Indicator color change table

O Soiution after severai small drops oi NaOH are added. The solution is now basic and the indicator is biue. The equivalence point occurs at pH = 7, which is the middie of the bromthymoi blue indicator color change (Table 19-4). There is a very iarge change in pH (3.76 to 8.65 for this solution) right around the equivaience point for strong acid/strong base titrations. 

The indicator method is especially convenient when the pH of a weU-buffered colorless solution must be measured at room temperature with an accuracy no greater than 0.5 pH unit. Under optimum conditions an accuracy of 0.2 pH unit is obtainable. A Hst of representative acid—base indicators is given in Table 2 with the corresponding transformation ranges. A more complete listing, including the theory of the indicator color change and of the salt effect, is also available (1). 

For those cases where a reaction occurred, write a complete and balanced equation. Indicate precipitates, gases, and color changes. Table 8.2 lists some insoluble salts. For those cases where no reaction took place, write No Reaction. ... 

To each of three different 13 x 100-mm test tubes, add 2 mL of distilled water and 1 drop of chlorophenol red indicator (color change from pH 5.2 to 6.8). Add 1 drop of 6 M aqueous eunmonia, NH3, solution to the first test tube and 1 drop of 3 M hydrochloric acid, HCl, solution to the third test tube so that the first and third test tubes will then serve as reference colors for the indicator. Bubble gas from your CO2 generator through the solution in the second test tube for a few minutes, (If CO2 generation is no longer strong, perform the bracketed procedure two paragraphs above.) Describe in TABLE 34.IB the colors of the three solutions. 

At pH = p a = 4.7, there is a 1 1 mixture of yellow and blue species, which appears green. As a crude rule of thumb, we predict that the solution will appear yellow when [Y]/[B ] 10/1 and blue when [B]/[Y] 10/1. (The symbol means is approximately equal to or greater than. ) From Equation 9-4, we predict that the solution will be yellow when pH p nin 1 (= 3.7) and blue when pH p nin + 1 (= 5.7). By comparison, Table 9-3 lists bromocresol green as yellow below pH 3.8 and blue above pH 5.4. Between pH 3.8 and 5.4, various shades of green are seen. Demonstration 9-2 illustrates indicator color changes and Box 9-2 describes an everyday application of 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. 

In Table 8.26, E° represents the redox potential at which the color change of the indicator would normally be perceived in a solution containing approximately 1A7H+. Lor a one-color indicator this is the potential at which the concentration of the colored form is just large enough to impart a visible color to the solution and depends on the total concentration of indicator added to the solution. If it is the reduced form of the indicator that is colorless, the potential at which the first visible color... 

A list of several common acid-base indicators, along with their piQs, color changes, and pH ranges, is provided in the top portion of Table 9.4. In some cases. 

Comments on the thermal nitration of enol silyl ethers with TNM. The strikingly similar color changes that accompany the photochemical and thermal nitration of various enol silyl ethers in Table 2 indicates that the preequilibrium [D, A] complex in equation (15) is common to both processes. Moreover, the formation of the same a-nitroketones from the thermal and photochemical nitrations suggests that intermediates leading to thermal nitration are similar to those derived from photochemical nitration. Accordingly, the differences in the qualitative rates of thermal nitrations are best reconciled on the basis of the donor strengths of various ESEs toward TNM as a weak oxidant in the rate-limiting dissociative thermal electron transfer (kET), as described in Scheme 4.40... 

An appropriate indicator would change color in the pH range, 7-10. From Table 19-4, phenolphthalein is the best indicator for this titration. 

The pH of a solution may be estimated from color indicators that change hue with pH, like litmus or phenolphthalein papers see Table 10-2. Where precise values are required, an electrical pH meter is utilized. 

Due to the low solubility of the concave pyridines 3 in diethyl ether, the corresponding pyridine buffers could not be compared with the experiments of Table 4. But when the protonations were carried out in other solvents, no influences of the acids (including the acids of Table 4) on the regioselectivity could be found. The exchange of diethyl ether by other solvents caused a color change of the allyl anion solution which indicated different structures for the anions in diethyl ether and in other, more polar solvents [44],... 

We will do one sample calculation for each region. Complete results are shown in Table 11-1 and Figure 11-1. As a reminder, the equivalence point occurs when the added titrant is exactly enough for stoichiometric reaction with the analyte. The equivalence point is the ideal result we seek in a titration. What we actually measure is the end point, which is marked by a sudden physical change, such as indicator color or an electrode potential. 

F. Select indicators from Table 11-4 that would be useful for the titrations in Figures 11-1 and 11-2 and the pKa= 8 curve in Figure 11-3. Select a different indicator for each titration and state what color change you would use as the end point. 

Metal ion indicators (Table 12-3) are compounds whose color changes when they bind to a metal ion. Useful indicators must bind metal less strongly than EDTA does. 

For ferroin, with E° = 1.147 V (Table 16-2), we expect the color change to occur in the approximate range 1.088 V to 1.206 V with respect to the standard hydrogen electrode. With a saturated calomel reference electrode, the indicator transition range will be... 

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+. 

Select indicators from Table 16-2 that would be suitable for finding the end point in Figure 16-3. What color changes would be observed ... 

For the pH ranges over which common indicators change color, see Table 11.2. 

A mixture of 1.42 g (0.01 mol) phosphorus pentoxide and 2.5 g of chromatography grade silica gel was placed in a flask and stirred for 30 min. A mixture of equimolar amounts (5 mmol) of the carboxylic acid and phenol was added. Usually an immediate color change was observed. After stirring for 6 h at the temperature indicated in the Table, methylene chloride (50 mL) was added. The mixture was stirred for 1 min and then filtered. The spent reagent was washed twice with methylene chloride (10 mL). The combined organics were washed with aqueous NaOH solution, water and dried over sodium sulfate, and the solvent was removed under reduced pressure. 

Table 5.17 lists ranges and colors of common pH indicators, and Figure 5.1 shows a graphical overview of the pH ranges and color changes. 

The difference in color change can clearly be seen by comparing the changes described in Tables 5.18 and 5.18 with those of Table 5.17 the changes of the mixtures are much more pronounced than those of the individual indicators. 

Fig. 4.2. The Miller-Urey apparatus for abiotic synthesis of biochemicals from primordial gases is shown. Before each experiment the system was thoroughly evacuated, flushed with interstellar-type gases, and sealed. Water is brought to a boil and vapors rise through an electric discharge chamber and are re-condensed and led back into the boiling water reservoir. It took only a few weeks to produce a color change in the water which indicated an accumulation of organic compounds shown in Table 4.1. On the young earth, of course, this experiment would have been carried on for a few million years.

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