When carrying out a chemical process, it is extremely important to follow the conditions of the reaction or to establish the achievement of its end. Sometimes this can be observed according to some external signs: the cessation of gas bubble evolution, the color of the solution, precipitation or, conversely, the transition of one of the components of the reaction to the solution, etc. In most cases, auxiliary agents are used to determine the end of the reaction, so called indicators, which are usually introduced into the analyzed solution in small quantities.

Indicators are chemical compounds that can change the color of a solution depending on environmental conditions, without directly affecting the test solution and the direction of the reaction. So, acid-base indicators change color depending on the pH of the medium; redox indicators - from the potential of the medium; adsorption indicators - on the degree of adsorption, etc.

Indicators are especially widely used in analytical practice for titrimetric analysis. They also serve as the most important tool for controlling technological processes in the chemical, metallurgical, textile, food and other industries. In agriculture, with the help of indicators, analysis and classification of soils are carried out, the nature of fertilizers and the amount of fertilizers needed for their application are established.

There are acid-base, fluorescent, redox, adsorption and chemiluminescent indicators.

ACID-ALKALINE (PH) INDICATORS

As is known from the theory of electrolytic dissociation, chemical compounds dissolved in water dissociate into positively charged ions - cations and negatively charged - anions. Water also dissociates to a very small extent into hydrogen ions, positively charged, and hydroxyl ions, negatively charged:

$\mathrm{H_2O_2 \rightleftarrows H^++HO_2^-}$
The concentration of hydrogen ions in solution is indicated by $\mathrm{pH}$ .

If the concentration of hydrogen and hydroxyl ions in the solution is the same, then such solutions are neutral and pH = 7. At a concentration of hydrogen ions corresponding to a pH of 7 to 0, the solution is acidic, but if the concentration of hydroxyl ions is higher (pH = 7 to 14), the solution alkaline.

Various methods are used to measure pH. Qualitatively, the reaction of the solution can be determined using special indicators that change their color depending on the concentration of hydrogen ions. These indicators are acid-base indicators that respond to changes in pH.

The overwhelming majority of acid-base indicators are dyes or other organic compounds whose molecules undergo structural changes depending on the reaction of the medium. They are used in titrimetric analysis for neutralization reactions, as well as for colorimetric determination of pH.

Indicator Coloring range pH Discoloration
Methyl violet 0.13-3.2 Yellow - purple
Thymol blue 1.2-2.8 Red - yellow
Tropeolin 00 1.4-3.2 Red - yellow
$\mathrm{\beta}$ - dinitrophenol 2.4-4.0 Colorless - Yellow
Methyl orange 3.1-4.4 Red - yellow
Naphthyl red 4.0-5.0 Red - Orange
Methyl red 4.2-6.2 Red - yellow
Bromothymol blue 6.0-7.6 Yellow - blue
Phenol red 6.8-8.4 Yellow - red
Metacresol Purple 7.4-9.0 Yellow - purple
Thymol blue 8.0-9.6 Yellow - blue
Phenolphthalein 8.2-10.0 Colorless - Red
Thymolphthalein 9.4-10.6 Colorless - Blue
Alizarin yellow P 10.0-12.0 Pale yellow - red orange
Tropeolin 0 11.0-13.0 Yellow - Medium
Malachite green 11.6-13.6 Greenish Blue - Colorless

If it is necessary to increase the accuracy of pH measurement, mixed indicators are used. For this, two indicators are selected with close intervals of the pH of the color transition, having additional colors in this interval. Using such a mixed indicator, it is possible to determine with an accuracy of 0.2 pH units.

Universal indicators are also widely used, capable of repeatedly changing color in a wide range of pH values. Although the accuracy of determination by such indicators does not exceed 1.0 pH units, but they allow the determination to be carried out in a wide pH range: from 1.0 to 10.0. Universal indicators are usually a combination of four to seven two-color or one-color indicators with different intervals of pH of the color transition, designed so that when the pH of the medium changes, a noticeable color change occurs.

For example, a universal PKC indicator produced by the industry is a mixture of seven indicators: bromocresol purple, bromocresol green, methyl orange, tropeolin 00, phenolphthalein, thymol blue and bromothymol blue.

Depending on pH, this indicator has the following color: at pH = 1 - raspberry, pH = 2 - pinkish-orange, pH = 3 - orange, pH = 4 - yellow-orange, pH = 5 yellow, pH = 6 - greenish yellow, pH = 7 - yellow-green. PH = 8 - green, pH = 9 - blue-green, pH = 10 - grayish-blue.

