Acid-base indicators in titrations

The intensity and colour of the fluorescence of many substances depend upon the pH of the solution indeed, some substances are so sensitive to pH that they can be used as pH indicators. These are termed fluorescent or luminescent indicators. Those substances which fluoresce in ultraviolet light and change in colour or have their fluorescence quenched with change in pH can be used as fluorescent indicators in acid-base titrations. The merit of such indicators is that they can be employed in the titration of coloured (and sometimes of intensely coloured) solutions in which the colour changes of the usual indicators would... 

Historically, pH sensitive dyes have been extensively used as indicators in acid-base titrations and in simple spot test papers, even leading to a common phrase in our everyday language, when people or topics are described as having passed the litmus test . The use of complexometric titrations for metal ions was a later but widely... 

Because of the striking color change between the bases (1) and monocations (4), it has been suggested that they could be useful indicators in acid-base titrations. A number of 2,4-disubstituted derivatives have been suggested for use over the pH range 5-9 (80MI2), and certain 3-aryl-... 

Salt formation generally favors the quinoid form and these salt solns are usually brightly Qolored. Presence of strong acids favors the colorless phenol form. Because of these color changes, nitrophenols have been used as indicators in acid-base titrations... 

PROBLEM 11.22 Methyl orange is an azo dye used as an indicator in acid-base titrations. (It is yellow-orange above pH 4.5 and red below pH 3.) Show how it can be synthesized from p-aminobenzenesulfonic acid (sulfanilic acid) and N,N-dimethylaniline. 

The following dye is used as indicators in acid-base titrations Rationalize the color changes. 

Many chemiluminescent compounds emit light only in neutral or, better in alkaline solution, e.g. luminol and lucigenin. Thus these can be used as neutralization indicators in acid-base titrations. Used in an end-point titration, light can be detected from strongly coloured or turbid or opaque solutions where colorimetric titrations are difficult. For example in the determination of the acidity of milk, of red wine, or mustard [4] or of dark coloured fats and oils [5,6]. It seems that luminol-fluorescein mixtures are better than luminol alone, because a hemin catalyst is not required here [7]. The system luminol - hemin leads to an irreversible destruction of hemin and is, therefore, not reversible with respect to the catalyst - in contrast to the luminol-fluorescein system [7]. 

Titrimetric (volumetric) factors for acids and bases are given in Table 11.28. Suitable indicators for acid-base titrations may be found in Tables 8.23 and 8.24. 

The utility of acid-base titrimetry improved when NaOH was first introduced as a strong base titrant in 1846. In addition, progress in synthesizing organic dyes led to the development of many new indicators. Phenolphthalein was first synthesized by Bayer in 1871 and used as a visual indicator for acid-base titrations in 1877. Other indicators, such as methyl orange, soon followed. Despite the increasing availability of indicators, the absence of a theory of acid-base reactivity made selecting a proper indicator difficult. 

In acid-base titrations the end point is generally detected by a pH-sensitive indicator. In the EDTA titration a metal ion-sensitive indicator (abbreviated, to metal indicator or metal-ion indicator) is often employed to detect changes of pM. Such indicators (which contain types of chelate groupings and generally possess resonance systems typical of dyestuffs) form complexes with specific metal ions, which differ in colour from the free indicator and produce a sudden colour change at the equivalence point. The end point of the titration can also be evaluated by other methods including potentiometric, amperometric, and spectrophotometric techniques. 

For end-point detection, any method usual in acid-base titration can be used with electrometric indication the precautions for protection against the... 

Figure 14 Some examples of endpoint determination in titrations using chemiluminescent indicators. (A) Acid-base titration the endpoint is detected by the emission of light (B) complexometric titration the endpoint is detected by disappearance of light. M, metal acting as a catalyst X, excited state from the CL precursor acting as indicator.

Observations that the presence of protein affects the colour change of some indicators used in acid-base titrations led to the development of methods for the quantitation of proteins based on these altered absorption characteristics of such dyes. As the presence of protein alters the colour produced by these indicators when measuring pH, so in the quantitation of proteins using dye-binding methods the control of pH is vital. 

Indicators are used in acid-base titrations as they change colour at i reaction. Indicators are usually weak acids in which the colour of the acid is different from that of its conjugate base. If we represent the Indicator as a weak acid of formula HIn, and say it has a red colour, and its conjugate base as In", and say it has a blue colour, its dissociation in water can be represented as Hln(aq) + HjO(l) H30" (aq) + In"(aq)... 

