Solutions for changing the color of indicators

The purpose of the work: determination of the active acidity of solutions by changing the color of various indicators.

Reagents: distilled water, solutions of hydrochloric acid HCl and sodium hydroxide NaOH concentration of 0.1 mol / l, solutions of phenolphthalein (alcohol solution), methyl orange, universal indicator and universal litmus papers.

Equipment: rack with test tubes, pipettes.

Method of work

To familiarize yourself with the color changes of indicators in various media, pour 5 ml of distilled water, hydrochloric acid and sodium hydroxide at a concentration of 0.1 mol/l into three test tubes.

Add 2 drops of methyl orange solution to each tube. Shake the contents of the tube or mix with a glass rod and compare the colors of the indicators in neutral, alkaline and acid solutions.

Repeat the experiment with the universal indicator, phenolphthalein and indicator litmus papers.

How does the color of solutions change? Record the observations in Table. 4.1.

Table 4.1

Results of changing the color of the indicators

Draw conclusions about the change in the color of the indicator in various environments.

Laboratory work 4.2. Determination of concentration

Alkali solution by titration

The purpose of the work: to determine the concentration of sodium hydroxide solution NaOH by titration with a solution of hydrochloric acid.

Reagents: distilled water, solutions of hydrochloric acid HCl (or HNO 3) and sodium hydroxide NaOH (or KOH) concentrations
0.1 mol/l, solutions of phenolphthalein (alcohol solution) or methyl orange.

Equipment: 50–100 ml conical flasks, 10 ml pipettes, 25 ml burettes, 50–100 ml beakers, droppers for indicators, measuring cups or measuring tubes, glass rods, filter paper.

Method of work

For operation, it is necessary to have a solution of hydrochloric acid HC1 with a concentration of 0.1 mol / l. During titration, a neutralization reaction occurs.

HCl + NaOH \u003d NaCl + H 2 O.

In order to fix the moment of neutralization, resort to the help of indicators (phenolphthalein or methyl orange).

Begin the titration by filling the burette with a solution of hydrochloric acid HCl with a concentration of 0.1 mol/l above zero division.

Pipette 10 ml of NaOH sodium hydroxide solution and pour into a 100 ml conical flask. Enter 2-3 drops of methyl orange indicator there.

Proceed to titration: lower the acid solution from the buret in small portions, approximately 0.2 ml each, into the alkali solution, which is continuously stirred.

The area of ​​the solution into which the acid enters turns pink, turning yellow when shaken. When the pink in the solution begins to turn yellow, start slowly draining the acid solution by 0.1 ml and continue until the solution in the flask takes on a persistent pinkish color from one drop. This completes the titration.

Determine by divisions of the burette the volume of the acid solution used to neutralize the alkali. The titration is repeated two more times.

The volume of alkali taken and the values ​​obtained for the volume of acid are recorded in Table. 4.2.

Table 4.2

Calculate the molar concentration of sodium hydroxide according to the formula (4.1), the resulting value is recorded in table. 4.2. They make a conclusion.

Laboratory work 4.3. Determination of the concentration of a solution

Acid titration

The purpose of the work: to determine the concentration of acid by titration.

Reagents: distilled water, solutions of hydrochloric acid HCl (or HNO 3) and sodium hydroxide NaOH (or KOH) with a concentration of 0.1 mol / l, solutions of phenolphthalein (alcohol solution) or methyl orange.

Equipment: 50–100 ml conical flasks, 10 ml pipettes, 25 ml burettes, 50–100 ml beakers, droppers for indicators, glass rods, filter paper.

Method of work

For operation, it is necessary to have a solution of sodium hydroxide NaOH with a concentration of 0.1 mol / l.

As in work 4.2. prepare a burette filled with alkali solution. Pipette 10 ml of the acid solution and pour into a 100 ml conical flask; add 2-3 drops of phenolphthalein indicator there. Place the flask on white paper under the burette.

Proceed to titration: lower the alkali solution from the burette in small portions, approximately 0.2 ml each, into the acid solution, which is continuously stirred.

