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Ministry of Education and Science of the Russian Federation

Federal State Budgetary Educational Institution of Higher Professional Education

"South Ural State University"

(national research university)

Department "Technology and catering"

Preparation of acid solutions

Completed by: Sharapova V.N.

Checked by: Sidorenkova L.A.

Chelyabinsk 2014

• 1. Preparation of acid solutions
• 2. Calculations in the preparation of solutions and features of the preparation of solutions of different concentrations
• 2.1 Calculations when preparing solutions of normal concentration
• 2.2 Calculations in the preparation of solutions, the concentration of which is expressed in grams per 1 liter
• 2.3 Calculations when preparing solutions of a certain percentage concentration

1. Preparation of acid solutions

In analyzes by the method of neutralization, 0.1 N. and 0.5 n. accurate solutions of sulfuric and hydrochloric acids, and in other methods of analysis, for example, redox, often use 2 N. approximate solutions of these acids.

For quick preparation of accurate solutions, it is convenient to use fixanals, which are weighed portions (0.1 g-eq or 0.01 g-eq) of chemically pure substances, weighed with an accuracy of four to five significant figures, located in sealed glass ampoules. When preparing 1 liter. solution from fixanal receive 0.1 n. or 0.01 n. solutions. Small amounts of solutions of hydrochloric and sulfuric acids 0.1 N. concentrations can be prepared from fixanals. Standard solutions prepared from fixanals are usually used to establish or verify the concentration of other solutions. Acid fixans can be stored for a long time.

To prepare an exact solution from fixanal, the ampoule is washed with warm water, washing off the inscription or label from it, and wipe it well. If the inscription is made with paint, then it is removed with a cloth moistened with alcohol. In a 1 L volumetric flask. insert a glass funnel, and into it - a glass striker, the sharp end of which should be directed upwards. After that, the ampoule with fixanal is lightly hit with a thin bottom on the tip of the striker or allowed to fall freely so that the bottom breaks when it hits the tip. Then, with a glass pin with a pointed end, the thin wall of the recess in the upper part of the ampoule is broken and the liquid contained in the ampoule is allowed to flow out. Then the ampoule in the funnel is thoroughly washed with distilled water from the washer, after which it is removed from the funnel, the funnel is washed and removed from the flask, and the solution in the flask is added to the mark with distilled water, closed with a stopper and mixed.

When preparing solutions from dry fixanals (for example, from oxalic acid fixanal), a dry funnel is taken so that the contents of the ampoule can be poured into the flask with gentle shaking. After the substance is transferred into the flask, the ampoule and funnel are washed, the substance is dissolved in the water in the flask, and the volume of the solution is adjusted to the mark with distilled water.

Large quantities of 0.1 N. and 0.5 n. solutions of hydrochloric and sulfuric acids, as well as approximate solutions of these acids (2 N, etc.) are prepared from concentrated chemically pure acids. First, a hydrometer or densimeter determines the density of concentrated acid.

According to the density in the reference tables, the concentration of the acid is found (the content of hydrogen chloride in hydrochloric acid or monohydrate in sulfuric acid), expressed in grams per 1 liter. The formulas calculate the volume of concentrated acid required to prepare a given volume of acid of the appropriate concentration. The calculation is carried out with an accuracy of two or three significant figures. The amount of water to prepare the solution is determined by the difference between the volumes of the solution and concentrated acid.

Table 1. Density and concentration of hydrochloric acid solutions (15°С)

Density g / cm 3		Density g / cm 3

Table.2 Density and concentration of sulfuric acid solutions (15°С)

Density g / cm 3

A solution of hydrochloric acid is prepared by pouring half of the required amount of distilled water into a vessel for preparing a solution, and then concentrated acid; after stirring, the solution is topped up to full volume with the remaining amount of water. Part of the second portion of water is rinsed with a beaker, which measured the acid.

A sulfuric acid solution is prepared by slowly adding concentrated acid with constant stirring (to prevent heating) to water poured into a heat-resistant glass vessel. At the same time, a small amount of water is left to rinse the beaker, which was used to measure the acid, pouring this residue into the solution after it has cooled.

Sometimes solutions of solid acids (oxalic, tartaric, etc.) are used for chemical analysis. These solutions are prepared by dissolving a sample of chemically pure acid in distilled water.

