Color indicators. Methods for determining the pH of solutions. Indicators. Change in color of indicators depending on pH
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Every schoolchild is familiar with litmus - it is used to determine the acidity of the environment. This substance is an acid-base indicator, that is, it has the ability to reversibly change color depending on the acidity of the solution: in an acidic environment, litmus becomes red, and in an alkaline environment, it turns blue. In a neutral environment, the litmus color violet is a combination of equal amounts of blue and red. Although litmus has been serving people faithfully for several centuries, its composition has not been fully studied. This is not surprising: after all, lac mousse is a complex mixture of natural compounds. It was already known in Ancient Egypt and Ancient Rome, where it was used as a violet paint - a substitute for expensive purple. Then the recipe for making litmus was lost. Only at the beginning of the 14th century. In Florence, the purple dye Orseille, identical to litmus, was rediscovered, and the method of its preparation was kept secret for many years.
When moving from an acidic to an alkaline environment, the color of litmus changes from red to blue.
Litmus was prepared from special types of lichens. The crushed lichens were moistened, and then ash and soda were added to this mixture. The thick mass prepared in this way was placed in wooden barrels, urine was added and kept for a long time. Gradually, the solution acquired a dark blue color. It was evaporated and in this form used for dyeing fabrics. In the 17th century, the production of orseili was established in Flanders and Holland, and lichens, which were brought from the Canary Islands, were used as raw materials.
A coloring substance similar to orseille was isolated in the 17th century. from heliotrope - a fragrant garden plant with dark purple flowers.
Famous physicist and chemist of the 17th century. Robert Boyle wrote about heliotrope: “The fruits of this plant yield a juice, which, when applied to paper or material, first has a fresh, bright green color, but suddenly changes it to purple. If the material is soaked in water and squeezed out, the water turns wine-colored; Such types of dye (they are usually called “tournesol”) are available from pharmacists, grocery stores and other places, which are used to color jelly, or other substances, as one wishes.” Since that time, orseil and heliotrope have been used in chemical laboratories. And only in 1704 the German scientist M. Valentin called this paint litmus.
Today, to produce litmus, crushed lichens are fermented in solutions of potash (potassium carbonate) and ammonia, then chalk or gypsum is added to the resulting mixture. It is believed that the coloring substances of litmus are indophenols, which exist in cationic form in an acidic environment, and in anionic form in an alkaline environment, for example:
In some countries, paint similar to litmus was also obtained from other plants. The simplest example is beet juice, which also changes color depending on the acidity of the medium.
In a strongly acidic environment, the methyl orange indicator is red, in a weakly acidic and neutral environment it is orange, and in an alkaline environment it is yellow.
Methyl orange in alkaline medium.
In the 19th century Litmus has been replaced by stronger and cheaper synthetic dyes, so the use of litmus is limited to only a rough determination of the acidity of the medium. For this purpose, strips of filter paper soaked in litmus solution are used. In analytical practice, the use of litmus is limited by the fact that as it becomes half-acidified, it changes color gradually, and not in a narrow pH range, like many modern indicators. Litmus was replaced in analytical chemistry by lakmoid - a resorcinol blue dye, which differs from natural litmus in structure, but is similar to it in color: in an acidic environment it is red, and in an alkaline environment it is blue.
As the pH increases to 8-8.5, the color of phenolphthalein changes from colorless to crimson.
Today, several hundred acid-base indicators are known, artificially synthesized since the mid-19th century. You can get acquainted with some of them in the school chemistry laboratory. The methyl orange indicator (methyl orange) is red in an acidic environment, orange in a neutral environment, and yellow in an alkaline environment. A brighter color range is characteristic of the thymol blue indicator: in an acidic environment it is crimson-red, in a neutral environment it is yellow, and in an alkaline environment it is blue. The indicator phenolphthalein (it is sold in pharmacies under the name “purgen”) is colorless in acidic and neutral environments, and has a crimson color in alkaline environments. Therefore, phenol-phthalein is used only to determine the alkaline environment. Depending on the acidity of the medium, the brilliant green dye also changes color (a hundred-alcohol solution is used as a disinfectant - “green”). In order to check this, you need to prepare a diluted solution of brilliant green: pour a few milliliters of water into a test tube and add one or two drops of the pharmaceutical preparation to it. The solution will acquire a beautiful green-blue color. In a strongly acidic environment, its color will change to yellow, and in a strongly alkaline solution it will become discolored.
