Indicator Method

Indicator Method for pH determination, colorimetric measurement of the active reaction of a solution according to the color of the indicator added to it, which is a weak acid or weak base (see Indicators). The greater the dissociation constant of an acid indicator, the higher the concentration of H-ions required to suppress its dissociation and cause a change in color; the same applies to alkaline indicators and OH-ions. Therefore, different indicators change their color at various pH values, which can deviate quite far from the neutral point in both acidic and alkaline directions. That pH value at which the degree of dissociation of the indicator is equal to one-half is called its transition point. At this point, it has an intermediate color between those it takes in a sharply acidic or alkaline medium, i.e., in the undissociated or in the completely dissociated state. On both sides of the transition point, the color of the indicator changes more or less sharply. This so-called transition zone, or indicator range zone. It is presented in the following table for several of the most commonly used indicators. pH of solution Methyl Orange Methyl-red o-nitrophenol Neutral-red α-naphtholphthalein Phenolphthalein Color of indicator 2 . . red - - - 3 . . orange-red violet-red - - 4 . . orange red - - 5 . . yellow orange colorless - - 6 . . - yellow greenish red - 7 . . - . yellow-green pink colorless - 8 . . - - orange greenish colorless 9 . . - - | yellow blue pink 10. . - - I - . red By examining a given solution with a series of indicators having different, sequentially differing transition ranges, its pH can be approximately determined. For example, if methyl orange takes on a yellow color in the solution being tested, its pH must be not lower than 5; the absence of reddening of phenolphthalein will show that the pH does not exceed 8 (more accurately-8.2); the greenish color of α-naphtholphthalein will lead to attributing to the pH approximately this last value. Buffer method. However, after such a preliminary determination, having found a suitable indicator whose transition range contains the desired pH, for its precise measurement it is necessary to perform a colorimetric comparison of the color of the indicator in this solution with its color in solutions of previously known reaction. For this purpose, buffer mixtures (see Buffer properties) are used, with the help of which a series of standard solutions with uniformly changing active reaction is prepared. Their pH can be easily calculated from the corresponding formulas or found in tables published by various authors. Moreover, in careful research to avoid inaccuracies possible in the preparation of buffer solutions, the latter should be checked using the hydrogen chain (see Gas chain), which is the basic and most accurate method for measuring pH. The accuracy of colorimetric measurement usually does not exceed 0.05 pH. The determination is somewhat complicated in the case of colored or turbid liquids, the natural color of which adds to the color of the indicator. In this case, in order to perform a colorimetric comparison of the color of the indicator in the test and in the standard solution, an equal layer of the test liquid (without indicator) is placed behind the latter as a "screen" on the path of the light rays. It acts as a light filter and affects the color of the passing rays in exactly the same way as if it were in the same solution with the indicator. Thanks to this ingenious device, proposed by Walpole, the influence of the natural pigment on the result of colorimetric comparison is eliminated. In Scheme I, the arrangement of the solutions is shown, and in the figure - a simple apparatus, the so-called comparator, commonly used for measurement (pure water is placed behind the test solution so that the absorption of light by water is also the same in both cases). Non-buffer method. A certain modification and simplification of the usual indicator method for measuring pH is the non-buffer method, proposed by Gillespie. The change in color of the indicator depends on the change in its degree of dissociation. The latter is expressed by a dissociation curve, and each pH corresponds to a certain point on the curve of a given indicator and a certain color. Instead of measuring pH with the help of standard buffer solutions, it can be found by directly determining the degree of dissociation of the indicator. For this purpose, let us place two solutions of the indicator on the path of a light ray - an acidic and an alkaline one. In one of them, all indicator molecules will be in the dissociated, in the other - in the undissociated state. The light ray, passing through both solutions, will obviously acquire the same color as if the dissociated and undissociated molecules of the indicator were together in one solution. Let us prepare, for example, with this indicator the following nine pairs of solutions: 1) 1 drop of indicator in an alkaline solution and 9 drops in an acidic one, 2) 2 drops in alkaline and 8 in acidic..., 9) 9 drops in alkaline and 1 drop in acidic. By colorimetric comparison, we will now find at what distribution of the indicator between the acidic and alkaline liquid a color identical to the color of 10 drops of indicator in the test solution will be obtained. The arrangement of the solutions is shown in Scheme II; to obtain it, Walpole's comparator, described above, is used. The obtained ratio directly gives the degree of dissociation of the indicator, and consequently the pH of the solution.