The weight of the sample of acid is calculated by the formula. The volume of water for dissolution is taken approximately equal to the volume of the solution (if the dissolution is not carried out in a volumetric flask). To dissolve these acids, use water that does not contain carbon dioxide.

In the density table we find the content of hydrogen chloride HCl in concentrated acid: G to = 315 g / l.

Calculate the volume of concentrated hydrochloric acid solution:

V to \u003d 36.5N * V / T to \u003d 36.5 * 0.1 * 10000 / 315 \u003d 315 ml.

The amount of water required to prepare the solution:

V H2O = 10000 - 115 = 9885 ml.

Weighing weight of oxalic acid H2C2O4*2H2O:

63.03N * V / 1000 \u003d 63.03 * 0.1 * 3000 / 1000 \u003d 12.6 g.

Establishment of the concentration of working solutions of acids can be carried out on sodium carbonate, borax, an exact solution of alkali (titrated or prepared from fixanal). When determining the concentration of solutions of hydrochloric or sulfuric acids by sodium carbonate or by borax, the method of titration of samples or (less often) the method of pipetting is used. In the weighed titration method, burettes with a capacity of 50 or 25 ml are used.

When determining the concentration of acids, the choice of indicator is of great importance. The titration is carried out in the presence of an indicator whose color transition occurs in the pH range corresponding to the equivalence point for the chemical reaction occurring during the titration. When a strong acid interacts with a strong base, methyl orange, methyl red, phenolphthalein and others can be used as indicators, in which the color transition occurs at pH = 4h10.

In the interaction of a strong acid with a weak base or with salts of weak acids and strong bases, indicators are used in which the color transition occurs in an acidic environment, for example, methyl orange. When weak acids interact with strong alkalis, indicators are used in which the color transition occurs in an alkaline medium, for example, phenolphthalein. The concentration of a solution cannot be determined by titration if a weak acid interacts with a weak base during the titration.

When establishing the concentration of hydrochloric or sulfuric acid by sodium carbonate three or four weighings of anhydrous chemically pure sodium carbonate are taken on an analytical balance in separate weighing bottles with an accuracy of 0.0002 g. To establish a concentration of 0.1 n. solution by titration from a burette with a capacity of 50 ml, the weight of the sample should be about 0.15 g. By drying in an oven at 150 ° C, the samples are adjusted to constant weight, and then transferred to conical flasks with a capacity of 200-250 ml and dissolved in 25 ml of distilled water . The weighing bottles with carbonate residues are weighed and the exact weight of each sample is determined by the difference in masses.

Titration of sodium carbonate solution with acid is carried out in the presence of 1-2 drops of a 0.1% solution of methyl orange (titration ends in an acidic medium) until the yellow color of the solution changes to orange-yellow. When titrating, it is useful to use a solution - a "witness", for the preparation of which one drop of acid from the buret and as many drops of the indicator as it is added to the titrated solution are added to distilled water, poured into the same flask as the flask in which the titration is performed.

The volume of distilled water for the preparation of the "witness" solution should be approximately equal to the volume of the solution in the flask at the end of the titration.

The normal concentration of the acid is calculated from the results of the titration:

N = 1000m n / Oe Na2CO3 V = 1000m n / 52.99V

where m n is the weight of the sample of soda, g;

V is the volume of acid solution (ml) used for titration.

From several experiments take the average convergent value of the concentration.

We expect to spend about 20 ml of acid for titration.

Soda weight:

52.99 * 0.1 * 20 / 1000 = 0.1 g

Example 4 A portion of sodium carbonate in 0.1482 g was titrated with 28.20 ml of hydrochloric acid solution. Determine the concentration of the acid.

Normal concentration of hydrochloric acid:

1000 * 0.1482 / 52.99 * 28.2 = 0.1012 n.

When determining the concentration of an acid solution by sodium carbonate by pipetting, a sample of chemically pure sodium carbonate, previously brought to constant weight by drying in an oven and weighed to the nearest 0.0002 g, is dissolved in distilled water in a calibrated volumetric flask with a capacity of 100 ml.