In an acidic solution pH< 7, в нейтральной среде рН = 7, в щелочной рН >7. The lower the pH, the more acidic the solution. At pH values > 7, the solution is said to be alkaline.
There are various methods for determining the pH of a solution. The nature of the solution environment is determined qualitatively using indicators. Indicators are substances that reversibly change their color depending on the solution environment. In practice, litmus, methyl orange, phenolphthalein and a universal indicator are most often used (Table 2).
table 2
Coloring of indicators in various solution environments
The pH value is very important for medicine; its deviation from normal values by even 0.01 units indicates pathological processes in the body. With normal acidity, gastric juice has a 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 disintegrate into ions, reactions in electrolyte solutions also occur between the 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 completion if it results in the formation of weak electrolytes (precipitate, gas, H 2 O, etc.). If there are no ions in the solution that can bond with 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 writing an ion exchange reaction in three forms:
K 2 SO 4 + BaCl 2 = BaSO 4 + 2KCl,
2K + + SO 4 2– + Ba 2+ + 2Cl – = BaSO 4 + 2K + + 2Cl – ,
Ba 2+ + SO 4 2– = BaSO 4 .
Rules for composing equations of ionic reactions
1. The formulas of weak electrolytes are written in molecular form, and the formulas of strong electrolytes in ionic form.
2. For reactions, solutions of substances are taken, so even slightly soluble substances in the case of reagents are written in the form of ions.
3. If a slightly 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 ion charges on the left side of the equation must be equal to the sum of the ion charges 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 abbreviated 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 substances listed below. 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) volume of solution; b) the nature of the electrolyte;
c) solvent; d) concentrations.
4. Abbreviated ionic equation
Al 3+ + 3OH – = 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. A weak electrolyte is:
a) barium hydroxide;
b) aluminum hydroxide;
c) hydrofluoric acid;
d) hydroiodic acid.
7. The sum of the coefficients in the short ionic equation for the interaction of barite water and carbon dioxide is equal to:
a) 6; b) 4; at 7; d) 8.
8. The following pairs of substances cannot be present in 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. What will a graph of the intensity of an electric light bulb connected to a circuit versus 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 decreasing, then increasing;
d) first increasing, then decreasing.
pH can be approximately estimated using indicators, measured accurately with a pH meter, or determined analytically by performing acid-base titration.
1. For a rough estimate of the concentration of hydrogen ions, acid-base indicators- organic dye substances, the color of which depends on the pH of the environment. The most well-known indicators include litmus, phenolphthalein, methyl orange (methyl orange) and others. Indicators can exist in two differently colored forms - either acidic or basic. The color change of each indicator occurs in its own acidity range, usually 1-2 units (Table 3.1). Their advantage is low cost, speed and clarity of research.
This method is not accurate enough, requires the introduction of salt and temperature corrections, and produces a significant error with very low mineralization of the water under study (less than 30 mg/l) and when determining the pH of colored and turbid waters. The method cannot be used for media containing strong oxidizing or reducing agents. It is usually used in field conditions and for approximate determinations.
Change in color of acid-base indicators
depending on the pH of the solution
2. To expand the working range of pH measurements, use the so-called universal indicator, which is a mixture of several indicators. The universal indicator changes color sequentially from red through yellow, green, blue to violet when moving from an acidic region to an alkaline one. The indicator changes color in the pH range 1.0-10.0 (Table 3.2).
Changing the color of the universal indicator
depending on the pH of the solution
3. Using a special device - pH meter- allows you to measure pH over a wider range and more accurately (up to 0.01 pH units) than using indicators. The ionometric method for determining pH is based on measuring the EMF of a galvanic circuit with a millivoltmeter ionometer, including a special glass electrode, the potential of which depends on the concentration of H + ions in the surrounding solution. The method is convenient and highly accurate, especially after calibrating the indicator electrode in a selected pH range; it allows you to measure the pH of opaque and colored solutions and is therefore widely used.
