INDICATORS in chemistry(lat. indicator pointer) - substances that change their color in the presence of certain chemical compounds in the medium under study (in solution, in air, in cells, in tissues), as well as when the pH or redox potential of the medium changes; are widely used in biochemical, clinical and sanitary laboratories.

I. is used to determine the end of the reaction (equivalence point) during titration, for the colorimetric determination of pH values ​​or redox potentials, for the detection of various kinds of substances in certain objects under study. For all these purposes, I. is used in the form of water or alcohol solutions or in the form of indicator papers, which are strips of filter paper soaked in I.

Depending on purpose and the mechanism of action And. subdivide into a number of groups.

Acid-base indicators are complex organic compounds that change color (two-color I.) or its intensity (one-color I.) depending on the pH of the medium. Two-color I. is, for example, a lakmoid: in an alkaline environment it has a blue color, and in an acidic environment it is red. An example of monochromatic I. is phenolphthalein, which is colorless in acidic environment and raspberry to alkaline.

According to the theory of Ostwald (W. Ostwald), acid-base I. are weak organic to-you or bases, the undissociated molecules of which have a different color in p-re than the anions and cations they form. Phenolphthalein, for example, is weak to - that, not dissociated molecules a cut are colorless, and anions paint solutions in raspberry color. In the solutions of I., which are weak to-you, dissociate according to the equation

where HA are non-dissociated I. molecules, H + are hydrogen ions, and A - are I. anions.

The ionization constant of such I. is equal to

Ka \u003d [H + ] [A - ] / [NA] (2)

(square brackets denote the molar concentrations of the corresponding particles).

I., which are weak bases, dissociate according to the equation

where BOH are non-dissociated I. molecules, B+ are I. cations, and OH- are hydroxyl ions.

The dissociation constant of these I. is

Kb = / (4)

It follows from equations 2 and 4 that the greater the dissociation constant, the more I. break down into ions and, consequently, at higher concentrations of H + ions (in cases where P. is a weak acid) or OH ions - (in cases where I.- weak base) its dissociation is suppressed and a color change occurs. Different I. have different values ​​of Ka and Kb. Therefore, they change their color at different pH values ​​of the medium. That interval of pH values, in which the color change of this I. occurs, is called the zone of action or the transition interval of I. The transition interval of I. is usually equal to pK ± 1, where pK is -lgK. The transition point of I. is called that pH value, at which the color change of I is visually most clearly perceived. The transition point is approximately equal to the pK value of this I.

Acid-base I. are widely used in titration to - t and alkalis, as well as for colorimetric measurement of the pH value of biol, liquids, cells, tissues, etc.

Titration of to-t and alkalis should be completed at the moment of reaching the equivalence point, i.e. at the moment when such a volume of titrant is added to the titrated solution to-you (alkali), in Krom contains an equivalent amount of to-you ( alkali). To do this, it is necessary to apply such an I., the transition point to-rogo is equal to the pH value of the titrated solution at the equivalence point (see Neutralization method). In table. I. are listed, most used in titration to - t and bases.

Qualitative determination of acidity and alkalinity is carried out using the so-called. neutral I., the transition point of which is practically at pH 7.0. These include, for example, litmus, which in an acidic environment (pH less than 7.0) is red, and in an alkaline environment (pH more than 7.0) Blue colour; neutral red, turning red in an acidic environment, and yellow in an alkaline environment.

An approximate measurement of the pH value of the medium (with an accuracy of 0.5-1.0 pH units) is usually carried out using a universal (combined) I., which is a mixture of several I., the transition intervals of which are close to each other and cover a wide range of values pH.

To 0.5 ml of the test liquid, add 1-2 drops of universal I solution. AND.

For a more accurate (0.1-0.5 pH unit) colorimetric determination of the pH value, one-color I. are usually used. in an acidic medium) to yellow (in an alkaline medium). With the same purpose use a number of two-color And. offered by Clark (W. M. Clark) and Labs (H. A. Lubs), which are sulfophthaleins. The acid and alkaline forms of these I. differ sharply in color, this is their advantage over Michaelis indicators.