The sample size when setting the concentration of 0.1 N. acid solution should be about 0.5 g (in order to obtain approximately 0.1 N solution when dissolved). For titration, take with a pipette 10-25 ml of sodium carbonate solution (depending on the capacity of the burette) and 1-2 drops of a 0.1% solution of methyl orange.

The pipetting method is often used to establish the concentration of solutions using semi-microburettes with a capacity of 10 ml with a division value of 0.02 ml.

The normal concentration of the acid solution when it is established by pipetting on sodium carbonate is calculated by the formula:

N \u003d 1000m n V 1 / 52.99V to V 2,

where m n is the mass of a sample of sodium carbonate, g;

V 1 - volume of carbonate solution taken for titration, ml;

V to - the volume of the volumetric flask in which the carbonate sample was dissolved;

V 2 - the volume of the acid solution used for titration.

Example 5 Determine the concentration of the sulfuric acid solution if, to establish it, 0.5122 g of sodium carbonate was dissolved in a volumetric flask with a capacity of 100.00 ml and 14.70 ml of the acid solution was used to titrate 15.00 ml of the carbonate solution (using a burette with a capacity of 25 ml) .

Normal concentration of sulfuric acid solution:

1000 * 0.5122 * 15 / 52.99 * 100 * 14.7 = 0.09860 n.

When establishing the concentration of sulfuric or hydrochloric acids by sodium tetraborate (storm) usually use the method of titration of weighed portions. Borax crystal hydrate Na 2 B 4 O 7 * 10H 2 O must be chemically pure and before establishing the acid concentration on it, it is subjected to recrystallization. For recrystallization, 50 g of borax is dissolved in 275 ml of water at 50-60°C; the solution is filtered and cooled to 25-30°C. Vigorously stirring the solution, cause crystallization. The crystals are filtered off on a Buchner funnel, redissolved and recrystallized. After filtration, the crystals are dried between sheets of filter paper at an air temperature of 20°C and a relative humidity of 70%; drying is carried out in air or in a desiccator over a saturated solution of sodium chloride. The dried crystals should not stick to the glass rod.

For titration, 3-4 weighed portions of borax are taken into a weighing bottle in turn with an accuracy of 0.0002 g and transferred to conical titration flasks, dissolving each portion in 40–50 ml of warm water with vigorous shaking. After transferring each sample from the weighing bottle to the flask, the weighing bottle is weighed. The difference in mass during weighing determines the value of each sample. The value of a separate sample of borax to establish a concentration of 0.1 N. acid solution when using a 50 ml burette should be about 0.5 g.

Borax solutions are titrated with acid in the presence of 1-2 drops of a 0.1% solution of methyl red until the yellow color of the solution changes to orange-red or in the presence of a solution of a mixed indicator consisting of methyl red and methylene blue.

The normal concentration of an acid solution is calculated by the formula:

N = 1000m n / 190.69V,

where m n is the weight of the sample of borax, g;

V is the volume of acid solution used for titration, ml.

It is supposed to use 15 ml of acid solution for titration.

Bur weight:

190.69 * 0.1 * 15 / 1000 = 0.3 g.

Example 7 Find the concentration of the hydrochloric acid solution if 24.38 ml of hydrochloric acid is used to titrate a 0.4952 g sample of borax.

1000 * 0,4952 / 190,624,38 = 0,1068

Establishing the concentration of acid in a solution of caustic soda or caustic potash is carried out by titration with an acid solution of an alkali solution in the presence of 1-2 drops of a 0.1% solution of methyl orange. However, this method of determining the acid concentration is less accurate than the one above. It is commonly used in acid concentration control tests. As the initial solution, an alkali solution prepared from fixanal is often used.

The normal concentration of acid solution N 2 is calculated by the formula:

N 2 \u003d N 1 V 1 / V 2,

where N 1 - normal concentration of alkali solution;

V 1 - volume of alkali solution taken for titration;

V 2 - the volume of the acid solution used for titration (the average value of convergent titration results).

Example 8 Determine the concentration of the sulfuric acid solution, if titration is 25.00 ml of 0.1000 N. sodium hydroxide solution consumed 25.43 ml of sulfuric acid solution.

Acid solution concentration:

0.1 * 25 / 25.43 = 0.09828 n.

