Color-change interval

5 Conditions for Use of Color Indicators 8.5.1 Color-Change Interval 

Henderson-Hasselbach s relation applied to the ionization of a pH indicator becomes pH=pK, + og [In-]/[InH]. 

The proton addition induces the disappearance of the basic form In in favor of the form InH. 

It is an experimental fact that the acidic form InH imposes its coloration when the two forms of concentration obey the rule 

In logarithmic terms, these inequalities imply that the color-change interval is such that 

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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. 

Answer (a) 0.989 V. (b) 0.67 V. (c) 0.32 V. (d) 0.043 V. (e) -0.209 V. Derive an expression for the minimum difference of formal potentials E — El for which a redox indicator would show an essentially complete change in color within the interval —0.1 to +0.1% of the end point. Assume that = equiv that the color-change interval is given by Equation (15-22). 

Thymolphthalein. The structure of this indicator is like that of phenolphthalein, with the exception that the phenol groups are replaced by two thymol groups. It melts at 253 . A convenient stock solution may contain 0.1% of the indicator in an 80% alcohol solution. Its color change interval lies between pH 9.3 and 10.5, the color changing from colorless to blue. 

Tetrdbromophenoltetraiodophthalein. Yellowish powder difficultly soluble in alcohol. Color-change interval from about 7.2 (colorless) to 9.0 (blue). 

The color change interval of phenol red is from pH 6.4 to 8.2, the color going from yellow to red. 

Heptamethoxytriphenylcarhinol Heptamethoxy red), C26H30O8, is prepared from 2,4-dimethoxyiodobenzene and the ethyl ester of 2,4,6-trimethoxybenzoic acid with a yield of 40%. Its melting point is 147°. The stock solutions should contain 0.1% of the indicator in alcohol. Appreciable time is required for the appearance of its color and in a colorimetric determination, solutions should be permitted to stand for a half hour before color comparisons are performed. The color change interval lies between pH 5.0 and pH 7.0. The indicator is of little use in titrations because of the time effect mentioned. H. Lund (l.c.) has studied the absorption spectrum of the indicator. 

Pentamethoxytriphenylcarhinol (Pentamethoxy red) was prepared by H. Lund from 2,4-dimethoxyiodobenzene and the methyl ester of o-methoxybenzoic acid, the yield being 40%. The compound melts at 146-147°. The alcoholic stock solution contains 0.1% of the indicator. This substance is very satisfactory, with a color change interval at pH 1.2-3.2 (red-violet to colorless). 

From the following experimental studies it is seen that the color change interval of any acid indicator may shift as much as two units on the pOH axis, and that the range of any alkaline indicator may be displaced two units on the pH axis. 

All of these experiments indicate tha< the color change interval of p-nitrophenol is displaced only slightly by heating. By extrapolating from the experiments of L. Michaelis and A. Gye-MANT, it can be estimated that the constant of p-nitrophenol at 100° is ten times as large as at room temperature. 

All of the preceding studies demonstrate the fact that the position of the color change interval of most indicators will be displaced appreciably as temperature increases. Only the sul-fonephthaleins and the phthaleins experience a negligible variation in their sensitivity towards hydrogen ions. The effect of warming solutions of the various indicators is summarized in the following table. 

Suitable indicator because equivalence point falls within the color-change interval... 

These relations show that the acid and base concentrations ratio depends on the pH through the indicator pKa value. The solution color depends on the pH value. Some indicators exhibit only one colored form they are called unicolor indicators. For example, this is the case of phenolphthalein. The theory concerning them is the same as that concerning bicolor indicators except for the influence of their concentration on the color-change interval (see Sect. 8.4). 

A first example is that of methyl orange (helianthin), whose bicolor indicator is the sodium salt of the 4-dimethylaminoazobenzen-4 -sulfonic acid whose color-change interval is 4.4 < pH < 6.2. The structures that participate in the equilibrium... 

We mention here just some indicators with their pH color-change intervals and their color changes by passing from the acidic to the basic form. 

