Indicators, acid-base errors

The acid-base error of indicators. Isohydric indicators. The measurement of pH in unbuffered or slightly buffered solutions. 

It has been shown that for most acid-base titrations the inflection point, which corresponds to the greatest slope in the titration curve, very nearly coincides with the equivalence point. The inflection point actually precedes the equivalence point, with the error approaching 0.1% for weak acids or weak bases with dissociation constants smaller than 10 , or for very dilute solutions. Equivalence points determined in this fashion are indicated on the titration curves in figure 9.8. 

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 has been well known for a relatively long time that micellar, i.e. association colloidal, systems have a considerable effect on such indicator equilibria. Indeed, in the 1920 s and early 1930 s experiments were carried out in order to elucidate the so-called colloid or indicator error (Hartley, 1934 Hartley and Roe, 1940). In addition, the protein error was noted in investigations involving acid-base titrations in the presence of proteins (Sorensen, 1929 cf. Hartley, 1934). These errors are, of course, the consequence of micellization and the subsequent effects of micelles on equilibrium (34). The importance of many indicators in the dye, textile, and photographic industries, and the analytical utility of the shifts in indicator equilibria prompted much of the research in this area. 

Use an acid—base indicator, which marks the endpoint of a titration by changing color. Although the equivalence point of a titration, defined by the stoichiometry, is not necessarily the same as the endpoint (where the indicator changes color), careful selection of the indicator will ensure only negligible error. 

Titration Errors with Acid/Base Indicators... 

We find two types of titration errors in acid/base titrations. The first is a determinate error that occurs when the pH at which the indicator changes color differs from the pH at the equivalence point. This type of error can usually be minimized by choosing the indicator carefully or by making a blank correction. 

Litmus (azolitmin) is the classical indicator for acid-base titrations, although by now it has been supplanted by much better indicators. Neither litmus nor azolitmin should be used in colorimetric determinations of pH because of their salt and protein errors. Litmus is of value only in the form of indicator paper (cf. Chapter Eleven). 

Erroneous results are obtained also in the colorimetric determination of the pH of solutions of very weak acids (or bases). When an indicator acid is employed, the acid error of the indicator itself must be taken into account and when the indicator salt is used, allowance must be made for the following reaction between the salt and the acid present ... 

Whereas the acid or base error of indicators can be eliminated by proper experimental conditions, the influence of electrolytes can not be avoided experimentally. The effect of electrolytes can be traced in general to two causes ... 

The ratio/i /2 in (16) increases with larger salt concentrations much more rapidly than the corresponding/o /i ratio in (14). Hence the salt error of dibasic indicator acids is much greater than of monobasic acids (such as the nitrophenols). Indicator bases may be treated in analogous fashion. 

In cases where the protein error is not known in advance, it is advisable to determine the pH both with an indicator acid and an indicator base. The result may be considered reliable if both measurements agree. 

The reason for systematic titration errors is that the equivalence point is indicated too early or too late. This happens when the transition point of the indicator does not exactly match the pH of the equivalence point of the titration (systematic errors caused by wrongly calibrated pipettes or burettes will not be discussed here). The transition point of an indicator gives the experimental endpoint of the titration. Because the term endpoint can also be applied in the sense of theoretical endpoint = equivalence point we shall use here the term transition point to be clear. The same can happen in case of instrumental methods of indication when these methods do not identify the equivalence point correctly, but systematically deviate from it. Color indicators are themselves acid-base systems Hl/1 (HI + H2O 1 + HsO ), the p a value of which is usually denoted as the pA) value, and it normally falls in the range of 2-12. There are bichromic and monochromic indicators. For example, a bichromic indicator may be red as an acid and blue as a base, and a monochromic may be colorless as an acid and violet as a base. In the case of bichromic indicators, the color changes when Chi = cr, that is at the buffer point of the indicator. Of course, the color change does not abruptly occur there, but it is smeared out in an interval (the so-called transition interval of an indicator), roughly in the... 

T raditionally, titration curve calculations are described in terms of equations that are valid only for parts of the titration. Equations will be developed here that reliably describe the entire curve. This will be done first for acid-base titration curves. In following chapters, titration curves for other reaction systems (metal complexation, redox, precipitation) will be developed and characterized in a similar fashion. For all, graphical and algebraic means of locating the endpoints will be described, colorimetric indicators and how they function will be explained, and the application of these considerations to (1) calculation of titration errors, (2) buffo design and evaluation, (3) sharpness of titrations, and finally, (4) in Chapter 18, the use of titration curve data to the determination of equilibrium constants will be presented. 

Errors (2) and (3) are negligibly small in comparison to error (1) and consequently in the selection of a suitable indicator only the magnitude of the chemical error is of great importance. Thus the appropriate acid-base indicator must have its transition pH range within the equivalence region. 

Omissions as well as errors of fact and interpretation are inevitable in dealing with a vast subject such as acid-base indicators. 1 shall be glad to have my attention drawn to errors and to incorporate suggestions for improvanent when a revision becomes possible. 

It is suspected that an acid-base titrimetric method has a significant indicator error and thus tends to give results with a positive systematic error (i.e. positive bias). To test this an exactly 0.1 M solution of acid is used to titrate 25.00 ml of an exactly 0.1 M solution of alkali, with the following results (ml) ... 

Fig. 3.8 Phenolic acids extracted from wheat stubble, wheat straw from half buried litter bags, and wheat stubble/soybean (no-till) soil. Phenolic acids isolated and quantified were caffeic acid (CAF), ferulic acid (FER), p-coumaric acid (PCO), p-hydroxybenzoic acid (POH), sinapic acid (SIN), syringic acid (SYR), and vanillic acid (VAN). Becausep-coumaric acid was so high in comparison to other phenolic acids in wheat residues, data are presented twice, once with p-coumaric acid (a) and once without p-coumaric acid (b). Because phenolic acids were so low in the soil they are also presented in (c). The absence of standard error bars for wheat straw and soil indicates that the error bars are too small to be visible. Figures based on data from Blum et al. (1991, 1992). Plenum Publishing Corporation, data used with permission of Springer Science and Business Media...

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