2. Calculations in the preparation of solutions and features of the preparation of solutions of different concentrations

solution acid concentration beaker

The accuracy of calculations in the preparation of solutions depends on how the solution is prepared: approximate or exact. When calculating approximate solutions, atomic and molecular masses are rounded to three significant figures. So, for example, the atomic mass of chlorine is taken equal to 35.5 instead of 35.453, the atomic mass of hydrogen is 1.0 instead of 1.00797, etc. Rounding is usually done up.

When preparing standard solutions, calculations are carried out with an accuracy of five significant figures. The atomic masses of the elements are taken with the same accuracy. In calculations, five-digit or four-digit logarithms are used. Solutions, the concentration of which will then be determined by titration, are prepared, as well as approximate ones.

Solutions can be prepared by dissolving solids, liquids, or diluting more concentrated solutions.

2.1 Calculations when preparing solutions of normal concentration

A weighed portion of a substance (g) for preparing a solution of a certain normality is calculated by the formula:

m n \u003d ENV / 1000,

where E is the chemical equivalent of the dissolved substance;

N - the required normality of the solution, g-eq / l;

V is the volume of the solution, ml.

A sample of the substance is usually dissolved in a volumetric flask. Dilute approximate solutions can be prepared by dissolving a sample of the substance in a volume of solvent equal to the volume of the solution. This volume can be measured with a measuring cylinder or beaker.

If the solution is prepared from a sample of the crystalline hydrate of a substance, then the value of the chemical equivalent of the crystalline hydrate is substituted into the calculation equation to determine the sample.

When preparing a solution with a certain normal concentration by diluting a more concentrated solution, the volume of a concentrated solution (ml) is calculated by the formula:

V to \u003d ENV / T to,

where T to - the concentration of a concentrated solution, g / l, or:

where N to - the normality of a concentrated solution, or:

V to \u003d ENV / 10 p to d to,

where p to - percentage concentration of a concentrated solution;

d to - the density of the concentrated solution, g / cm 3.

Concentrated solutions are diluted in volumetric flasks. When preparing precise solutions (for example, standard solutions from a more concentrated standard solution), concentrated solutions are measured with pipettes or poured from burettes. When preparing approximate solutions, dilution can be done by mixing a concentrated solution with a volume of water equal to the difference between the volumes of the diluted and concentrated solutions:

2.2 Calculations in the preparation of solutions, the concentration of which is expressed in grams per 1 liter

The weight of the substance (g) for such solutions is calculated by the formula:

where T is the concentration of the solution, g/l;

V is the volume of the solution, ml.

The dissolution of a substance is usually carried out in a volumetric flask, bringing the volume of the solution after dissolution to the mark. Approximate solutions can be prepared by dissolving a sample in a volume of water equal to the volume of the solution.

If the solution is prepared from a sample of crystalline hydrate, and the concentration of the solution is expressed based on an anhydrous substance, the sample of crystalline hydrate is calculated by the formula:

m n \u003d TVM k / 1000M,

where M k is the molecular weight of the crystalline hydrate;

When preparing solutions by diluting more concentrated ones, the volume of a concentrated solution is determined by the formula:

where T k is the concentration of the concentrated solution, g/l, or:

V k \u003d 100VT / 1000p k d k,

where p k is the percentage concentration of the concentrated solution;

d k - density of the concentrated solution, g/cm 3 ;

V k \u003d VT / EN k,

where N k is the normal concentration of a concentrated solution; E is the chemical equivalent of a substance.

Solutions are prepared in the same way as in the preparation of solutions of a certain normal concentration by diluting more concentrated solutions.

For approximate calculations related to the preparation of solutions by diluting more concentrated ones, you can use the dilution rule ("rule of the cross"), which states that the volumes of mixed solutions are inversely proportional to the difference in concentrations of the mixed and mixed solutions. This is expressed in diagrams:

where N 1 , T 1 , N 3 , T 3 - concentrations of mixed solutions;

N 2 , T 2 - the concentration of the solution obtained by mixing;

V 1 , V 3 - volumes of mixed solutions.

