When carrying out a chemical process, it is extremely important to monitor the conditions for the course of the reaction or to establish the achievement of its completion. Sometimes this can be observed by some external signs: the cessation of the evolution of gas bubbles, a change in the color of the solution, precipitation, or, conversely, the transition of one of the reaction components into the solution, etc. In most cases, auxiliary reagents are used to determine the end of the reaction, so called indicators, which are usually introduced into the analyzed solution in small quantities.
indicators called chemical compounds that can change the color of the 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 environment; 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 the control of 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 required amount of them to be applied to the soil are established.
Distinguish acid-base, fluorescent, redox, adsorption and chemiluminescent indicators.
ACID-BASE (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 positively charged hydrogen ions and negatively charged hydroxyl ions:
The concentration of hydrogen ions in a solution is denoted by the symbol .
If the concentration of hydrogen and hydroxide ions in the solution is the same, then such solutions are neutral and pH = 7. At a concentration of hydrogen ions corresponding to pH from 7 to 0, the solution is acidic, but if the concentration of hydroxide ions is higher (pH = from 7 to 14), the solution alkaline.
Various methods are used to measure the pH value. Qualitatively, the reaction of the solution can be determined using special indicators that change their color depending on the concentration of hydrogen ions. Such indicators are acid-base indicators that respond to changes in the pH of the medium.
The vast majority of acid-base indicators are dyes or other organic compounds, the molecules of which undergo structural changes depending on the reaction of the medium. They are used in titrimetric analysis in neutralization reactions, as well as for colorimetric determination of pH.
| Indicator | Color transition pH range | Color change |
|---|---|---|
| methyl violet | 0,13-3,2 | Yellow - purple |
| thymol blue | 1,2-2,8 | Red - yellow |
| Tropeolin 00 | 1,4-3,2 | Red - yellow |
| - 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 improve the accuracy of pH measurement, then mixed indicators are used. To do this, two indicators are selected with close pH intervals of the color transition, having additional colors in this interval. With this mixed indicator, determinations can be made with an accuracy of 0.2 pH units.
Widely used are also universal indicators that can repeatedly change color in a wide range of pH values. Although the accuracy of determination by such indicators does not exceed 1.0 pH units, they allow determinations in a wide pH range: from 1.0 to 10.0. Universal indicators are usually a combination of four to seven two-color or single-color indicators with different color transition pH ranges, designed in such a way that when the pH of the medium changes, a noticeable color change occurs.
For example, the commercially available universal indicator PKC is a mixture of seven indicators: bromocresol purple, bromocresol green, methyl orange, tropeolin 00, phenolphthalein, thymol blue, and bromothymol blue.
This indicator, depending on pH, 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 added a few drops to the test solution. By changing the color of the solution, the pH value is judged. In addition to alcohol-soluble indicators, water-soluble forms are also produced, which are ammonium or sodium salts of these indicators.
In many cases, it is more convenient to use not indicator solutions, but indicator papers. The latter are prepared as follows: the filter paper is passed through a standard indicator solution, the paper is squeezed out of the excess solution, dried, cut into narrow strips and bookleted. To carry out the test, an indicator paper is dipped into the test solution or one drop of the solution is placed on a strip of indicator paper and a change in its color is observed.
FLUORESCENT INDICATORS
Some chemical compounds, when exposed to ultraviolet rays, have the ability, at a certain pH value, to cause the solution to fluoresce or change its color or shade.
