Indicators

INDICATORS, indicators of the state of a given system. In chemistry, during various operations, primarily in volumetric analysis, indicators are understood to be substances that make it possible to determine the end of a reaction. Thus, in diazotization, the so-called iodine-starch paper (paper impregnated with starch and potassium iodide) is used as an indicator; when a sufficient quantity of nitrous acid salt has been added for the diazotization reaction to reach completion, then upon further addition, the remaining free nitrous acid will oxidize the potassium iodide with the release of free iodine, which causes the indicator paper to turn blue. In volumetric analysis, in some cases, the solution used for titration itself serves as the indicator. Thus, when titrating with permanganate, as soon as a free excess of KMnO₄ appears in the reaction mixture, it imparts a pink color to the liquid. When titrating chlorides according to Mohr, a solution of K₂CrO₄ serves as the indicator: an excess of the silver solution used for precipitating the chlorides causes the formation of a red precipitate of AgCrO₄; when titrating according to Volhard, the indicator is ammonium thiocyanate, which gives a color with iron alum. In acidimetry and alkalimetry, indicators are substances of the nature of weak acids or bases that sharply change their color with the transition from the undissociated form to the ionized form, with the second differing in its structure from the first. The degree of ionization, and consequently the color of the indicator, depends on the active reaction of the medium; this has led to the widespread use of indicators in the colorimetric determination of the concentration of H-ions (C_H), as well as in acidimetry and alkalimetry.

To explain the change in the color of an indicator depending on the C_H of the solution, two theories have been proposed: 1) the ionic theory of Ostwald and 2) the chemical, or chromophoric, theory of Hantzsch. Ostwald believes that the undissociated molecules of the indicator cause one color of the solution, while the ions that appear during the formation of salt-like compounds of the indicator cause another. For example, phenolphthalein in an acidic solution is practically in an undissociated state and is colorless, while in an alkaline solution it breaks down into ions, giving a red color. Ostwald's point of view has met with serious objections, an example of which is the indication of the red color of phenolphthalein salt in a solid, and therefore undissociated, state. According to Hantzsch, the change in the color of an indicator is caused by the formation of compounds with a different structural formula than the original ones. It is not the breakdown of the molecule into ions, but intramolecular rearrangements that cause the change in color. During the further development of Hantzsch's theory, it became clear that the tautomeric forms of the indicator molecule formed with a change in the concentration of H⁺, in contrast to the original ones, are substances that dissociate well in solution. Thus, Hantzsch's theory approached the clarification of the mechanism of the formation of the ionized form of the indicator, deepening but not changing the basic tenets of Ostwald's theory. Between the number of ions and the molecules of the indicator that have not broken down into ions, according to the law of mass action, there exists the ratio: [H⁺] · [Ind^-] / [HInd] = K_HInd. Here, the indicator is considered as an acid; [H⁺] is the concentration of H-ions; [Ind^-] is the concentration of the ions of the indicator, or its alkaline form; [HInd] is the concentration of the molecules of the indicator, or its acidic form; K_HInd is the dissociation constant of the given indicator.

For the majority of indicators in the concentrations used during titration, the change in color becomes noticeable when about 10% of the total amount of indicator molecules transitions from one form to another. This interval, called the color change zone, is characteristic for each indicator and must be taken into account when selecting an indicator for titration. Three cases are distinguished: 1) titration of a strong acid with a strong base, 2) a weak acid with a strong base, and 3) a weak acid with a weak base. Titration curves (see figure) show that in the 1st case, upon neutralization of the entire acid, a sharp jump in pH occurs—from 3 to 11. In this case, any indicator whose color change zone lies between pH=4 and 10 can be used. In the second case, the jump of the titration curve covers a relatively small pH zone, for example, for acetic acid by means of n/10 NaOH; in these cases, only indicators whose color change occurs at a pH between 8 and 10 are suitable. In the third case, titration is especially difficult, because the pH jump will be very insignificant; the change in color will occur gradually, and one can judge the end of the reaction only by comparing the color of the titrated solution with the color of a solution of the medium salt of the acid and base that participated in the titration. The pH of the latter solution [in our case NH₄(C₂H₃O₂)] must be known for the correct selection of an indicator during titration. Instead of an indicator solution, one can use so-called indicator papers, which are high-quality filter paper impregnated with an indicator solution and dried in the air. Indicators belong to the most diverse classes of chemical substances: to nitro compounds, phthaleins, sulfophthaleins, azo dyes, etc. Many indicators give incorrect readings in liquids containing protein, in alcoholic solutions, and in the presence of salts in solutions.