The area of ​​the solution into which the alkali enters turns pink, which disappears when shaken. When the pink color in the solution begins to fade slowly, start to lower the acid solution by 0.1 ml and continue until the solution in the flask from one drop appears a weak but fairly stable color of the solution. This completes the titration.

Determine by divisions of the burette the volume of the alkali solution used to neutralize the acid.

The titration is repeated two more times, each time starting from zero division of the burette.

The taken volume of acid and the obtained values ​​of the volume of alkali are recorded in table. 4.3.

Table 4.3

Results of experiment and calculation

Calculate the content of hydrochloric acid according to the formula (4.1), the value is recorded in table. 4.3. They make a conclusion.

Questions for self-preparation and control

1. What is titration? How is this operation performed?

2. What solution is called standard?

3. What is the equivalence point and how is it fixed?

4. Why is there a sharp change in the color of the indicator at the equivalence point?

5. Why are indicators used when titrating a solution?

6. What is the transition area of ​​the indicator?

7. In what medium do methyl orange and phenolphthalein change their color?

8. In what volume ratios do solutions of the same and different molar concentrations react?

Similar information.

Methyl orange lower limit p. H 3. 1 Red upper limit p. H 4. 4 ↔ Yellow 4 - (4-dimethylaminophenylazo) sodium benzenesulfonate

Phenolphthalein Getting into acid for others is a failure, But he will endure without sighs, without crying. But in the alkalis of phenolphthalein, not life will come, but solid raspberries!

Oxoacids of chlorine Ca r. H 0.1 M Hypochlorous HCl solution. O 2.9*10 -8 7.2 2.8 Chloride HCl. O 2 1.1*10 -2 2 1.5 0.1 -1 1.0*1010 -10 1.0 Chloric acid HCl. O 3 Chloric HCl. O 4

R. Ka of some weak acids (HO) n. XO formula p. Ka HCl. O 7, 2 HCl. O 2 2.0 HBr. O 8.7 HNO 2 3. 3 HJO 11.0 H 2 SO 3 1.9 H 3 As. O 3 9, 2 H 2 Se. O 3 2, 6 H 4 Ge. O 4 8.6 H 3 PO 4 2. 1 H 3 As. O 4 2, 3

R. Ka of some weak acids (HO) n. XO formula p. Ka HCl. O 7, 2 HCl. O 2 2.0 H 2 CO 3 6.3 HBr. O 8.7 HNO 2 3. 3 HJO 11.0 H 2 SO 3 1.9 H 3 As. O 3 9, 2 H 2 Se. O 3 2, 6 H 4 Ge. O 4 8.6 H 3 PO 4 2. 1 H 3 As. O 4 2, 3

R. Ka of some weak acids (HO) n. XO formula p. Ka HCl. O 7, 2 HCl. O 2 2.0 H 2 CO 3 6.3 HBr. O 8, 7 HNO 2 3, 3 H 3 PO 3 1, 8 HJO 11, 0 H 2 SO 3 1, 9 H 3 PO 2 2, 0 H 3 As. O 3 9, 2 H 2 Se. O 3 2, 6 H 4 Ge. O 4 8.6 H 3 PO 4 2. 1 H 3 As. O 4 2, 3

Consider the following titration cases.

Titration of a strong acid with a strong base

HCl + NaOH ® NaCl + H 2 O

H + + OH - ® H 2 O

At the equivalence point, a salt of a strong acid and a strong base is formed, which does not undergo hydrolysis. The reaction of the medium will be neutral (рН=7). In this case, litmus can serve as an indicator.

Titration of a weak acid with a strong base

CH 3 COOH + NaOH ® CH 3 COONa + H 2 O

CH 3 COOH + OH - ® CH 3 COO - + H 2 O

The resulting salt of a weak acid and a strong base in solution undergoes hydrolysis:

CH 3 COO - + HOH ® CH 3 COOH + OH -

The equivalence point in this case will be in an alkaline environment, so you should use an indicator that changes color at pH> 7, for example, phenolphthalein.