If the solution is prepared by diluting a concentrated solution with water, then N 3 \u003d 0 or T 3 \u003d 0. For example, to prepare a solution of concentration T 2 \u003d 50 g / l from solutions of concentration T 1 \u003d 100 g / l and T 3 \u003d 20 g / l it is necessary to mix the volume V 1 \u003d 50 - 20 \u003d 30 ml of a solution with a concentration of 100 g / l and V 3 \u003d 100 - 50 \u003d 50 ml of a solution with a concentration of 20 g / l:

2.3 Calculations when preparing solutions of a certain percentage concentration

Sample weight (g) is calculated by the formula:

where p is the percentage concentration of the solution;

Q is the mass of the solution, g.

If the volume of the solution V is given, the mass of the solution is determined by:

where d is the density of the solution, g / cm 3 (can be found in the reference tables).

The weight of the sample for a given volume of solution is calculated:

The mass of water to dissolve the sample is determined by:

Since the mass of water is numerically approximately equal to its volume, water is usually measured with a graduated cylinder.

If the solution is prepared by dissolving the crystalline hydrate of the substance, and the concentration of the solution is expressed as a percentage of the anhydrous substance, then the mass of the crystalline hydrate is calculated by the formula:

m n \u003d pQM k / 100M,

where M k is the molecular weight of the crystalline hydrate;

M is the molecular weight of the anhydrous substance.

It is convenient to prepare solutions by diluting more concentrated ones by measuring certain volumes of solutions and water, while the volume of a concentrated solution is calculated by the formula:

V k \u003d pdV / p k d k,

where d k is the density of the concentrated solution.

Solutions of a certain percentage concentration are prepared as approximate, and therefore weighed samples of substances with an accuracy of two or three significant figures are weighed on technical scales, and beakers or measuring cylinders are used to measure volumes.

If a solution is obtained by mixing two other solutions, one of which has a higher concentration and the other a lower one, then the mass of the initial solutions can be determined using the dilution rule ("rule of the cross"), which for solutions of a certain percentage concentration says: the masses of the mixed solutions are inversely proportional differences in percentage concentrations of mixed and obtained solutions. This rule is expressed by the scheme:

For example, to obtain a solution at a concentration of p 2 \u003d 10% from solutions of concentration p 1 \u003d 20% and p 3 \u003d 5%, you need to mix the amount of initial solutions: m 1 \u003d 10-5 \u003d 5 g of a 20% solution and m 3 \u003d 20 -10=10g 5% solution. Knowing the density of the solutions, you can easily determine the volumes required for mixing.

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When preparing solutions of percentage concentration, the substance is weighed on a techno-chemical balance, and liquids are measured with a measuring cylinder. Therefore, a hitch! substances are calculated with an accuracy of 0.1 g, and the volume of 1 liquid with an accuracy of 1 ml.

Before proceeding with the preparation of the solution, | | it is necessary to make a calculation, i.e., calculate the amount of solute and solvent to prepare a certain amount of a solution of a given concentration.

CALCULATIONS IN THE PREPARATION OF SALT SOLUTIONS

Example 1. It is necessary to prepare 500 g of a 5% solution of potassium nitrate. 100 g of such a solution contains 5 g of KN0 3; 1 We make up the proportion:

100 g solution-5 g KN0 3

500 » 1 - X» KN0 3

5-500 "_ x \u003d -jQg- \u003d 25 g.

Water should be taken 500-25 = 475 ml.

Example 2. It is necessary to prepare 500 g of a 5% CaCl solution from CaCl 2 -6H 2 0 salt. First, we calculate for anhydrous salt.

100 g solution - 5 g CaCl 2 500 "" - X "CaCl 2 5-500 _ x = 100 = 25 g -

The molar mass of CaCl 2 \u003d 111, the molar mass of CaCl 2 - 6H 2 0 \u003d 219 *. Therefore, 219 g of CaCl 2 -6H 2 0 contain 111 g of CaCl 2 . We make a proportion:

219 g CaC1 2 -6H 2 0-111 g CaC1 2

X "CaCl 2 -6H 2 0-26" CaCI,

219-25 x \u003d -jjj- \u003d 49.3 g.

The amount of water is 500-49.3=450.7 g, or 450.7 ml. Since water is measured with a graduated cylinder, tenths of a milliliter are not taken into account. Therefore, you need to measure 451 ml of water.

CALCULATIONS IN THE PREPARATION OF ACID SOLUTIONS

When preparing acid solutions, it must be taken into account that concentrated acid solutions are not 100% and contain water. In addition, the required amount of acid is not weighed, but measured with a graduated cylinder.

