HYDROGEN ION INDICATORS 



159 



HYDROGEN ION INDICATORS 



indicator is very important. That 

 chosen should be one which has its 

 sharpest color transition near the true 

 equivalence point. For example, when 

 titrating the weak base ammonium 

 hydroxide with hydrochloric acid, the 

 resulting ammonium chloride is an 

 acidic salt, and the indicator selected 

 should be one that shifts below neu- 

 trality, methyl red for example. 

 Again, when titrating a weak acid such 

 as lactic acid with sodium hydroxide, 

 the salt formed is alkaline, and the 

 indicator of choice should shift above 

 neutrality, bromthymol blue, for ex- 

 ample. In titrating strong acids and 

 strong bases against each other, the 

 selection of indicator is not critical, as 

 here the pH change near the end point 

 is very great for only a small increment 

 of the added reagent. One should 

 never titrate weak acids and weak bases 

 together, for the results cannot be 

 accurate. 



There are several possible sources of 

 error in the use of hydrogen ion indi- 

 cators. Some are present at all times 

 and others especially so in biological 

 fluids. An occasional indicator, with 

 two groups sensitive to acid or alkali, 

 has two ranges of color transition at 

 perhaps widely separately pH values. 

 For example, th3^mol blue shifts from 

 red to yellow at a low range (1.2-2.8) 

 and from yellow to blue at a much higher 

 one (8.0-9.6) . A very few indicators ex- 

 hibit "dichroism" (or "dichromatism") 

 in which the color varies with the depth 

 and concentration of the solution. 

 Bromcresol purple and bromphenol blue 

 are examples. 



In attempting to determine the pH 

 of a very dilute solution, a false result 

 may be obtained by the use of indicators 

 as they are themselves acids or bases. 

 (Recall, for example, that water in con- 

 tact with the carbon dioxide of the 

 normal atmosphere has a pH of about 

 5.7.) For such cases, a very small 

 amount of buffer should be present to 

 offset this effect. 



Indicators are intended for aqueous 

 sj^stems, and the presence of other 

 solvents such as alcohol decreases the 

 dissociation constant. Acidic indi- 

 cators then become more sensitive to 

 hydrogen ions, and basic ones less sensi- 

 tive. A control solution of the same 

 solvent composition may be used for 

 comparison, however. 



Temperature errors are slight over 

 the usual ranges. 



Indicators may be altered or de- 

 stroyed by the presence of oxidizing and 

 reducing agents, and thej^ niay unite 

 with heavy metal ions. Fortunately, 



these are negligible considerations in 

 biological fluids, but greater potential 

 errors exist. 



Proteins and their hydrolysis prod- 

 ucts are usually amphoteric and may 

 combine with the indicator. Congo 

 red, for example, is almost worthless in 

 protein solutions. 



The presence of much neutral salt 

 will affect the color of indicator solu- 

 tions, partly by influencing the light 

 absorption and partly by shifting the 

 ratio between dissociated and non- 

 dissociated forms of the indicator. 



The use of mixed indicators deserves 

 greater attention than it has yet re- 

 ceived. As employed for titrations, 

 they are of two general types. One 

 sort consists of two acid-base indicators 

 which have color transitions in op- 

 posite directions, resulting in a very 

 sharp change at a narrow pH zone. 

 The other kind utilizes for contrast 

 color a dye which itself is not influenced 

 by hydrogen ion concentration. The 

 composite color change resulting is 

 usually much sharper than that of the 

 indicator alone. In recent years, there 

 have appeared "universal indicators" 

 consisting of a mixture of half a dozen 

 compounds with a "spectrum" of colors 

 which may vary over the entire pH 

 range. Such indicators are not very 

 accurate and should be used only as a 

 first rough test on an unknown solu- 

 tion. These mixtures are more com- 

 mon as test papers. 



It is not surprising that hj^drogen ion 

 indicators have been employed as a 

 sort of vital stain to determine the re- 

 action of various living components. 



In 1893 Ehrlich injected neutral red 

 in an attempt to determine the reaction 

 about phagocj'tosed granules. Since 

 then, other workers have applied other 

 dyes, striving to estimate the approxi- 

 mate pH of tissues, of the fluids bathing 

 them, and even of individual cells. 

 Alizarin red and litmus have been much 

 used, the later especially with lower 

 organisms. Thus, Steiglitz applied all 

 three dyes mentioned above to estimate 

 the reaction of living kidney (E. J., 

 Arch. Int. Med., 1924, 33, 483-496) and 

 confi.rmed the contention that alkaline 

 urine can be formed by an acidic cortex. 

 Harvey and Benslev (B. C. H. and R. 

 R., Biol. Bull., 1912, 23, 225-249) used 

 pH indicators to indicate that gastric 

 fluid does not arise directly within the 

 cells of the mucosa. Margaria (R., J. 

 Physiol., 1934, 82, 496-497) injected 

 bromcresol purple and bromphenol blue, 

 and claimed to have measured pH 

 changes upon stretching a muscle. 

 Orr (J. W., J. Path. & Bact., 1937, 44, 