Individual, mixed and universal acid-base indicators are usually dissolved in ethanol and a few drops are added to the test solution. By changing the color of the solution, a pH value is judged. In addition to alcohol-soluble indicators, water-soluble forms are also available, which are ammonium or sodium salts of these indicators.

In many cases, it is more convenient to use indicator solutions rather than indicator solutions. The latter are prepared as follows: filter paper is passed through a standard indicator solution, the paper is squeezed from the excess solution, dried, cut into narrow strips and stitched into booklets. To carry out the test, the indicator paper is lowered into the test solution or one drop of the solution is placed on a strip of indicator paper and its color change is observed.

FLUORESCENT INDICATORS

Some chemical compounds, when exposed to ultraviolet rays, have the ability, at a certain pH value, to cause fluorescence of a solution or change its color or shade.

This property is used for acid-base titration of oils, turbid and highly colored solutions, since conventional indicators are unsuitable for these purposes.

Work with fluorescent indicators is carried out by illuminating the test solution with ultraviolet light.

 Indicator Fluorescence pH range (in ultraviolet light) Fluorescence color change 4-ethoxyacridone 1.4-3.2 Green - blue 2-naphthylamine 2.8-4.4 The increase in violet fluorescence Dimetnlnafteirodin 3.2-3.8 Purple - Orange 1-Naphthylamine 3.4-4.8 Blue fluorescence buildup Acridine 4.8-6.6 Green - Purple 3,6-Dioxiftalimide 6.0-8.0 Yellow green yellow 2,3-dicyanohydroquinone 6.8-8.8 Blue; green Euchrysin 8.4-10.4 Orange - Green 1,5-Naphthylamine sulfamide 9.5-13.0 Yellow green SS-acid (1,8-aminonaphthol 2,4-disulfonic acid) 10.0-12.0 Violet - Green

Redox indicators

Redox indicators are chemical compounds that change the color of a solution depending on the value of the redox potential. They are used in titrimetric methods of analysis, as well as in biological studies for colorimetric determination of the redox potential.

 Indicator Normal redox potential (at pH = 7), V Solution color oxidative form restored form Neutral red —0.330 Red violet Colorless Safranin T —0.289 Brown Colorless Potassium indigomonosulfonate —0.160 Blue Colorless Potassium indigodisulfonate -0.125 Blue Colorless Potassium indigotrisulfonate —0.081 Blue Colorless Potassium indngotetrasulfonate —0.046 Blue Colorless Toluidine Blue +0.007 Blue Colorless Tnonin +0.06 Purple Colorless Sodium o-Cresolindophenolate +0.195 Reddish blue Colorless Sodium 2,6-dichlorophenolindophenolate +0.217 Reddish blue Colorless m-Bromophenolindophenolate sodium +0.248 Reddish blue Colorless Dipheinlbenzidine +0.76 (acidic solution) Purple Colorless

 Indicator Detectable ion Ion precipitator Discoloration Alizarin Red C $\mathrm{Fe(CN){}^{4-}_6}$ $\mathrm{Pb{}^{2+}}$ Yellow - Rose Red Bromophenol Blue $\mathrm{Tl{}^+}$ $\mathrm{I{}^-}$ Yellow - green $\mathrm{Hg{}^{2+}}$ $\mathrm{Cl{}^{-}}$ Lilac - Yellow $\mathrm{SCN{}^{-}}$ $\mathrm{Ag{}^+}$ Violet - Blue Green Diphenylcarbazide $\mathrm{Cl{}^{-}}$ , $\mathrm{Br{}^{-}}$ , $\mathrm{Hg{}^{+}}$ Colorless - Violet Congo Red $\mathrm{Br{}^{-}}$ , $\mathrm{Cl{}^{-}}$ , $\mathrm{Ag{}^+}$ Red - Blue $\mathrm{Ag{}^+}$ $\mathrm{Cl{}^{-}}$ Blue - red Fluorescein $\mathrm{Cl{}^{-}}$ , $\mathrm{Br{}^{-}}$ $\mathrm{Ag{}^+}$ Yellow green pink Eosin $\mathrm{Br{}^{-}}$ , $\mathrm{I{}^-}$ $\mathrm{Ag{}^+}$ Yellow Red - Red Violet Erythrosine $\mathrm{MoO{}^{2-}_4}$ $\mathrm{Pb{}^{2+}}$ Red Yellow - Dark Red