A glass electrode is a glass tube with a blown ball at its end with a very thin wall, into which a suspension of AgCl in an HCl solution is poured and a silver wire is immersed. Thus, inside the tube with the ball there is a silver chloride electrode. To measure pH, a glass electrode is immersed in the solution being tested (thereby without introducing any foreign substances into it). A reference electrode is immersed in the same solution directly or through an electrolytic switch. In the resulting system, the transfer of electrons from the silver chloride electrode to the reference electrode, which occurs under the influence of a directly measured potential difference, is inevitably accompanied by the transfer of an equivalent number of protons from the inside of the glass electrode to the test solution. If we consider the concentration of H + ions inside the glass electrode constant, then the measured emf is a function of only the activity of hydrogen ions, i.e. pH of the test solution.
4. Analytical volumetric method- acid-base titration - also gives accurate results for determining the acidity of solutions. A solution of known concentration (titrant) is added dropwise to the test solution. When they are mixed, a chemical reaction occurs. The equivalence point - the moment when there is exactly enough titrant to completely complete the reaction - is recorded using an indicator. Next, knowing the concentration and volume of the added titrant solution, the acidity of the solution is calculated.
INDICATORS(Late Lat. indicator - pointer), chemical. substances that change color or form a precipitate when the substance changes. component in the solution. Indicate a certain state of the system or the moment when this state is reached. There are reversible and irreversible indicators. A change in the color of the former when the state of the system changes (for example, when the pH of the medium changes) could be. repeated many times. Irreversible indicators are subject to irreversible chemistry. transformations, for example, with BrO 3 - are destroyed. Indicators that are introduced into the solution under study are called. internal, in contrast to external, the procedure with which 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 soaked in an indicator, or they are mixed on a white porcelain plate with a drop of indicator. AND Indicators are most often used to determine the end of a population. chem. districts, ch. arr. end point (t.t.t.). In accordance with titrimetric methods distinguish between acid-base, adsorption, oxidation-reduction. and complexometric. indicators. are r-rime org compounds, which change their color or depending on H + (pH of the environment). Appl. to establish the end of the circuit between the circuits and (including at) or other circuits, if H + are involved in them, as well as for colorimetric. determining the pH of water solutions. Naib. important ones are given in table. 1. The reason for the change in color of 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 has a weak value HIn, then in the aqueous solution the following occurs: 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 form 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. of this process K ln = / called. indicator. The color of the solution depends on the ratio /, which is determined by the pH of the solution.It is considered that the color of one form of the indicator is noticeable if it is 10 times higher than the other form, i.e. if the ratio / = /K ln is 0.1 or 10. A change in the color of the indicator is noted in the region pH = pK ln b 1, which is called. indicator transition interval. Change in max. clearly when = and K ln = [H 3 O] +, i.e. at pH = pK ln. The pH value at which it usually ends is called. pT 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 coincide with the pH of the indicator used, which leads to the so-called. indicator error. If an excess of untitrated weak or compound remains in the c.t.t., the error is called. resp. basic or acidic.
Indicator sensitivity - (v/l) determined (in this case H + or OH -
) at the point of max. abrupt color transition. There are: indicators sensitive to acids with a transition interval in the region of alkaline pH values (for example, thymolphthalein); sensitive indicators with a transition interval in the acidic region (like dimethyl yellow, etc.); neutral indicators, the transition interval of which is approx. pH 7 (neutral red, etc.). AND Indicators come in one or two colored shapes; such indicators are called resp. single-color and two-color. Naib. a clear change in color would be observed in those indicators whose acidic and basic forms are colored complementary. colors. However, such indicators do not exist. Therefore, by adding, the colors of both forms are changed accordingly. Thus, for 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 if you use a mixture of two indicators, the colors of which complement the other. friend. Such indicators are called mixed (Table 2).
Mixtures of indicators that continuously change their color over the entire pH range from 1 to 14 are called. universal. They are used for approx. assessment of pH of solutions.
The color change of the indicator is influenced by it. For two-color indicators, the higher , the less dramatic the color change, because The absorption spectra of both forms overlap each other to a greater extent and the color change becomes more difficult to detect. Usually the same minimum (several drops of solution) quantity of indicator is used.
The transition interval of many indicators depends on the temperature. Thus, it changes 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. This is due to change.