For the titration of a strong base with a weak acid, the equivalence point is reached when the pH is greater than 7. The half equivalence point is when half of the total amount of base needed to neutralize the acid has been added. It is at this point that the pH = pK of the weak acid. In acid-base titrations, a suitable acid-base indicator is used to detect the endpoint from the change of colour of the indicator used. An acid-base indicator is a weak acid or a weak base. The following table contains the names and the pH range of some commonly used acid-base indicators. 

As in acid-base titrations, indicators and electrodes are commonly used to find the end point of a redox titration. 

An important question is whether we can use any indicator electrode. A redox electrode, i.e. inert in the range of potential where measurements are being done, is a possibility, especially for redox titrations. In other cases, the use of electrodes selective to the ion being titrated is better, such as pH electrodes in acid-base titrations. The method of analysis of the data obtained is, naturally, the same in all cases and independent of electrode material. 

Figure 22.7 Examples of indicators used in acid-base titrations thymol blue and phenolphthalein...

More than brief discussion of the numerous ways in which end points can be taken other than by visual methods is beyond our scope. For example, end-point techniques may involve photometry, potentiometry, amperometry, conductometry, and thermal methods. In principle, many physical properties can be used to follow the course of a titration in acid-base titrations, use of the pH meter is common. In terms of speed and cost, visual indicators are usually preferred to instrumental methods when they give adequate precision and accuracy for the purposes at hand. Selected instrumental methods may be used when a suitable indicator is not available, when higher accuracy under unfavorable equilibrium conditions is required, or for the routine analysis of large numbers of samples. 

It is important to know the dissociation constant of an indicator in order to use it properly in acid-base titrations. Spectrophotometry can be used to measure the concentration of these intensely colored species in acidic versus basic solutions, and from these data the equilibrium between the acidic and basic forms can be calculated. In one such study on the indicator wj-nitrophenol, a 6.36 X 10 M solution was examined by spectrophotometry at 390 nm and 25°C in the following experiments. In highly acidic solution, where essentially all the indicator was in the form HIn, the absorbance was 0.142. In highly basic solution, where essentially all of the indicator was in the form In , the absorbance was 0.943. In a further series of experiments, the pH was adjusted using a buffer solution of ionic strength I, and absorbance was measured at each pH value. The following results were obtained ... 

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. 

Although indicators are still widely used in acid/base titrations, the glass pH electrode and pH meter allow the direct measurement of pH as a function of titrant volume. The glass pH electrode is discussed in detail in Chapter 21. The titration curve for the titration of 50.00 mL of 0.1000 M weak acid = 1.0 X 10 ) with 0.1000 M NaOH is shown in Figure 14F-4a. The end point can be located in several ways from the pH versus volume data. 

Indicator For acid-base titrations, an organic compound that exhibits its different colors in solutions of different acidities used to determine the point at which the reaction between two solutes is complete. 

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). 

The procednre for the titration is shown in Fignre 4.21. First, a known amount of KHP is transferred to an Erlenmeyer flask and some distilled water is added to make up a solution. Next, NaOH solntion is carefnlly added to the KHP solution from a buret until we reach the equivalence point, that is, the point at which the acid has completely reacted with or been neutralized by the base. The eqnivalence point is usually signaled by a sharp change in the color of an indicator in the acid solution. In acid-base titrations, indicators are substances that have distinctly different colors in acidic and basic media. One commonly nsed indicator is phenolphthalein, which is colorless in acidic and nentral solntions bnt reddish pink in basic solntions. At the equivalence point, all the KHP present has been nentralized by the added NaOH and the solution is still colorless. However, if we add jnst one more drop of NaOH solntion from the buret, the solution will immediately tnm pink becanse the solntion is now basic. Example 4.9 il-Instrates snch a titration. 

Many acid-base indicators are plant pigments. For example, by boiling chopped red cabbage in water we can extract pigments that exhibit many different colors at various pHs (Figure 16.6). Table 16.1 lists a number of indicators commonly used in acid-base titrations. The choice of a particular indicator depends on the strength of the acid and base to be titrated. Example 16.7 illustrates this point. 

In acid-base titrations, close to the equivalence point, there is a rapid change in pH which may be assessed by adding a suitable indicator. The acid-base indicator must change its colour over a range of pH which is as near as possible to that of equivalence point. 

As we mentioned, for a titration to be successful there must be some indication of when it has finished. Chemical indicators are a common way of determining whether a titration is complete. In acid-base titrations, chemical indicators are typically small molecules whose colors change depending on the pH of the solution (or, in other words, on their protonation state). In other cases, the indicator may change color based on whether it is coordinated to another species in a solution. It is also possible to use a device like a pH-sensitive electrode to determine when the titration has ended this doesn t rely on you (or another person) being able to notice a visible change in the color of the solution. Such a device can be more accurate, since it doesn t rely on a qualitative interpretation of a color change. 

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