Redox or redox indicators, are organic dyes, the color of which in the oxidized and reduced state is different. Such I. are used in oxidimetric titration (see Oxidimetry), as well as for the colorimetric determination of the redox potentials of liquids (see Redox potential), individual cells and tissues in cytochemical and cytol laboratories. Most redox indicators turn into colorless compounds upon reduction, and become colored upon oxidation. Oxidized and reduced forms of I. are in solutions in a state of dynamic equilibrium:

oxidized form + ne<->reduced form, where n is the number of electrons.

The ratio between the equilibrium concentrations of the two forms of this I., and therefore the color of the solution, in Krom is I., depends on the magnitude of the redox potential of the solution. If the value of the potential of the solution is greater than the normal redox potential (E0) of this redox indicator, then most of the I. in this solution passes into the oxidized form (usually colored), if the redox potential of the medium under study is less than E0, then I. turns into a reduced form (usually colorless). At equality of values ​​of redox potential of the environment and E0 of the indicator of concentration of the oxidized and restored forms I. are equal to each other. Having a series of I. with different values ​​of E0, it is possible to judge the magnitude of the redox potential of a given medium by their color in a given environment. The redox indicators proposed by Michaelis, which have the common name "viologens" and are derivatives of gamma and gamma "-dipyridyls, have low toxicity and are widely used to measure redox potentials in biol, systems; in these I., the reduced form is colored.

The normal redox potential of viologens does not depend on the pH value of the solution. This is what distinguishes them from other redox indicators.

Complexometric indicators (metal indicators) are water-soluble organic dyes capable of forming colored complex compounds with metal ions. These I. are used to establish the equivalence point in complexometric titration (see Complexometry).

Adsorption indicators- These are organic dyes adsorbed on the surface of precipitates formed during titration by the precipitation method, and change their color when the equivalence point is reached. For example, when chlorides are titrated with silver nitrate, tropeolin 00 changes color at the equivalence point from yellow to pink.

Chemiluminescent (fluorescent) indicators- organic compounds (for example, lumenol, luceginin, silaxen, etc.) that have the ability to luminesce in natural light or when irradiated with ultraviolet light. The intensity and color of luminescence depend both on the pH value of the medium and on the value of its redox potential; these I. are used in titration (during neutralization and oxidimetry) of highly colored or turbid liquids, when a change in the color of ordinary I. is imperceptible.

And. are used in many biochemical. the methods applied in klin. - biochem. laboratories. The most commonly used of them are bromthymol blue (when determining the activity of fructose diphosphate aldolase in the blood serum, the activity of acetylcholinesterase and cholinesterase in the blood serum according to A. A. Pokrovsky, as well as the activity of carboxylesterase in the blood according to A. A. Pokrovsky and L. G. Ponomareva), bromophenol blue (in the electrophoretic separation of various proteins for coloring electrophoregrams along with amido black and acid blue-black), universal I., phenol red (when determining the activity of aspartate and alanine aminotransferases in blood serum, cholinesterase activity in blood serum, etc. .), phenolphthalein, nitrosine tetrazolium used for quality and quantification activity of various dehydrogenases (see. Dehydrogenases), etc.

Name of the indicator	Indicator transition interval, in pH units
		Acid form of the indicator	Alkaline indicator
Tropeolin 00 (sodium diphenylaminoazo-n-benzenesulfonate)
Dimethyl yellow (dimethylaminoazobenzene)		orange red
Methyl orange (sodium)
Methyl red (dacid)
Phenol red (phenolsulfophthalein)
Phenolphthalein		Colorless	Crimson
thymolphthalein		Colorless

Bibliography: Vinogradova E. N. Methods for determining the concentration of hydrogen ions, M., 1956, bibliogr.; Indicators, ed. E. Bishop and I. N. Marov, trans. from English, vol. 1-2, M., 1976, bibliogr.