Titration of a weak base with a strong acid

NH 4 OH + HCl ® NH 4 Cl + H 2 O

NH 4 OH + H + ® NH 4 + + H 2 O

The resulting salt in solution undergoes hydrolysis:

NH 4 + + HOH ® NH 4 OH + H +

The equivalence point will be in an acidic environment, so methyl orange can be used.

SAMPLE design of laboratory work in

Titrimetric analysis

Lab No. ... Date

“Name of the lab”

Primary standard - C E (NaOH) = …………mol/l

Determined substance (titrant) - С E (HCl) = ?, T(HCl) = ?

Indicator - methyl orange

Titration conditions - (pH of the medium, heating, etc.)

Reaction equation (in molecular and ion-molecular forms):

The results of the experiment are entered in the table:

Calculations:

PREPARATION AND STANDARDIZATION OF SOLUTIONS OF TITRANTS FOR ACID-BASE TITRATION

1. Preparation and standardization of 0.1 M HCl

Using a hydrometer, determine the density of the concentrated hydrochloric acid solution given to you (let's say that r \u003d 1.179 g / ml).

According to the density table of solutions (Appendix Table 3), find the mass fraction of acid in this solution (w = 36%). Calculate what volume of 36% HCl solution you need to take to prepare 250 ml of a 0.1 mol/l solution.

The molar mass of an HCl equivalent is 36.46 g/mol, so 250 ml of a 0.1 mol/l solution should contain 0.912 g of anhydrous HCl:

m(HCl) \u003d M E C E V \u003d 36.46 0.1 0.25 \u003d 0.912 g

The mass of a 36% HCl solution containing this amount of acid is:

The volume of the initial acid solution can be found by the formula:

Use a small graduated cylinder to measure the calculated volume (≈ 2.0 ml) of 36% hydrochloric acid solution and pour into the large cylinder. Bring the volume of the solution to 250 ml with distilled water, pour it into a 250 ml bottle and mix.

1.2 Preparation of the primary standard solution

Sodium tetraborate and HCl react with each other as follows

Na 2 B 4 O 7 + 2HCl + 5H 2 O ® 2NaCl + 4H 3 BO 3

The equivalence factor Na 2 B 4 O 7 in this reaction is 1/2, M (1/2 Na 2 B 4 0 7 × 10H 2 O) = 190.686 g / mol.

To prepare 100 ml of 0.1M 1/2 Na 2 B 4 O 7, you need to take m = C × V × M = 0.1 × 0.1 × 190.686 = 1.9 g Na 2 B 4 0 7 × 10H 2 O .

First, about 1.9 g of Na 2 B 4 0 7 × 10H 2 O are weighed using a hand balance. The mass of the sample taken is then specified using an analytical balance and dissolved in about 50 ml of hot water in a volumetric flask with a capacity of 100 ml. After cooling, the solution is adjusted with water at a temperature environment to the mark and mix. Calculate the exact concentration of sodium tetraborate in solution

,

where m is the weight of the sample of sodium tetraborate decahydrate, measured using an analytical balance.

1.3. Titration of a primary standard solution with standardized HCl solution

Pipette 10.00 ml of sodium tetraborate solution into 3 titration flasks, add 2 drops of 1% methyl orange solution to each flask and titrate with the prepared HCl solution. The titration is carried out until the pure yellow color of the solution acquires an orange tint. For comparison, you can take 10 ml of distilled water and add to it the same amount of indicator as to the analyzed solution. Titration is carried out until there is a difference in the colors of both solutions.

The molar concentration of HCl in solution is

Where 10.00 ml, - the average value of the volume (ml) of the HCl solution used for 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 production 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. Of the large number of color indicators in chemistry, the most famous and used 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 donor 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 hydrochloric acid HCl is added. Next, a separate solution of dimethylaniline in HCl is prepared, which is poured into the first solution when cooled, obtaining a dye. It is further 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 medium, litmus turns blue, which makes it possible to use it as a universal indicator for the general determination of the medium 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.