Titration redox indicators

A selected list of redox indicators will be found in Table 8.26. A redox indicator should be selected so that its if" is approximately equal to the electrode potential at the equivalent point, or so that the color change will occur at an appropriate part of the titration curve. If n is the number of electrons involved in the transition from the reduced to the oxidized form of the indicator, the range in which the color change occurs is approximately given by if" 0.06/n volt (V) for a two-color indicator whose forms are equally intensely colored. Since hydrogen ions are involved in the redox equilibria of many indicators, it must be recognized that the color change interval of such an indicator will vary with pH. 

The most important class of redox indicators, however, are substances that do not participate in the redox titration, but whose oxidized and reduced forms differ in color. When added to a solution containing the analyte, the indicator imparts a color that depends on the solution s electrochemical potential. Since the indicator changes color in response to the electrochemical potential, and not to the presence or absence of a specific species, these compounds are called general redox indicators. 

A partial list of general redox indicators is shown in Table 9.18. Examples of appropriate and inappropriate indicators for the titration of Fe + with Ce + are shown in Figure 9.37. 

AH classes of azine dyes are vattable, ie, they are reduced to "colorless" forms, then oxidized back to the dye (2). They therefore offer good fastness to oxidation (3). Because of this property, many find uses as redox indicators in titrations. 

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. 

For the titration of colourless or slightly coloured solutions, the use of an indicator is unnecessary, since as little as 0.01 mL of 0.02 M potassium permanganate imparts a pale-pink colour to 100 mL of water. The intensity of the colour in dilute solutions may be enhanced, if desired, by the addition of a redox indicator (such as sodium diphenylamine sulphonate, AT-phenylanthranilic acid, or ferroin) just before the end point of the reaction this is usually not required, but is advantageous if more dilute solutions of permanganate are used. 

The green colour due to the Cr3+ ions formed by the reduction of potassium dichromate makes it impossible to ascertain the end-point of a dichromate titration by simple visual inspection of the solution and so a redox indicator must be employed which gives a strong and unmistakable colour change this procedure has rendered obsolete the external indicator method which was formerly widely used. Suitable indicators for use with dichromate titrations include AT-phenylanthranilic acid (0.1 per cent solution in 0.005M NaOH) and sodium diphenylamine sulphonate (0.2 per cent aqueous solution) the latter must be used in presence of phosphoric) V) acid. 

The introduction of reversible redox indicators for the determination of arsenic(III) and antimony(III) has considerably simplified the procedure those at present available include 1-naphthoflavone, and p-ethoxychrysoidine. The addition of a little tartaric acid or potassium sodium tartrate is recommended when antimony(III) is titrated with bromate in the presence of the reversible... 

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. 

A weighed amount of sample is dissolved in a mixture of propanone and ethanoic acid and titrated potentiometrically with standard lead nitrate solution, using glass and platinum electrodes in combination with a ferro-ferricyanide redox indicator system consisting of 1 mg lead ferrocyanide and 0.5 ml 10% potassium ferricyanide solution. The endpoint of the titration is located by graphical extrapolation of two branches of the titration plot. A standard solution of sodium sulfate is titrated in the same way and the sodium sulfate content is calculated from the amounts of titrant used for sample and standard. (d) Water. Two methods are currently available for the determination of water. 

To leam why the ideal choice of a redox indicator is one which has a standard redox potential as close as possible to the end point of the titration system. 

The end point of a redox titration can be detected by use of a redox indicator, that is, a redox couple (0,R) in which the oxidized and reduced forms of the couple have distinctly different UV-visible spectra and hence, present different colours to the naked eye ... 

Suppose we add a small quantity of redox indicator to a relatively large volume of our redox titration mixture, e.g. aqueous ferrous and ferric ions. In addition, suppose that the starulard electrode potential of the indicator couple (we will call it E ) is greater than that of the ferrous-ferric... 

The larger the difference in standard potential between titrant and analyte, the greater the break in the titration curve at the equivalence point. A redox titration is usually feasible if the difference between analyte and titrant is a 0.2 V. However, the end point of such a titration is not very sharp and is best detected potentiometrically. If the difference in formal potentials is a 0.4 V. then a redox indicator usually gives a satisfactory end point. 

Numerous analytical procedures are based on redox titrations involving iodine. Starch8 is the indicator of choice for these procedures because it forms an intense blue complex with iodine. Starch is not a redox indicator it responds specifically to the presence of I2, not to a... 

Ce4+ is yellow and Ce3+ is colorless, but the color change is not distinct enough for cerium to be its own indicator. Ferroin and other substituted phenanthroline redox indicators (Table 16-2) are well suited to titrations with Ce4+. 

B. Would indigo tetrasulfonate be a suitable redox indicator for the titration of Fe(CN)g with Tl3+ in 1 M HC1 Hint The potential at the equivalence point must be between the potentials for each redox couple.)... 

V(phen)32+ has been used as a redox indicator in the titration of FeCl3 by Cr11 and E0 for Ru(bipy)32+ is very high (1.33 V at pH = 0) and the colour change (colourless to yellow) is well marked in the titration of oxalate in 2 M perchloric acid with ceric nitrate. 

Because of the clinical significance of vitamin C, it is essential to In-able to detect and quantify its presence in various biological materials. Ana lytical methods have been developed to determine the amount of ascorbic acid in foods and in biological fluids such as blood and urine. Ascorbic acid may be assayed by titration with iodine, reaction with 2,4-dinitrophenylhy-drazine, or titration with a redox indicator, 2,6-dichlorophenolindophenol (DCIP) in acid solution. The latter method will be used in this experiment because it is reasonably accurate, rapid, and convenient and can be applied to many different types of samples. 

In this experiment several samples will be prepared, and the amount of vitamin C in each will be determined. Each sample will be titrated with standardized redox indicator, DCIP. The vitamin C content will be calculated in terms of milligrams per milliliter (mg/mL) of sample. 

The sodium salt of these acids may be used to prepare aqueous solutions of indicators. Other examples of redox indicators include starch, potassium thiocyanate, methylene blue, and phenosafranine. Some selected general indicators in redox titrations are listed in Table 1.6.3. The properties of starch as an indicator in iodometric titration are discussed in the following section. 

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... 

The electrical circuit consists of two electrodes a redox indicator electrode and a reference electrode that also passes current. A fixed potential difference is applied and the equivalence point is calculated from the intersection of the two straight lines that show the variation of current before and after the endpoint in a plot of current as a function of added titrant volume. The plots can have various forms, depending on whether the titrated species or titrant are or are not electroactive. Figure 14.1 shows the four possible cases. Sometimes the potential difference applied is less than that necessary to reach the mass-transport-limited current, but sufficient to give good results. 

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... 

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