INDICATORS(Late Latin indicator - pointer), chem. in-va, changing color, or forming a precipitate when changing to.-l. component in solution. They indicate a certain state of the system or at the moment of reaching this state. There are reversible and irreversible indicators. The change in color of the first when the state of the system changes (for example, when the pH of the medium changes) can be. repeated many times. Irreversible indicators undergo irreversible chem. transformations, for example, at BrO 3 - are destroyed. Indicators, to-rye injected into the test solution, called. internal, unlike external, p-tion with to-rymi is carried out outside the analyzed mixture. In the latter case, one or more drops of the analyzed solution are placed on a piece of paper impregnated with an indicator, or they are mixed on a white porcelain plate with a drop of indicator. And Indicators are most often used to establish the end of c.-l. chem. r-tion, Ch. arr. end point (k.t.t.). According to the titrimetric methods distinguish between acid-base, adsorption, oxidizing-reducing. and complexometric. indicators. are p-rime org comp., to-rye change their color or depending on H + (pH of the medium). Appl. to establish the end of the p-tion between to-tami and (including at) or other p-tions, if they involve H +, as well as for colorimetric. determination of pH of water solutions. Naib. important are given in table. 1. The reason for the change in the color of the indicators is that its addition or release is associated with the replacement of some chromophore groups by others or with the appearance of new chromophore groups. If the indicator is weak to-ta HIn, then in the aqueous solution takes place: HIn + H 2 O D In- + H 3 O +. If the indicator is weak In, then: In + H 2 O D HIn + + OH - . In general terms, we can write: In a + H 2 O D In b + H 3 O + , where In a and In b - respectively. acidic and basic forms of the indicator, which are colored differently. this process K ln = / naz. indicator. The color of the solution depends on the ratio /, a cut is determined by the pH of the solution.

It is believed that the color of one form of the indicator is noticeable if it is 10 times higher than other forms, i.e. if the ratio / \u003d /K ln is 0.1 or 10. The change in the color of the indicator is noted in the region of pH \u003d pK lp b 1, to-ry called. indicator transition interval. Change max. distinctly when = and K ln = [H 3 O] +, i.e. at pH = pKln. The pH value, at Krom, usually ends, called. RT indicator. Indicators for are selected in such a way that the color transition interval includes the pH value that the solution should have at the equivalence point. Often this pH value does not match the pT of the indicator used, which leads to the so-called. indicator error. If an excess of untitrated weak or to-you remains in K. t. t., the error is called. resp. basic or acid. Indicator sensitivity - (in / l) determined (in this case H+ or OH - ) at the point of max. abrupt color change. Distinguish: indicators, sensitive to there, with an interval of transition in the field of alkaline pH values ​​(eg, thymolphthalein); sensitive to indicators with a transition interval in the acidic region (as in dimethyl yellow, etc.); neutral indicators, the transition interval to-rykh is approx. pH 7 (neutral red, etc.). And indicators come with one or two colored forms; such indicators are called resp. one-color and two-color. Naib. a clear change in color would be observed in those indicators, the acid and basic forms of which are colored in addition. colors. However, there are no such indicators. Therefore, by adding , the colors of both forms are changed accordingly. So, in methyl red, the transition from red to yellow occurs in the range of 2 pH units, and if you add to the solution, then the color transition from red-violet to green is observed sharply and clearly at pH 5.3. A similar effect can be achieved by using a mixture of two indicators, the colors of which complement each other. friend. Such indicators are called mixed (Table 2).