Within this range, the indicator will appear to change from one color to the other. The color-change interval is hence... 

Between these limits, the color change will, of course, be gradual since it depends on the ratio of the two colored forms. According to some authors, the color-change interval is narrower (see Table 8.1). 

Trivial name Color-change interval Color change (acidic to basic form) ... 

Influence of the Indicator Concentration on the Color-Change Interval... 

The color-change interval of a bicolor pH indicator is independent of its total concentration. The assertion is correct only when the Henderson relation is legitimate. 

The color-change interval depends on temperature since its pKa value, which is present in Henderson s relation, depends on it (as do all the equilibrium constants). 

We must use indicators whose color-change interval is located in basic pH values. Their choice depends on thepK value of the acid. For example, for the titration of acetic acid in the above conditions, pH = 8.7 at the equivalence point. Using phe-nolphthalein, thymolphthalein, and thymol blue is satisfactory. For the second acid (pKa 7.00), pH = 9.9 at the equivalence point, only thymolphthalein is satisfactory. (Recall that in this second example, the color-change interval is narrower than in the first one—see the curve s shape). For very weak acids (pKa > 7), no simple indicator can be used. Only some mixtures of judiciously chosen indicators can be used. This is due to the fact that the pH change at the equivalence point is very weak. The very deep reason behind this fact is that the neutralization reaction can no longer be considered quantitative (see Chap. 10). 

From a practical standpoint, the neutralization indicators to be used must have their color-change interval located in the acid range. For example, for the above-mentioned titration of ammonia, methyl red and helianthin are suitable. This is also the case of bromophenol blue and bromocresol green. 

From a practical standpoint, we can infer from these considerations that there are two ways to titrate phosphoric acid We base it on either the first or second sharp endpoint. In the first case, only the first acidity is neutralized and the color-change interval is located near pH=4.8 Congo red is suitable. In the second case, the first two acidities are neutralized and the color-change interval is located near pH=9.8 thymolphthalein and phenolphthalein are suitable. The volumes at the equivalence point, which must be taken into account to calculate the titer, vary, of course, from simple to double according to the chosen strategy. 

The equivalence point is detected by the color change of phenol red. In strictly aqueous medium, the color change interval is 6.8 < pH< 8.4. Because in hydroalcoholic medium, the present species at the equivalence point is the base RCOO as in water, it means that the color change of phenol red occurs in a medium slightly more basic in ethanol than in water. Incidentally, we follow this kind of reasoning to compare the absolute acidity levels of different solvents. 

The disodium phosphate Na2HP04 can be titrated with a hydrochloric solution with a mixed indicator. The mixed indicator is a mixture of bromocresol green (color-change interval 3.8 color-change interval 4.2 < pH < 6.3) ... 

A strong acid is released during the reaction. An important point to note is that the formed oxime does not exhibit any basic character also, the initial solution was neutral. This is easily explained by the fact that hydroxylamine is a strong base and, hence, its conjugate acid a very weak one. It is interesting also to note that the reaction medium is nonaqueous for the most part, since it only contains 10% water (in mass). Water is added essentially in order to dissolve the hydroxylamine hydrochloride. The medium also contains pyridine in order to displace the oxime formation equilibrium toward the right by formation of the pyridinium ion. The hydrochloric acid released is titrated with a methanolic sodium hydroxide solution. The indicator chosen is bromophenol blue, whose color-change interval is located on the acidic side. Actually, it is the pyridinium cation that is titrated with the methanolic sodium hydroxide solution. The pyridinium cation results from the protonation of the pyridine by the released hydrochloric acid. 

The indication of the equivalence point is normally carried out with metal ion indicators (see Chaps. 26 and 29). We have already mentioned that not only do they exhibit the property to form chelates with metallic ions but also that they are acid-base indicators. Thus, their coloration depends on the presence of metal cations but also on the pH of the solution. We have already seen that their behavior may be described and explained by introducing their proper conditional constants. The point that remains to be investigated is their color change interval, which depends on several parameters. 

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