This property is used for acid-base titration of oils, turbid and strongly 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 (under ultraviolet light) | Fluorescence color change |
| 4-Ethoxyacridone | 1,4-3,2 | Green - blue |
| 2-Naphthylamine | 2,8-4,4 | Increasing violet fluorescence |
| Dimetnlnaphteirodine | 3,2-3,8 | Lilac - orange |
| 1-Naphthylam | 3,4-4,8 | Increase in blue fluorescence |
| Acridine | 4,8-6,6 | Green - purple |
| 3,6-Dioxyphthalimide | 6,0-8,0 | yellow-green - yellow |
| 2,3-Dicyanhydroquinone | 6,8-8,8 | Blue; green |
| Euchrysin | 8,4-10,4 | Orange - green |
| 1,5-Naphthylaminesulfamide | 9,5-13,0 | Yellow green |
| CC-acid (1,8-aminonaphthol 2,4-disulfonic acid) | 10,0-12,0 | Purple - green |
REDOX INDICATORS
Redox indicators- chemical compounds that change the color of the solution depending on the value of the redox potential. They are used in titrimetric methods of analysis, as well as in biological research for the colorimetric determination of redox potential.
| Indicator | Normal redox potential (at pH=7), V | Mortar coloring | |
| oxidizing form | restored form | ||
| Neutral red | -0,330 | Red-violet | Colorless |
| Safranin T | -0,289 | brown | Colorless |
| Potassium indihomonosulfonate | -0,160 | Blue | Colorless |
| Potassium indigodisulfonate | -0,125 | Blue | Colorless |
| Potassium indigotrisulfonate | -0,081 | Blue | Colorless |
| Potassium inngtetrasulfonate | -0,046 | Blue | Colorless |
| Toluidine blue | +0,007 | Blue | Colorless |
| Tnonin | +0,06 | purple | Colorless |
| o-cresolindophenolate sodium | +0,195 | reddish blue | Colorless |
| Sodium 2,6-Dnchlorophenolindophenolate | +0,217 | reddish blue | Colorless |
| m-Bromophenolindophenolate sodium | +0,248 | reddish blue | Colorless |
| dipheinlbenzidine | +0.76 (acid solution) | purple | Colorless |
ADSORPTION INDICATORS
Adsorption indicators- substances in the presence of which the color of the precipitate formed during titration by the precipitation method changes. Many acid-base indicators, some dyes and other chemical compounds are able to change the color of the precipitate at a certain pH value, which makes them suitable for use as adsorption indicators.
| Indicator | Defined ion | Ion precipitant | Color change |
| Alizarin Red C | Yellow - rose red | ||
| Bromophenol blue | Yellow - green | ||
| Lilac - yellow | |||
| Purple - blue-green | |||
| Diphenylcarbazide | , , | Colorless - violet | |
| Congo red | , , | Red - blue | |
| Blue - red | |||
| Fluorescein | , | yellow-green - pink | |
| Eosin | , | yellow-red - red-violet | |
| Erythrosine | Red-yellow - dark red |
CHEMILUMINESCENT INDICATORS
This group of indicators includes substances capable of emitting visible light at certain pH values. Chemiluminescent indicators are convenient to use when working with dark liquids, since in this case a glow appears at the end point of the titration.
INDICATORS(from lat. indicator - pointer) - substances that allow you to monitor the composition of the environment or the progress of a chemical reaction. One of the most common is acid-base indicators, which change color depending on the acidity of the solution. This happens because in an acidic and alkaline environment, the indicator molecules have a different structure. An example is the common indicator phenolphthalein, which was previously also used as a laxative called purgen. In an acidic medium, this compound is in the form of undissociated molecules, and the solution is colorless, and in an alkaline medium, in the form of singly charged anions, and the solution has a crimson color ( cm. ELECTROLYTIC DISSOCIATION. ELECTROLYTES). However, in a strongly alkaline environment, phenolphthalein becomes colorless again! This happens due to the formation of another colorless form of the indicator - in the form of a three-charged anion. Finally, in a medium of concentrated sulfuric acid, a red color appears again, although not as intense. Its culprit is the phenolphthalein cation. This little-known fact can lead to an error in determining the reaction of the environment.