Mixtures of indicators, to-rye continuously change their color in the entire range of pH values ​​from 1 to 14, called. universal. They are used for approx. assessment of pH solutions. The color change of the indicator is influenced by it. For two-color indicators, the higher , the less sharply the change in color, because. the absorption spectra of both forms overlap more and the color change becomes more difficult to detect. Usually use the same minimum (several drops of solution) amount of the indicator. The transition interval of many indicators depends on t-ry. So, it changes its color at room temperature in the pH range of 3.4-4.4, and at 100 ° C in the pH range of 2.5-3.3. It has to do with change. Colloidal particles present in the solution adsorb indicators, which leads to a complete change in its color. To avoid errors in the presence positively charged colloidal particles, indicators-bases should be used, and in the presence. negatively charged - acid indicators. Under normal conditions, the effect of dissolved CO 2 must be taken into account, especially when using indicators with pK ln > 4 (eg, methyl red, ). Sometimes CO 2 is previously removed by boiling or titrated with a solution in the absence of contact with. The influence of extraneous neutrals (salt effect) is manifested in the shift of indicators. In the case of acid indicators, the transition interval shifts to a more acidic region, and in the case of base indicators, to a more alkaline one. Depending on the nature of the solvent, the colors of the indicators, their pK ln and sensitivity change. Thus, methyl red in gives a color transition at higher H + values ​​than bromophenol blue, and vice versa in ethylene glycol solution. In water-methanol and water-ethanol solutions, the change in comparison with the aqueous medium is insignificant. In an alcoholic medium, acid indicators are more sensitive to H + than base indicators. Although when in non-toxic environments, usually k.t.t. is set potentiometrically using a glass indicator, they are also used (Table 3). Most often, for the weak, methyl red is used in or in anhydrous CH 3 COOH; at weak to-t- in DMF. The behavior of indicators in non-aqueous and aqueous media is similar. For example, for a weak to-you HIn in p-solvent SH can be written: HIn + SH D In- + SH 2 + . The mechanism of action of the indicators is the same as in, only in non-aqueous media they use the corresponding acidity scales (pH p, pA; see). They are also used in quality, changing color and intensity depending on pH and allowing titration of strongly colored and cloudy solutions.

For weak to-t apply t called. cloudiness indicators in-va, forming reversible, coagulating in a very narrow pH range (for example, isonitroacetyl-n-aminobenzene gives turbidity at pH 10.7-11.0). As you can use complexes with (see below); these complexes, collapsing, change the color of the solution in a narrow pH range. To determine org. to-t and in in the presence. a solution immiscible with it is used so-called. amphi-indicators, to-rye are acid indicators (eg, 00) with decomp. org. (e.g. ). These indicators are well sol. in org. p-parents, bad at; are highly sensitive. Adsorption indicators in the islands that can be adsorbed on the surface of the sediment and change color or intensity at the same time. These indicators are usually reversible and are used in precipitation. which the indicator is adsorbed. A large group of indicators (Table 4) adsorbed by the surface of the sediment with the formation of c contained in the sediment.


For example, rr pink colors, to-ry does not change when AgNO 3 is added. But at p-rum KBr, the precipitate adsorbs Ag +, which attach to themselves. The precipitate becomes red-violet. In c.t.t., when all Ag + is titrated, the color of the precipitate disappears and the solution becomes pink again. Inorg. adsorption indicators form a colored precipitate or complex from the titrant (as, for example, used as indicators CrO 4- and SCN - in ). as an adsorbent. indicators are also used nek-ry acid-base, oxidize.-restore. and complexometric. indicators, St. Islands to-rykh (acid, redox. potentials and stability of complexes with) in adsorbed. condition depend on the nature and on the surface of the sediment. Oxidation-reduction indicators - in-va, capable of changing color depending on the oxidation.-restore. r-ra potential. Used to establish K. t. t. oxidize-restore. and for colorimetric definitions okislit.-restore. potential (primarily in biology). Such indicators are, as a rule, in-va, to-rye themselves undergo or, and the oxidized (In Ox) and reduced (In Red) forms have different colors. For reversible oxidizing.-restore. indicators can be written: In Ox + ne D In Red, where n is a number. At potential E, the ratio of both forms of the indicator is determined by:
,
where E ln - real okislit.-restore. indicator potential, depending on the composition of the solution. The color transition interval is practically observed when the ratio / changes from 0.1 to 10, which at 25 °C corresponds to D E (in V) = E ln b (0.059/n). The potential corresponding to the sharpest color change is E ln . When choosing an indicator, take into account Ch. arr. values ​​E ln , coefficient molar redemption of both forms of the indicator and the potential of the solution at the equivalence point. At strong (K 2 Cr 2 O 7, KMnO 4, etc.), indicators are used that have relatively high Eln, for example, and its derivatives; with strong [Ti(III), V(II), etc.], indicators with relatively low Eln are used, for example, (Table 5).


