vary during a titration and then derive equations that can be used to predict titration curves. One reason to calculate titration curves is to understand the chemistry that occurs during titrations. A second reason is to learn how experimental control can be exerted to influence the quality of an analytical titration. For example, certain titrations conducted at the wrong pH could give no discernible end point. In precipitation titrations, the concentrations of analyte and titrant and the size of Ksp influence the sharpness of the end point. For acid-base titrations (Chapter 11) and oxidation-reduction titrations (Chapter 16). the theoretical titration curve enables us to choose an appropriate indicator. 

The appearance of a cathodic limiting current after the equivalence point reflects the reduction of excess Ag+ titrant. Figure 3.43B shows resulting amperometric titration curves for two values of applied potential. Their shapes are determined by the behavior of the limiting current of the voltammogram at the particular potential during the titration. 

Figure 4.5 illustrates the titration curve that is obtained with an amperometric indicating system for the titration of lead ion with dichromate ion. If the applied potential is set on the plateau for the reduction of lead ion (approximately -0.5 V vs. SCE), curve a in Figure 4.5, will result. In contrast, if the applied potential is set at 0 V versus SCE, no current will flow until the point when excess chromate ion exists in the solution curve b is indicative of the titration curve that would be obtained. 

None of the redox titration results in previous papers showed the completely synchronized titration curves in the three wavelength regions as given in Fig. 9. This inconsistency is most likely caused by the quality of the purified preparation. The fast-form preparation was used for the experiments published in 1999 (Fig. 9), whereas other published results were likely to be obtained from the slow form or mixtures of slow and fast forms. The reduction rates of hemes a and as with dithionite are identical in the fast form, but reduction of heme as is much slower than that of heme a in the slow form (Moody, 1996). 

Fig. 9. A reductive titration of the crystalline bovine heart cytochrome c oxidase with dithionite. Absolute spectra for each oxidation state are shown for the Soret (A) and visible (B) regions. The difference spectra against the spectrum in the fully reduced state are given for the near-infrared region (C). The insets show titration curves against the electron equivalent per enzyme. The reaction mixture contained 7.5 jlM bovine heart cytochrome c oxidase in 0.1 M sodium phosphate buffer, pH 7.4. The enzyme preparation was stabilized with a synthetic non-ionic detergent, CH3(CH2)ii(0CH2CH2)80H. The light path was 1 cm.

BrOj" and reductants such as FeJ+, AsOj", Ti34 and S20 ". The titration curves are often unsymmetrical due to the stoichiometry of the reaction, but the error is small (<1 %) or can be eliminated by titrating standards. If the reduction potentials of two species differ by >0.2 V, separate inflections will be observed when titrating the mixture, e.g. the titration of Fe2 and Ti3 with potassium permanganate. 

Redox potentials have been determined for each of the steps of reduction of the trypsin-solubilized reductase (403) step 1, one electron consumed, Eo = —109 mV step 2, two electrons consumed. Eg = —276 mV and step 3, one electron consumed. Eg = —371 mV at pH 7.0, 26°. As expected, the redox potential of step 3 is more negative than the potential of the NADPH-NADP+ couple and was determined from the dithio-nite titration. The overall potentiometric—spectrophotometric titration curves could be very closely fitted with a computer-generated curve based on the assumptions of four one-electron reduction steps and octinction coefficients of 4.9 and 4.5 mM cm for the semiquinones, FliH and rijH the Eg values assumed for steps 2 and 3 were —270 and —290 mV. The precise fit was very sensitive to all of the assumptions (40 ) ... 

Figure 15-2 (left) depicts several titration curves of Fe(II) with permanganate. Beyond the end point the experimental curves differ from the theoretical shape, which is nearly flat beyond the end point (5-equivalent reduction). The essential symmetry of the curves suggests that the potential is determined by the Mn(III)-Mn(II) couple beyond the end point. Evidence for this behavior can be seen in solutions containing sulfate or phosphate, which tend to stabilize Mn(III) (Section 17-1). That sulfuric and phosphoric acids have about the same effect before and after the end point is consistent with the similarity of the behavior of the Mn(III)-Mn(II) and the Fe(III)-Fe(II) systems with respect to changes in activity coefficients as well as with respect to hydrolysis and complex formation. 

In Fe(II)-dichromate titrations, Winter and Moyer observed a time dependence of the potential after the end point. When potential readings were taken soon after each addition, an asymmetrical titration curve was observed, but when a time interval of 10 to 15 min was allowed after each addition, the curve approached the theoretical shape. We have noted that automatically recorded titration curves for the Fe(II)-dichromate titration show a considerably smaller potential jump than manually observed curves, the difference being due to lower potentials after the end point. But curves plotted with 15 s of waiting for each point differed only slightly from curves plotted with 150 s of waiting. Ross and Shain also studied the drift in potential of platinum electrodes with time and noted hysteresis effects in recorded potentiometric titration curves. These effects, due to oxidation and reduction of the platinum surface, are discussed below. 

Surface oxide formation undoubtedly is involved in the Fe(II)-dichromate titration curves, which Smith and Brandt found to be different when the direction of titration was reversed (Figure 15-2, right). Kolthoff and Tanaka found that the rate of oxidation with dichromate was slow, whereas the rate of reduction with Fe(II) was fast. Ross and Shain found the same sort of behavior and noted also that the rates of oxidation and reduction decreased in more dilute solutions. The oxidized surface in a dichromate solution may be largely covered with adsorbed dichromate, as chromium surfaces have been shown to be in some experiments with radio-chromium, so that it is relatively ineffective as an electron-transfer surface for the Fe(III)-Fe(II) system. 

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