I trophenols; in a solution more acidic than their transition zone, they are colorless. Therefore Water Tested liquid + 10 drops of indicator Acidic solution + (10-p) drops of indicator Alkaline solution + p drops of indicator Scheme II. If the indicator represents a weak acid, then in the alkaline solution the salt formed from it is completely dissociated. The distribution of the indicator between the alkaline and acidic solution, equal for example 3:7, corresponds to a degree of dissociation of 0.3. With a ratio of 5:5 the degree of dissociation equals 0.5, and the pH equals the transition point of the given indicator. For several indicators, the pH values determined in this way are given in the table (according to measurements by Hayspie, at 25-30°). Measurement is further simplified when using one-color indicators, of which only one form is colored, while the other is completely colorless. Such are for example phenolphthalein and various nitro- and di-Brom-phenol-blue Methyl-1. rot Brom-:cresol-purple Brom-thymol-blue 3.1 3.5 3.7 3.9 4.1 4.3 4.5 4.7 5.0 4.05 4.4 4.6 4.8 5.0 5.4 5.6 5.95 5.3 5.7 5.9 6.1 6.3 6.5 6.7 6.9 7.2 6.15 6.5 6.7 6.9 7.1 7.3 7.5 7.7 8.05 Phenol-Cre-Thymol-rot rot blue 6.75 7.15 7.85 7D 7.5 8.2 7.3 7.7 8.4 7.5 7.9 8.6 8.1 8.8 7.9 9.0 8.1 8.5 9.2 8.3 8.7 9.4 8.65 9.05 9.75 the solutions in scheme II can be discarded. The degree of dissociation of the indicator in this case is simply equal to the ratio of its amount in the alkaline solution to its content in the tested solution (of course provided that the indicator represents a weak monobasic acid). In view of the fact that nitrophenols and dinitrophenols are distinguished among one-color indicators by considerable stability, Michaelis proposed to prepare from them according to the method described here constant series of solutions (containing various amounts of a given indicator in an alkaline medium), comparison with which directly gives the desired pH. The bufferless method and its simplifications described above are especially convenient for rapid approximate measurement of the reaction. Colorimetry of tissues. For measuring pH, usually 5-10 cm³ of the tested solution is taken. By reducing the size of the comparator, this number can be lowered to 1 cm³. To measure even smaller amounts, for example one drop of liquid, it is compared with a series of drops of standard mixtures, adding to them if possible the same drops of indicator. Differences in color are more noticeable if observed against a clean white background, for example on a white porcelain plate. Graff utilized this method for measuring the reaction of tissues, small pieces of which he quickly rubbed to release the tissue juice. However, such crushing of tissues causes significant postmortem mixing of their reaction (in the acidic direction) and does not give a correct picture of the pH of living tissue (coarser differences between the pH of various tissues can still be captured by this method). The error can be significantly reduced if the tissue to be crushed is first strongly cooled and frozen (with liquid air). In this case, however, Graff's method, without becoming fully reliable and flawless, loses its main advantage—simplicity and ease of execution. The danger of postmortem change in intracellular reaction can only be avoided by measuring the pH of living tissue. This became possible with the development of microsurgical methods (see Microsurgery), which allow introducing any (non-toxic to it!) indicators into a living cell. Schmidtmann, by pricking the cell with a micro-needle, introduced a grain of dry, water-soluble indicator attached to it. The color taken on by the cell juice adjacent to the indicator and partially dissolving it was compared with the color imparted by such a grain of indicator to various buffer solutions. Other researchers (Chambers and his colleagues, D. and J. Needham, etc.) injected a droplet of indicator solution of known concentration into the cell. By using a series of indicators having bright and sharply changing colors and sufficiently close transition points, it is possible with sufficient approximation to determine the pH of the cell sap or even individual parts of it. Errors of the I. m. When using indicators, it is necessary to take into account a number of conditions that can more or significantly distort the measurement results. This is first of all—the concentration of the indicator itself. Indicators are not neutral substances—most often they are acids. In the case of low buffer capacity of the solution, their addition can significantly shift the initial reaction. This effect will naturally be the stronger the higher the concentration of the indicator. Therefore, when studying liquids with very small buffer capacity (for example, pure water), one should limit oneself to the smallest possible amounts of indicator. The temperature of the solution also affects the transition zone of the indicator. However, when using the buffer method, this effect can be eliminated, since it is the same in the tested liquid and in the solution used for comparison. It is only necessary to know the pH of the corresponding buffer solutions at a given temperature. In the bufferless method, the change in transition point with temperature—if the change is sufficiently large—cannot be neglected. The so-called salt error of indicators plays a much greater role. All indicators to a greater or lesser degree change their degree of dissociation and transition point in the presence of salts. Indicators in which this effect is strong are completely unsuitable for studying solutions rich in salts, for example sea water. But for all others, in this case, a corresponding correction must be introduced, the value of which at the present time is known for many indicators. The protein error is more difficult to take into account—and therefore much more dangerous. In the presence of various protein substances and products of their decomposition, indicators can give completely incorrect readings. Some of them turn out to be absolutely unsuitable under these conditions, others—inaccurate. Since the combination of protein substances with indicators, causing the protein error, occurs gradually, the indicator immediately after addition may give for several moments a more correct reading, after which its color quickly changes. Thus, only in the absence of proteins and products of their decomposition, as well as large amounts of salts, colorimetric methods give sufficiently reliable results. In the presence of these substances, only the application of the hydrogen chain allows in each case to establish the suitability of one or another indicator and the magnitude of the corresponding correction.