The colloidal particles present in the solution adsorb indicators, which leads to a complete change in its color. To eliminate errors in the presence. positively charged colloidal particles, base indicators should be used, and if present. negatively charged - acid indicators.
Under normal conditions, it is necessary to take into account the influence of dissolved CO 2, especially when using indicators with pK ln > 4 (for example, methyl red). Sometimes CO 2 is first removed by boiling or the solution is titrated in the absence of contact with.
The influence of extraneous neutrals (salt effect) is manifested in a 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 region.
Depending on the nature of the solvent, the colors of the indicators, their pK ln and sensitivity change. Thus, methyl red gives a color transition at higher H + values than bromophenol blue, and in ethylene glycol solution it is the other way around. In water-methanol and water-ethanol solutions, the change compared to the aquatic environment is insignificant. In an alcoholic environment, acid indicators are more sensitive to H + than base indicators.
Although in non-natural environments it is usually established potentiometrically using a glass indicator, it is also used (Table 3).
Most often, for the weak, methyl red is used in or in anhydrous CH 3 COOH; for weak acids - in DMF.
The behavior of indicators in non-aqueous and aquatic environments is similar. For example, for weak values HIn in the solution SH you can write: HIn + SH D In - + SH 2 + . The mechanism of action of 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 as products that change color and intensity depending on pH and allow the titration of highly colored and turbid solutions.
For weak ones, the so-called are used. indicators of turbidity, forming reversible, coagulating in a very narrow pH range (for example, isonitroacetyl-n-aminobenzene produces turbidity at pH 10.7-11.0). You can use complexes with (see below); These complexes, when destroyed, change the color of the solution in a narrow pH range.
To determine org. to-t and in present. The so-called solvent that does not mix with it is used. amphibious indicators, which are indicator acids (eg, 00) with decomp. org. (eg, ). These indicators are well soluble. in org. r-retailers, bad in; are highly sensitive.
Adsorption indicators are substances that can be adsorbed on the surface of the sediment and change color or intensity. These indicators are usually reversible and are used in sedimentation. First of all, they are adsorbed by the sediment, identical to those that are part of the sediment itself, after what the indicator is adsorbed. A large group of indicators (Table 4) are adsorbed by the surface of the sediment with the formation of s contained in the sediment.
For example, the solution is pink, which does not change when AgNO 3 is added. But with a solution of KBr, the precipitate that falls out adsorbs Ag +, which is added to itself. The precipitate becomes red-violet. At the c.t.t., when all Ag + have been titrated, the color of the precipitate disappears and the solution becomes pink again.
Inorg. adsorbts. indicators form a colored precipitate or complex from the titrant (such as, for example, CrO 4 used as indicators- and SCN - in). As an adsorbts. Some indicators are also used: acid-base, oxidation-reduction. and complexometric. indicators, properties of which (acid, oxidation-reduction potentials and stability of complexes with) in the adsorbir. condition depend on the nature and surface of the sediment.
Oxidation-reduction indicators - substances that can change color depending on oxidation-reduction. solution potential. Used to establish the oxidation-reduction temperature. and for colorimetric determination of oxidation-reduction. potential (primarily in biology). Such indicators are, as a rule, substances that themselves are exposed to or, and the oxidized (In Ox) and reduced (In Red) forms have different colors.
For reversible oxidation-reduction. 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 oxidation-reduction. 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 equal to Eln. When choosing an indicator, take into account ch. arr. values Eln, coefficient. molar extinction of both forms of the indicator and the potential of the solution at the equivalence point. When 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 substances change their color irreversibly, for example, when they are destroyed with the formation of colorless. products, such as under the influence or naphthol blue-black under the influence of BrO 3.
Complexometric indicators are substances that form colored complexes with (M), differing in color from the indicators themselves. They are used to establish the quality of the indicator. The stability of complexes with indicators (In) is less than that of the corresponding complexonates,
therefore, in K.T.T. they displace indicators from complexes with. At the moment of color change at the equivalence point = and, therefore, рМ = - log K Mln, where рМ = - log[M] is called. transition point of the indicator, K Mln - stability of the complex with the indicator. The error 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 carrying out a chemical process, it is extremely important to monitor the conditions of the reaction or determine whether it has reached completion. Sometimes this can be observed by some external signs: the cessation of the release of gas bubbles, a change in the color of the solution, the formation of a precipitate 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, such as called indicators, which are usually introduced into the analyzed solution in small quantities.