Acid-base indicators are very diverse; many of them are easily accessible and therefore known for more than one century. These are decoctions or extracts of colored flowers, berries and fruits. So, a decoction of iris, pansies, tulips, blueberries, blackberries, raspberries, black currants, red cabbage, beets and other plants turns red in an acidic environment and green-blue in an alkaline one. This is easy to see if you wash the pot with the remnants of borscht with soapy (i.e. alkaline) water. Using an acidic solution (vinegar) and an alkaline solution (drinking, or better, washing soda), you can also make inscriptions on the petals of various colors in red or blue.
Ordinary tea is also an indicator. If you drop lemon juice or dissolve a few crystals of citric acid into a glass of strong tea, the tea will immediately become lighter. If you dissolve baking soda in tea, the solution will darken (of course, you should not drink such tea). Tea made from flowers (“karkade”) gives much brighter colors.
Probably the oldest acid-base indicator is litmus. Back in 1640, botanists described the heliotrope (Heliotropium Turnesole) - a fragrant plant with dark purple flowers, from which a dye was isolated. This dye, along with the juice of violets, became widely used by chemists as an indicator, which was red in an acidic environment and blue in an alkaline one. This can be read in the writings of the famous 17th century physicist and chemist Robert Boyle. Initially, with the help of a new indicator, mineral waters were investigated, and from about 1670 they began to use it in chemical experiments. “As soon as I add a slightly small amount of acid,” the French chemist Pierre Pomet wrote about “tournesol” in 1694, “it turns red, so if anyone wants to know if something contains acid, it can be used.” In 1704, a German scientist M. Valentin called this paint litmus, this word has remained in all European languages except French, in French litmus is tournesol, which literally means "turning after the sun". the same thing, only in Greek.It soon turned out that litmus can be extracted from cheaper raw materials, for example, from certain types of lichens.
Unfortunately, almost all natural indicators have a serious drawback: their decoctions deteriorate rather quickly - turn sour or moldy (alcoholic solutions are more stable). Another disadvantage is the too wide range of color change. In this case, it is difficult or impossible to distinguish, for example, a neutral medium from a slightly acidic one or a slightly alkaline one from a strongly alkaline one. Therefore, in chemical laboratories, synthetic indicators are used that sharply change their color within fairly narrow pH limits. There are many such indicators, and each of them has its own scope. For example, methyl violet changes color from yellow to green in the pH range of 0.13 - 0.5; methyl orange - from red (pH< 3,1) до оранжево-желтой (рН 4); бромтимоловый синий – от желтой (рН < 6,0) до сине-фиолетовой (рН 7,0); фенолфталеин – от бесцветной (рН < 8,2) до малиновой (рН 10); тринитробензол – от бесцветной (pH < 12,2) до оранжевой (рН 14,0).
In laboratories, universal indicators are often used - a mixture of several individual indicators, selected so that their solution alternately changes color, passing through all the colors of the rainbow when the acidity of the solution changes over a wide pH range (for example, from 1 to 11). Strips of paper are often impregnated with a solution of a universal indicator, which allows you to quickly (albeit with not very high accuracy) determine the pH of the analyzed solution by comparing the color of the strip moistened with the solution with a reference color scale.
In addition to acid-base indicators, other types of indicators are also used. So, redox indicators change their color depending on whether an oxidizing or reducing agent is present in the solution. For example, the oxidized form of diphenylamine is purple, while the reduced form is colorless. Some oxidizing agents can themselves serve as an indicator. For example, when analyzing iron(II) compounds in the course of the reaction
10FeSO4 + 2KMnO4 + 8H2SO4? 5Fe 2 (SO 4) 3 + 2MnSO 4 + K 2 SO 4 + 8H 2 O
the added permanganate solution becomes colorless as long as Fe 2+ ions are present in the solution. As soon as the slightest excess of permanganate appears, the solution acquires a pink color. By the amount of permanganate consumed, it is easy to calculate the iron content in the solution. Similarly, in numerous analyzes using the iodometry method, iodine itself serves as an indicator; to increase the sensitivity of the analysis, starch is used, which makes it possible to detect the slightest excess of iodine.
