Universal indicator in a neutral environment. Color indicators

There are various methods for determining the concentration (more precisely, activity) of hydrogen ions (and, accordingly, the concentration of hydroxide ions). One of the simplest (colorimetric) is based on the use acid-base indicators. Many organic acids and bases, which change their color in a certain narrow range of pH values, can serve as such indicators.

Indicators are weak acids or bases that have different colors in their undissociated and dissociated (ionic) forms.

Example.

1.Phenolphthalein is an acid which in molecular form (HJnd) is colorless at pH8.1. Phenolphthalein anions (Jnd -) at pH9.6 have a red-violet color:

H Jnd  H + + Jnd -

Colorless  red-violet

pH8.1 pH9.6

With a decrease in the concentration of H + ions and an increase in the concentration of OH ions, the molecular form of phenolphthalein becomes anionic due to the detachment of the hydrogen ion from the molecules and its binding to the hydroxide ion in water. Therefore, at pH9.6, the solution in the presence of phenolphthalein acquires a red-violet color. On the contrary, in acidic solutions at pH  8.1, the equilibrium shifts towards the molecular form of the indicator, which has no color.

2.Methyl orange is a weak base JndOH , which in molecular form at pH 4.4 has yellow. Jnd + cations at pH3.0 color the solution red:

JndOH  Jnd + + OH -

yellow  red

pH4.4 pH3.0

acid form indicator is called the form that prevails in acid solutions, and basic form - the one that exists in basic (alkaline) solutions. In a certain range of pH values ​​in the solution, a certain amount of both forms of the indicator can be simultaneously in equilibrium, as a result of which a transitional color of the indicator occurs - this is the pH range of the indicator color transition, or simply indicator transition interval.

Table 1 shows the transition intervals of some commonly used indicators.

Table 1

Acid-base indicators

For a quick determination of pH, it is also convenient to use a solution of a universal indicator, which is a mixture of various indicators and has a large transition range (pH values ​​from 1 to 10). On the basis of a universal indicator, the industry produces special paper tapes for determining the pH of solutions by comparison with a special scale for changing their color under the action of the test solution.

In the colorimetric method, standard buffer solutions are used to accurately determine pH, the pH value of which is precisely known and constant.

Buffer solutions are mixtures of weak acids or bases with their salts. Such mixtures maintain a certain pH value both when diluted and when small amounts of strong acids or alkalis are added.

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.

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.

Acid-base indicators (otherwise they are called indicators) are substances that change their color depending on the environment where they are located. Usually, such substances are weak acids or weak bases. When dissolved in water, they dissociate weakly, forming ions. As an example consider an indicator, which is a weak acid having the formula general view When dissolved in water, the following equilibrium is established between this weak acid and its conjugate base:

An acid is used as an indicator, the color of which differs markedly from the color of its conjugate base. At low values, the ion concentration in the solution is high, and hence the equilibrium position is shifted to the left. Under these conditions, the equilibrium solution has color A. At high

Rice. 8.1. Phenolphthalein.

At pH values, the concentration in the solution is low, and, therefore, the equilibrium position is shifted to the right, which means that the equilibrium solution has color B.

An example of an indicator in which an equilibrium of this type is established in an aqueous solution is phenolphthalein (Fig. 8.1). Phenolphthalein is a colorless weak acid which, when dissolved in water, forms pink anions. In an acidic medium, the equilibrium between an acid and its anion is shifted to the left. The concentration of anions is so low that their pink color is imperceptible. However, in an alkaline (basic) medium, the equilibrium shifts to the right, and the concentration of anions becomes sufficient to reveal their pink color.

If we apply the law of mass action to the equilibrium of an indicator in an aqueous solution, then in the general case for an indicator, which is a weak acid, we obtain the following expression for the equilibrium constant:

The value is called the dissociation constant of the indicator.

The color of the indicator changes from A to B at some color transition point. At this point

Therefore, from equation (5)

The pH value of the solution at the color transition point of the indicator is denoted by . Thus, it represents a pH value at which half of the indicator is in the form of an acid, and half is in the form of its conjugate base.

Indicator color change range

At low values, the indicator, which is a weak acid, is almost completely in the form and, therefore, the color of this form predominates in the solution. As the intensity of color A, inherent in the form, decreases, and the equilibrium described by equation (4) shifts to the right. Thus, the intensity of color B, the inherent form increases. The observed color change from A to B actually occurs over a range of values. An indicator is most effective when a distinct observable

Table 8.5. Indicators

my color change occurs in a narrow range of values ​​For most indicators, this range is within the limits of the value (Table 8.5).

A universal indicator is a mixture of indicators that gives a gradual change in color over a wide range of changes. If you add a few drops of a universal indicator to the solution, then you can approximately determine it by the color of the solution.