Some in-va change their color irreversibly, for example, when destroyed with the formation of colorless. products, as under the action or naphthol blue-black under the action of BrO 3 . Complexometric indicators - in-va, forming with (M) colored complexes, differing in color from the indicators themselves. They are used to establish c.t.t. The stability of complexes with indicators (In) is less than that of the corresponding complexonates, therefore, in the c.t.t., indicators are displaced from complexes with . At the moment of color change at the equivalence point = and, therefore, pM = - lg K Mln , where pM = - lg[M] is called. transition point of the indicator, K Mln - stability of the complex with the indicator. The error at is due to the fact that a certain amount can be attached to the indicator, and not to the titrant. Naib. often use the so-called.

When conducting chemical process it is extremely important to follow the conditions 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, capable of changing the color of the solution depending on the 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 essential tool for control of technological processes in chemical, metallurgical, textile, food and other industries. AT agriculture with the help of indicators, they analyze and classify soils, establish the nature of fertilizers and required amount them for incorporation into the soil.

Distinguish acid-base, fluorescent, redox, adsorption and chemiluminescent indicators.

ACID-BASE (PH) INDICATORS

As is known from the theory 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, whose molecules 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, select two indicators 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 ethyl alcohol and add 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 excess solution is squeezed out of the paper, dried, cut into narrow strips and booklets. 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 displaying at certain pH values. visible light. 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.

Among the diversity organic matter there are special compounds that are characterized by color changes in different environments. Before the advent of modern electronic pH meters, indicators were indispensable "tools" for determining the acid-base indicators of the environment, and continue to be used in laboratory practice as auxiliary substances in analytical chemistry and also in the absence of the necessary equipment.

What are indicators for?

Initially, the property of these compounds to change color in various media was widely used to visually determine the acid-base properties of substances in solution, which helped to determine not only the nature of the medium, but also to draw a conclusion about the resulting reaction products. Indicator solutions continue to be used in laboratory practice to determine the concentration of substances by titration and allow you to learn how to use improvised methods in the absence of modern pH meters.

There are several dozens of such substances, each of which is sensitive to a rather narrow area: usually it does not exceed 3 points on the informativeness scale. Thanks to such a variety of chromophores and their low activity among themselves, scientists managed to create universal indicators that are widely used in laboratory and industrial conditions.

Most used pH indicators

It is noteworthy that in addition to the identification property, these compounds have a good dyeing ability, which allows them to be used for dyeing fabrics in the textile industry. From a large number The most famous and used color indicators in chemistry are methyl orange (methyl orange) and phenolphthalein. Most of the other chromophores are currently used in admixture with each other, or for specific syntheses and reactions.

methyl orange

Many dyes are named for their primary colors in a neutral environment, which is also characteristic of this chromophore. Methyl orange is an azo dye having a grouping - N = N - in its composition, which is responsible for the transition of the color of the indicator to red in and to yellow in alkaline. Azo compounds themselves are not strong bases, however, the presence of electron-donating groups (‒ OH, ‒ NH 2 , ‒ NH (CH 3), ‒ N (CH 3) 2, etc.) increases the basicity of one of the nitrogen atoms, which becomes able to attach hydrogen protons according to the donor-acceptor principle. Therefore, with a change in the concentration of H + ions in a solution, a change in the color of the acid-base indicator can be observed.