Complesonometric indicators are widely used - substances that form colored complex compounds with metal ions (many of which are colorless). An example is eriochrome black T; the solution of this complex organic compound has a blue color, and in the presence of magnesium, calcium and some other ions, complexes are formed that are colored in an intense wine-red color. The analysis is carried out as follows: a solution containing the analyzed cations and an indicator is added dropwise to a stronger complexing agent than the indicator, most often Trilon B. As soon as Trilon completely binds all metal cations, there will be a distinct transition from red to blue. From the amount of trilon added, it is easy to calculate the content of metal cations in the solution.
Other types of indicators are also known. For example, some substances are adsorbed on the surface of the sediment, changing its color; such indicators are called adsorption. When titrating cloudy or colored solutions, in which it is almost impossible to notice a change in the color of conventional acid-base indicators, fluorescent indicators are used. They glow (fluoresce) in different colors depending on the pH of the solution. For example, the fluorescence of acridine changes from green at pH = 4.5 to blue at pH = 5.5; it is important that the luminescence of the indicator does not depend on the transparency and intrinsic color of the solution.
Ilya Leenson
Color change of indicators depending on pH
Acid-base indicators are compounds whose color changes depending on the acidity of the medium.
For example, litmus is red in an acidic environment and blue in an alkaline environment. This property can be used to quickly evaluate the pH of solutions.
Acid-base indicators are widely used in chemistry. It is known, for example, that many reactions proceed differently in acidic and alkaline media. By adjusting the pH, the direction of the reaction can be changed. Indicators can be used not only for qualitative, but also for quantitative assessment of the acid content in a solution (acid-base titration method).
The use of indicators is not limited to "pure" chemistry. The acidity of the environment must be controlled in many production processes, when assessing the quality of food products, in medicine, etc.
IN table 1 the most "popular" indicators are indicated and their color in neutral, acidic and alkaline media is noted.
Table 1
Methyl orange
Phenolphthalein
In fact, each indicator is characterized by its own pH interval in which the color change occurs (transition interval). The change in color occurs due to the transformation of one form of the indicator (molecular) into another (ionic). As the acidity of the medium decreases (with an increase in pH), the concentration of the ionic form increases, and that of the molecular form decreases. Table 2 lists some acid-base indicators and their respective transition ranges.
table 2Substances that change color when the reaction of the medium changes are indicators - most often complex organic compounds - weak acids or weak bases. Schematically, the composition of indicators can be expressed by the formulas НInd or IndOH, where Ind is a complex organic anion or indicator cation.
In practice, indicators have been used for a long time, but the first attempt to explain their action was made in 1894 by Ostwald, who created the so-called ionic theory. According to this theory, undissociated indicator molecules and its Ind ions have different colors in solution, and the color of the solution changes depending on the position of indicator dissociation equilibrium. For example, phenolphthalein (an acid indicator) has colorless molecules and crimson anions; methyl orange (main indicator) - yellow molecules and red cations.
phenolphthalein methyl orange
HIndH + + Ind–IndOH 
Ind + +OH-
colorless raspberries. yellow red
A change in accordance with Le Chatelier's principle leads to a shift in equilibrium to the right or to the left.
According to the chromophore theory (Hanch), which appeared later, the change in the color of indicators is associated with a reversible rearrangement of atoms in the molecule of an organic compound. Such a reversible rearrangement in organic chemistry is called tautomerism. If, as a result of a tautomeric change in the structure, special groups called chromophores appear in the molecule of an organic compound, then the organic substance acquires a color. Chromophores are groups of atoms that contain one or more multiple bonds that cause selective absorption of electromagnetic vibrations in the UV region. Groupings of atoms and bonds, such as −N=N− , =C=S , −N=O, quinoid structures, etc., can act as chromophore groups.
When a tautomeric transformation leads to a change in the structure of the chromophore, the color changes; if, after rearrangement, the molecule no longer contains a chromophore, the color will disappear.