Acid-base titrations

Acid-base titration is an experimental technique for determining the concentration of an acid or base, used primarily in quantitative chemical analysis. Typically, an acid of known concentration is gradually added from a buret to an alkaline solution of unknown concentration in a conical flask. The equivalence point of the titration is reached when exactly the stoichiometric amount of acid is added to the base. At this point, all the alkali is neutralized and there is neither excess acid nor excess base in the solution. The solution consists only of salt and water. For example, when adding of hydrochloric acid with the concentration of the sodium hydroxide solution with the concentration, the titration equivalence point is reached at the moment when exactly hydrochloric acid is added to the solution. This follows from the stoichiometric equation

In acid-base titrations, indicators are often used to determine the equivalence point. However, the equivalence point can also be determined potentiometrically using a meter or conductometric methods (see Chap. 10).

Let us now assume that the titration is carried out by adding a base to an acid. If you build a graph of the change in solution as the volume of added base increases, then depending on whether the acid and base are strong or weak, four types of curves will be obtained. These four types of titration curves are shown in Fig. 8.2. It should be noted that reaching the equivalence point is characterized by a sharp increase. The only exception in this respect is the titration of a weak acid with a weak base. If for

To determine the equivalence point of an acid-base titration, you have to use an indicator, then it should be chosen so that the pH range in which the color change occurs falls on the vertical part of the titration curve. This provides a sharp change in the color of the indicator at the moment the equivalence point of the titration is reached.

Titration of a strong acid with a strong base. For example,

The vertical part of the curve of this titration falls on the range of pH changes from 4 to 10. Therefore, at the titration equivalence point, adding another drop of base to the acid causes an increase in pH by 6 units at once. Therefore, for such titrations, indicators can be used that have a color change range between pH 4 and 10. Examples of such indicators are methyl red and phenolphthalein. Note that if methyl orange is used as an indicator for the titration of a strong acid with a strong base, then the color change is not so sharp.

Titration of a strong acid with a weak base. For example,

The vertical part of the curve of this titration falls on the range of pH changes from 4 to 8. Convenient indicators for it are methyl red or

bromothymol blue, but not phenolphthalein, since its range of color change falls on the flat part of the titration curve.

Titration of a weak acid with a strong base. For example,

The vertical part of this titration curve falls within the range of pH values ​​from 6.5 to 11. Therefore, phenol red or phenolphthalein are convenient indicators for it. Indicators with a color change range below pH 6, like methyl orange, are not suitable for this titration,

Rice. 8.2. Titration curves of 25.00 cm3 acid with a concentration of 0.10 mol/dm3 with a base with a concentration of 0.10 mol/dm3: a - titration of a strong acid with a strong base; b - titration of a strong acid with a weak base; c - titration of a weak acid with a strong base; d - titration of a weak acid with a weak base. I-phenolphthalein, II-methyl orange.

since their range of color change falls on the flat part of the titration curve and, therefore, does not allow an accurate detection of the equivalence point.

Titronation of a weak acid with a weak base. For example,

This type of titration is characterized by the absence of a sudden change in pH at the moment the equivalence point is reached. Changes in pH occur smoothly over the entire range of accepted values. Therefore, it is not possible to select an indicator for this type of titrations.

So let's do it again

1. A strong electrolyte, when dissolved or in a molten state, is completely ionized.

2. A weak electrolyte, when dissolved or in a molten state, dissociates into ions only partially.

3. The Ostwald dilution law relates the dissociation constant of an electrolyte to its degree of dissociation a and concentration c:

4. According to the Bronsted-Lowry theory, an acid is a substance that donates (gives off) protons, and a base is a substance that accepts (attaches) protons.

5. A strong acid has a weak conjugate base.

6. A weak acid has a strong conjugate base.

7. Amphoteric substance can react both as an acid and as a base.

8. Lewis acid is a substance capable of accepting an electron pair provided by a base.

9. A Lewis base is a substance that has an unshared electron pair.

10. where is the acid dissociation constant.

11. , where is the dissociation constant of the base.

12. , where is the ionic product of water.

16. The equilibrium in an aqueous solution of an indicator, which is a weak acid, is determined by the equation

17. The equivalence point of an acid-base titration is reached when a stoichiometric amount of base is added to the acid.

18. The range of color change of the visual indicator should fall on the vertical part of the titration curve.

INDICATORS(from lat. indicator - pointer) - substances that allow you to monitor the composition of the environment or the flow 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 may 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, began to be 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 mold (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; solution of this complex organic compound has Blue colour, 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: to a solution containing the analyzed cations and an indicator, a stronger complexing agent, compared to the indicator, is added dropwise, 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) different color 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

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 for some outward signs: stopping the evolution of gas bubbles, changing the color of the solution, precipitation or, conversely, the transition of one of the reaction components into the solution, etc. In most cases, to determine the end of the reaction, auxiliary reagents are used, the so-called indicators, which are usually injected 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 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, they analyze and classify soils, establish the nature of fertilizers and the required amount for their application to 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.

To measure the pH value, use various methods. 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.

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 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.

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.

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.

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 end point titration, a glow occurs.