More on getting methyl orange

Get methyl orange in the reaction with the diazotization of sulfanilic acid C 6 H 4 (SO 3 H)NH 2 followed by a combination with dimethylaniline C 6 H 5 N(CH 3) 2 . Sulfanilic acid is dissolved in a sodium alkali solution by adding sodium nitrite NaNO 2, and then cooled with ice to carry out the synthesis at temperatures as close as possible to 0 ° C and poured hydrochloric acid HCl. Next, a separate solution of dimethylaniline in HCl is prepared, which is poured into the first solution when cooled, obtaining a dye. It is additionally alkalized, and dark orange crystals precipitate from the solution, which, after several hours, are filtered off and dried in a water bath.

Phenolphthalein

This chromophore got its name from the addition of the names of the two reagents that are involved in its synthesis. The color of the indicator is notable for its change in color in an alkaline medium with the acquisition of a raspberry (red-violet, raspberry-red) hue, which becomes colorless when the solution is strongly alkalized. Phenolphthalein can take several forms depending on the pH of the environment, and in strongly acidic environments it has an orange color.

This chromophore is obtained by the condensation of phenol and phthalic anhydride in the presence of zinc chloride ZnCl 2 or concentrated sulfuric acid H 2 SO 4 . In the solid state, phenolphthalein molecules are colorless crystals.

Previously, phenolphthalein was actively used in the creation of laxatives, but gradually its use was significantly reduced due to the established cumulative properties.

Litmus

This indicator was one of the first reagents used on solid carriers. Litmus is a complex mixture of natural compounds that is obtained from certain types of lichens. It is used not only as but also as a means for determining the pH of the medium. This is one of the first indicators that began to be used by man in chemical practice: it is used in the form of aqueous solutions or strips of filter paper impregnated with it. Litmus in the solid state is a dark powder with a slight ammonia odor. When dissolved in clean water the color of the indicator takes on a violet color, and when acidified, it turns red. In an alkaline environment, litmus turns blue, which allows it to be used as a universal indicator for general definition environment indicator.

It is not possible to accurately establish the mechanism and nature of the reaction that occurs when the pH changes in the structures of the litmus components, since it can include up to 15 different compounds, some of which may be inseparable active substances, which complicates their individual studies of chemical and physical properties.

Universal indicator paper

With the development of science and the advent of indicator papers, the establishment of environmental indicators has become much simpler, since now it was not necessary to have ready-made liquid reagents for any field research, which scientists and forensic scientists still successfully use. So, solutions were replaced by universal indicator papers, which, due to their wide spectrum of action, almost completely eliminated the need to use any other acid-base indicators.

The composition of the impregnated strips may vary from manufacturer to manufacturer, so an approximate list of ingredients may be as follows:

• phenolphthalein (0-3.0 and 8.2-11);
• (di)methyl yellow (2.9-4.0);
• methyl orange (3.1-4.4);
• methyl red (4.2-6.2);
• bromthymol blue (6.0-7.8);
• α-naphtholphthalein (7.3-8.7);
• thymol blue (8.0-9.6);
• cresolphthalein (8.2-9.8).

The packaging necessarily contains color scale standards that allow you to determine the pH of the medium from 0 to 12 (about 14) with an accuracy of one integer.

Among other things, these compounds can be used together in aqueous and aqueous-alcoholic solutions, which makes the use of such mixtures very convenient. However, some of these substances may be poorly soluble in water, so it is necessary to select a universal organic solvent.

Due to their properties, acid-base indicators have found their application in many fields of science, and their diversity has made it possible to create universal mixtures that are sensitive to a wide range of pH values.

In an acidic pH solution< 7, в нейтральной среде рН = 7, в щелочной рН >7. The lower the pH, the greater the acidity of the solution. At pH values ​​> 7, one speaks of the alkalinity of the solution.

There are various methods for determining the pH of a solution. Qualitatively, the nature of the solution medium is determined using indicators. Indicators are substances that reversibly change their color depending on the medium of the solution. In practice, litmus, methyl orange, phenolphthalein, and a universal indicator are most often used (Table 2).

table 2

Coloring of indicators in various solution media

The hydrogen index is very important for medicine, its deviation from normal values ​​even by 0.01 units indicates pathological processes in the body. With normal acidity, gastric juice has pH = 1.7; human blood has pH = 7.4; saliva - pH = 6.9.