Indicators are chemical compounds that can change the color of a solution depending on environmental conditions, without directly affecting the test solution or the direction of the reaction. Thus, acid-base indicators change color depending on the pH of the environment; 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 monitoring technological processes in the chemical, metallurgical, textile, food and other industries. In agriculture, with the help of indicators, soils are analyzed and classified, the nature of fertilizers and the required quantity to be applied to the soil are determined.
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 ones - anions. Water also dissociates to a very small extent into hydrogen ions, which are positively charged, and hydroxyl ions, which are negatively charged:
The concentration of hydrogen ions in a solution is indicated by the symbol.
If the concentration of hydrogen and hydroxyl ions in the solution is the same, then such solutions are neutral and pH = 7. At a concentration of hydrogen ions corresponding to pH from 7 to 0, the solution is acidic, but if the concentration of hydroxyl ions is higher (pH = from 7 to 14), the solution alkaline.
Various methods are used to measure the pH value. Qualitatively, the reaction of a 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 environment.
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 environment. They are used in titrimetric analysis for 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 increase the accuracy of pH measurements, then mixed indicators are used. To do this, select two indicators with close pH intervals of color transition, having additional colors in this interval. Using such a mixed indicator, determinations can be made with an accuracy of 0.2 pH units.
Universal indicators that can change color many times over a wide range of pH values are also widely used. 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 pH color transition intervals, designed in such a way that a noticeable color change occurs when the pH of the medium changes.
For example, the industrially produced universal indicator PKS 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 - crimson, 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 a few drops are added to the test solution. The pH value is determined by the change in color of the solution. 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 indicator papers rather than indicator solutions. The latter are prepared as follows: filter paper is passed through a standard indicator solution, the excess solution is squeezed out of the paper, dried, cut into narrow strips and bound into booklets. To carry out the test, the indicator paper is dipped into the test solution or one drop of the solution is placed on a strip of indicator paper and the 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 fluorescence of a solution or change its color or shade.
This property is used for acid-base titration of oils, cloudy and highly colored solutions, since conventional indicators are unsuitable for these purposes.
Work with fluorescent indicators is carried out by illuminating the test solution with ultraviolet light.
Indicator | pH range of fluorescence change (in ultraviolet light) | Fluorescence color change |
4-Ethoxyacridone | 1,4-3,2 | Green - blue |
2-Naphthylamine | 2,8-4,4 | Increase in violet fluorescence |
Dimetnlnaphteyrodine | 3,2-3,8 | Lilac - orange |
1-Naftilamnn | 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-Naphthylaminesulfonamide | 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 a solution depending on the value of the redox potential. They are used in titrimetric methods of analysis, as well as in biological studies for the colorimetric determination of redox potential.
Indicator | Normal redox potential (at pH=7), V | Coloring the solution | |
oxidative form | restored form | ||
Neutral red | -0,330 | Red-violet | Colorless |
Safranin T | -0,289 | Brown | Colorless |
Potassium indigomonosulfonate | -0,160 | Blue | Colorless |
Potassium indigodisulfonate | -0,125 | Blue | Colorless |
Potassium indigotrisulfonate | -0,081 | Blue | Colorless |
Potassium indigo tetrasulfonate | -0,046 | Blue | Colorless |
Toluidine blue | +0,007 | Blue | Colorless |
Tnonin | +0,06 | Purple | Colorless |
Sodium o-Cresolindophenolate | +0,195 | Reddish blue | Colorless |
Sodium 2,6-Dnchlorophenolindophenolate | +0,217 | Reddish blue | Colorless |
Sodium m-bromophenolindophenolate | +0,248 | Reddish blue | Colorless |
Diphenylbenzidine | +0.76 (acidic solution) | Purple | Colorless |
ADSORPTION INDICATORS
Adsorption indicators- substances in the presence of which a change in the color of the precipitate formed during titration by precipitation occurs. Many acid-base indicators, some dyes and other chemical compounds are capable of changing the color of a precipitate at a certain pH value, which makes them suitable for use as adsorption indicators.
Indicator | Ion to be detected | Ion precipitator | Color change |
Alizarin red C | Yellow - pink-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 that can emit 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.