Modern ideas are based on the ion-chromophoric theory, according to which the change in the color of the indicators is due to the transition from the ionic form to the molecular one, and vice versa, accompanied by a change in the structure of the indicators. Thus, one and the same indicator can exist in two forms with different molecular structures, and these forms can transform into one another, and an equilibrium is established between them in solution.
As an example, we can consider structural changes in the molecules of typical acid-base indicators - phenolphthalein and methyl orange under the action of alkali and acid solutions (at different pH values).
The reaction, as a result of which, due to the tautomeric rearrangement of the structure of the phenolphthalein molecule, a chromophore group arises in it, which causes the appearance of color, proceeds according to the following equation:
colorless colorless colorless
crimson
Indicators, as weak electrolytes, have small dissociation constants. For example, K d of phenolphthalein is 2 ∙ 10 -10 and in neutral media it is found mainly in the form of its molecules due to a very low concentration of ions, which is why it remains colorless. When alkali is added, H + -ions of phenolphthalein bind, "tighten" with OH - alkali ions, forming water molecules, and the indicator dissociation equilibrium position shifts to the right - towards an increase in the concentration of Ind - ions. In an alkaline medium, a disodium salt is formed, which has a quinoid structure, which causes the color of the indicator. The shift in equilibrium between tautomeric forms occurs gradually. Therefore, the color of the indicator does not change immediately, but passes through a mixed color to the color of the anions. When acid is added to the same solution simultaneously with the neutralization of alkali - at a sufficient concentration of H + -ions - the equilibrium position of the dissociation of the indicator shifts to the left, towards molarization, the solution becomes discolored again.
Similarly, the color of methyl orange changes: neutral molecules of methyl orange give the solution a yellow color, which, as a result of protonation, turns into red, corresponding to the quinoid structure. This transition is observed in the pH range 4.4–3.1:
yellow Red
Thus, the color of the indicators depends on the pH environment. The color intensity of such indicators is quite high and is clearly visible even with the introduction of a small amount of the indicator, which is not able to significantly affect the pH of the solution.
A solution containing an indicator changes color continuously as the pH changes. The human eye, however, is not very sensitive to such changes. The range in which the color change of the indicator is observed is determined by the physiological limits of color perception by the human eye. With normal vision, the eye is able to distinguish the presence of one color in a mixture of it with another color only if there is at least some threshold density of the first color: a change in the color of the indicator is perceived only in the area where there is a 5-10-fold excess of one form in relation to another. Taking HInd as an example and characterizing the state of equilibrium
Hind 
H + + Ind-
corresponding constant
,
it can be written that the indicator shows its purely acid color, usually captured by the observer, when
,
and a purely alkaline color at
Within the interval determined by these values, a mixed color of the indicator appears.
Thus, the eye of the observer distinguishes a change in color only when the reaction of the medium changes in the range of about 2 pH units. For example, for phenolphthalein, this pH range is from 8.2 to 10.5: at pH = 8.2, the eye observes the beginning of the appearance of a pink color, which intensifies to pH = 10.5, and at pH = 10.5, an increase in red color already invisible. This range of pH values, in which the eye distinguishes a change in the color of the indicator, is called the transition interval of the color of the indicator. For methyl orange, K D = 1.65 10 -4 and pK = 3.8. This means that at pH = 3.8, the neutral and dissociated forms are in equilibrium in approximately equal concentrations.
The specified pH range of approximately 2 units for various indicators does not fall in the same region of the pH scale, since its position depends on the specific value of the dissociation constant of each indicator: the stronger the acid HInd , the more acidic the transition interval of the indicator is . In table. 18 shows the transition intervals and colors of the most common acid-base indicators.
For a more accurate determination of the pH value of solutions, a complex mixture of several indicators applied to filter paper (the so-called "Kolthoff universal indicator") is used. A strip of indicator paper is dipped in the test solution, placed on a white waterproof substrate, and the color of the strip is quickly compared with the reference scale for pH.