Ion exchange reactions and conditions for their occurrence

Since electrolyte molecules in solutions decompose into ions, reactions in electrolyte solutions proceed between ions. Ion exchange reactions- these are reactions between ions formed as a result of the dissociation of electrolytes. The essence of such reactions is the binding of ions through the formation of a weak electrolyte. In other words, the ion exchange reaction makes sense and proceeds almost to the end if weak electrolytes (precipitate, gas, H 2 O, etc.) are formed as a result of it. If there are no ions in the solution that can bind to each other to form a weak electrolyte, then the reaction is reversible; equations for such exchange reactions are not written.

When recording ion exchange reactions, molecular, full ionic, and abbreviated ionic forms are used. An example of recording an ion exchange reaction in three forms:

K 2 SO 4 + BaCl 2 \u003d BaSO 4 + 2KCl,

2K + + SO 4 2– + Ba 2+ + 2Cl – = BaSO 4 + 2K + + 2Cl – ,

Ba 2+ + SO 4 2– \u003d BaSO 4.

Rules for compiling equations of ionic reactions

1. Formulas of weak electrolytes are written in molecular form, strong ones in ionic form.

2. For the reaction, solutions of substances are taken, therefore, even poorly soluble substances in the case of reagents are recorded in the form of ions.

3. If a poorly soluble substance is formed as a result of a reaction, then when writing the ionic equation, it is considered insoluble.

4. The sum of the charges of the ions on the left side of the equation must be equal to the sum of the charges of the ions on the right side.

Test on the topic “Theory of electrolytic dissociation. Ion exchange reactions»

1. The reaction that occurs when magnesium hydroxide is dissolved in sulfuric acid is described by the reduced ionic equation:

a) Mg 2+ + SO 4 2– = MgSO 4;

b) H + + OH - = H 2 O;

c) Mg(OH) 2 + 2H + = Mg 2+ + 2H 2 O;

d) Mg(OH) 2 + SO 4 2– = MgSO 4 + 2OH –.

2. Four vessels contain one liter of 1M solutions of the following substances. Which solution contains the most ions?

a) Potassium sulfate; b) potassium hydroxide;

c) phosphoric acid; d) ethyl alcohol.

3. The degree of dissociation does not depend on:

a) the volume of the solution; b) the nature of the electrolyte;

c) solvent; d) concentration.

4. Reduced ionic equation

Al 3+ + 3OH - \u003d Al (OH) 3

corresponds to the interaction:

a) aluminum chloride with water;

b) aluminum chloride with potassium hydroxide;

c) aluminum with water;

d) aluminum with potassium hydroxide.

5. An electrolyte that does not dissociate stepwise is:

a) magnesium hydroxide; b) phosphoric acid;

c) potassium hydroxide; d) sodium sulfate.

6. Weak electrolyte is:

a) barium hydroxide;

b) aluminum hydroxide;

c) hydrofluoric acid;

d) hydroiodic acid.

7. The sum of the coefficients in the brief ionic equation for the interaction of barite water and carbon dioxide is:

a) 6; b) 4; at 7; d) 8.

8. The following pairs of substances cannot be in a solution:

a) copper chloride and sodium hydroxide;

b) potassium chloride and sodium hydroxide;

c) hydrochloric acid and sodium hydroxide;

d) sulfuric acid and barium chloride.

9. A substance whose addition to water will not change its electrical conductivity is:

a) acetic acid; b) silver chloride;

c) sulfuric acid; d) potassium chloride.

10. How will the dependence of the incandescence of an electric bulb included in the circuit on time look like if the electrodes are immersed in a solution of lime water, through which carbon dioxide is passed for a long time?

a) Linear increase;

b) linear decrease;

c) first decrease, then increase;

d) first increase, then decrease.