The most commonly used indicators are the following: 1. Methyl violet - pentamethylpararosaniline hydrochloride C₂₄H₂₈N₃Cl, used for a qualitative test for free HCl in gastric juice. 0.1% aqueous solution. Color change from violet to green between pH=1.5–3.2; from green to yellow between pH=1.5–0.1. For 10 cm³ of liquid, 2–10 drops of the indicator solution are taken. It gives a significant error in liquids containing protein and at a high concentration of salts.

2. Tropaeolin 00 (Orange IV) - phenylaminoazobenzene-p-sulfonic acid sodium salt - C₆H₄(SO₃Na)N=NC₆H₄NHC₆H₅ (for use, see item 1). 0.1% aqueous solution changes color from red to orange-yellow between pH=1.3–3.0. For 10 cm³ of the liquid being tested, 2 drops of the indicator. The error in the presence of salts is insignificant. 3. Dimethylaminoazobenzene (dimethyl yellow) - C₆H₅N=NC₆H₄N(CH₃)₂. Convenient for the titration of weak bases, as well as alkalis bound to weak acids. Used when titrating free hydrochloric acid in gastric juice. 0.01% solution in 90° alcohol. The color transition zone from red to yellow is between pH=2.9–4.0. For 10 cm³, 1–4 drops of the solution are taken. 4. Congo red (congorot), one of the benzidine azo dyes, SO₃NaC₁₀H₆N=NC₆H₄C₆H₄N=NC₁₀H₆SO₃NaNH₂.

^XNH₂ is used most often in the form of indicator papers. It is used for a qualitative test for free HCl in gastric juice. 0.1% aqueous solution. Changes color from red to violet-blue between pH=3.0-5.2. For 10 cm³, 1-3 drops of Indicator 5 are taken. Methyl orange (helianthin B, orange III) - sodium dimethylaminoazobenzene-p-sulfonate - SO₃Na.C₆H₄.N=N.C₆H₄.N(CH₃)₂, is used in acidimetry. Not sensitive to weak acids. 0.1% aqueous solution. Changes color from red to orange-yellow between pH=3.0-4.4. For 10 cm³ of liquid - 1-4 drops of indicator. 6. Methyl red - dimethylaminoazobenzene-o-carboxylic acid - HOOC.C₆H₄.N=N.C₆H₄.N(CH₃)₂; a 0.2% solution is used: 1 g of Indicator is dissolved in 30 cm³ of alcohol and diluted with water to 500 cm³. A very commonly used indicator in acidimetry; it changes color from red to yellow between pH=4.4-6.2. For 10 cm³ of liquid - 1-4 drops of indicator. 7. Lacmoid (resorcin blue) - C₁₂H₉O₃N. A sign of the purity of the commercial preparation is its solubility in boiling water; a significant amount of the indicator should pass into the solution, coloring the water a beautiful blue color; the alcoholic solution should have a blue-violet color without a brown tint. To obtain a good preparation from commercial lacmoid, the latter is extracted with boiling water, filtered, and the dye is precipitated by slightly acidifying the blue liquid. After a few hours, the precipitate is collected on a filter and washed with cold water. It is dried at a low temperature or the filter with the precipitate is transferred into alcohol, which is allowed to evaporate on a water bath. With good commercial preparations, the yield is 40% of the starting material. Good commercial preparations are also purified by treatment with 96% alcohol and subsequent evaporation of the alcoholic filtrate in a vacuum over sulfuric acid. It is used in acidimetry. 0.2% alcoholic solution. The zone of color change from red to blue is between pH=4.4-6.4. For 10 cm³ of liquid - 1-5 drops of indicator. 8. Azolitmin - a component (4-5%) of commercial litmus. Azolitmin is used in acidimetry; papers impregnated with an azolitmin solution are used in the formol titration of amino acids according to Sorensen. 1 g of azolitmin is dissolved in 100 cm³ of slightly alkalized water and the resulting solution is carefully acidified to a violet color. The zone of color change from red to blue is between pH=5.0-8.0. For 10 cm³ of liquid - 1-10 drops of the indicator solution. 9. Phenolphthalein - C₂₀H₁₄O₄. Melting point 250°. Widely used in acidimetry. Convenient for the titration of weak acids with strong bases. 1% or 0.1% solution. Preparation: 5 or 0.5 g is dissolved in 300 cm³ of alcohol and made up with water to 500 cm³. Color change from colorless to red - between pH=8.2 and 10.0. In strongly alkaline liquids, the indicator becomes colorless again.