Table 18
Transition intervals and coloring in various media
the most common acid-base indicators
|
Name |
Indicator color in different environments |
||
|
Phenolphthalein |
colorless |
crimson 8.0 < pH < 9.8 |
crimson |
|
purple 5 < рН < 8 | |||
|
Methyl Orange |
Orange | ||
|
3.1< рН < 4.4 | |||
|
Methyl purple |
purple | ||
|
Bromocresol | |||
|
Bromothymol | |||
|
thymol |
2,5 < pH < 7,9 | ||
Indicators- organic compounds that can change color in solution with a change in acidity (pH). Indicators are widely used in titration in analytical chemistry and biochemistry. Their advantage is the low cost, speed and visibility of the study.
Indicators are usually used by adding a few drops of an aqueous or alcoholic solution, or a little powder, to a sample of the test solution. So, during titration, an indicator is added to an aliquot of the test solution, and the color changes at the equivalence point are observed.
Indicator Color Transition Intervals
The figure shows approximate data on the existence of different color forms of indicators in aqueous solutions.
For more precise information (multiple transitions, numerical pH value), see the next section.
Table of pH transition values for the most common indicators
Acid-base indicators common in laboratory practice are given in ascending order of pH values that cause a color change. Roman numerals in square brackets correspond to the color transition number (for indicators with multiple transition points).
| Indicator and transition number | X | Color more acid form | pH interval and transition number | Color more alkaline form |
||
|---|---|---|---|---|---|---|
| methyl violet | yellow | 0.13-0.5 [I] | green | |||
| Cresol Red [I] | Red | 0.2-1.8 [I] | yellow | |||
| methyl violet | green | 1,0-1,5 | blue | |||
| Thymol blue [I] | to | Red | 1.2-2.8 [I] | yellow | ||
| Tropeolin 00 | o | Red | 1,3-3,2 | yellow | ||
| methyl violet | blue | 2,0-3,0 | purple | |||
| (Di)methyl yellow | o | Red | 3,0-4,0 | yellow | ||
| Bromophenol blue | to | yellow | 3,0-4,6 | blue-violet | ||
| Congo red | Red | 3,0-5,2 | blue | |||
| methyl orange | o | Red | 3,1-(4,0)4,4 | (orange-)yellow | ||
| Bromocresol green | to | yellow | 3,8-5,4 | blue | ||
| Bromocresol blue | yellow | 3,8-5,4 | blue | |||
| Lakmoid | to | Red | 4,0-6,4 | blue | ||
| methyl red | o | Red | 4,2(4,4)-6,2(6,3) | yellow | ||
| Chlorophenol red | to | yellow | 5,0-6,6 | Red | ||
| Litmus (azolithine) | Red | 5,0-8,0 (4,5-8,3) | blue | |||
| Bromocresol purple | to | yellow | 5,2-6,8(6,7) | bright red | ||
| Bromothymol blue | to | yellow | 6,0-7,6 | blue | ||
| Neutral red | o | Red | 6,8-8,0 | amber yellow | ||
| Phenol red | about | yellow | 6,8-(8,0)8,4 | bright red | ||
| Cresol Red | to | yellow | 7,0(7,2)-8,8 | Dark red | ||
| α-Naphtholphthalein | to | yellow-pink | 7,3-8,7 | blue | ||
| thymol blue | to | yellow | 8,0-9,6 | blue | ||
| Phenolphthalein [I] | to | colorless | 8.2-10.0 [I] | crimson red | ||
| thymolphthalein | to | colorless | 9,3(9,4)-10,5(10,6) | blue | ||
| Alizarin yellow LJ | to | pale lemon yellow | 10,1-12,0 | brown yellow | ||
| Nile blue | blue | 10,1-11,1 | Red | |||
| diazo violet | yellow | 10,1-12,0 | purple | |||
| indigo carmine | blue | 11,6-14,0 | yellow | |||
| Epsilon Blue | Orange | 11,6-13,0 | dark purple | |||