Example 1. It is necessary to prepare 500 g of a 10% hydrochloric acid solution, based on the available 58% acid, the density of which is d=l,19.

1. Find the amount of pure hydrogen chloride that should be in the prepared acid solution:

100 g solution -10 g HC1 500 » » - X » HC1 500-10 * = 100 = 50 g -

* To calculate the solutions of the percentage concentration of the mole, the mass is rounded to whole numbers.

2. Find the number of grams of concentrated)
acid, which will contain 50 g of HC1:

100 g acid-38 g HC1 X » » -50 » HC1 100 50

X gg—"= 131.6 G.

3. Find the volume that this amount occupies 1
acids:

V--— 131 ‘ 6 110 6 sch

4. The amount of solvent (water) is 500-;
-131.6 = 368.4 g or 368.4 ml. Since the necessary co-
the amount of water and acid is measured with a measuring cylinder
rum, then tenths of a milliliter are not taken into account
ut. Therefore, to prepare 500 g of a 10% solution
hydrochloric acid, you need to take 111 ml of hydrochloric acid I
acids and 368 ml of water.

Example 2 Usually, in calculations for the preparation of acids, standard tables are used that indicate the percentage of an acid solution, the density of a given solution at a certain temperature, and the number of grams of this acid contained in 1 liter of a solution of a given concentration (see Annex V). In this case, the calculation is simplified. The amount of prepared acid solution can be calculated for a certain volume.

For example, you need to prepare 500 ml of a 10% hydrochloric acid solution, based on a concentrated 38% j solution. According to the tables, we find that a 10% hydrochloric acid solution contains 104.7 g of HC1 in 1 liter of solution. We need to prepare 500 ml of I, therefore, the solution should be 104.7: 2 \u003d 52.35 g of HO.

Calculate how much you need to take concentrated I acids. According to the table, 1 liter of concentrated HC1 contains 451.6 g of HC1. We make up the proportion: 1000 ml-451.6 g of HC1 X » -52.35 » HC1

1000-52.35 x \u003d 451.6 \u003d "5 ml.

The amount of water is 500-115 = 385 ml.

Therefore, to prepare 500 ml of a 10% hydrochloric acid solution, you need to take 115 ml of a concentrated HC1 solution and 385 ml of water.

Preparation of a sulfuric acid solution with a mass fraction of 5%. 28.3 cm 3 of concentrated sulfuric acid are mixed with 948 cm 3 of distilled water.

Preparation of a solution of mass concentration of manganese 0.1 mg/cm 3 . Potassium permanganate weighing 0.288 g is dissolved in a small amount of sulfuric acid solution with a mass fraction of 5% in a volumetric flask with a capacity of 1000 cm 3. The volume of the solution in the flask was adjusted to the mark with the same solution of sulfuric acid. The resulting solution is decolorized by adding a few drops of hydrogen peroxide or oxalic acid and stirred. The solution is stored for no more than 3 months at room temperature.

Preparation of reference solution. In volumetric flasks with a capacity of 50 cm 3 place a solution of the mass concentration of manganese 0.1 mg/cm 3 in the volumes indicated in the comparison table of solutions.

Table 1

Comparison table for manganese solutions

Add 20 cm3 of distilled water to each flask. Solutions are prepared on the day of the test.

Preparation of a solution of silver nitrate with a mass fraction of 1%. Silver nitrate weighing 1.0 g is dissolved in 99 cm 3 of distilled water.

Testing: Focusing on the premix formulation, take the volume of the test solution containing from 50 to 700 μg of manganese, place in glass beakers with a capacity of 100 cm 3 and evaporate to dryness on a sand bath or electric stove with asbestos mesh. The dry residue is moistened with drops of concentrated nitric and then sulfuric acids, the excess of which is evaporated. The treatment is repeated twice. The residue is then dissolved in 20 cm 3 of hot distilled water and transferred to a 50 cm 3 volumetric flask. The glass is washed several times with small portions of hot distilled water, which are also poured into a volumetric flask. 1 cm 3 of phosphoric acid, 2 cm 3 of a solution of silver nitrate with a mass fraction of 1% and 2.0 g of ammonium persulphate are added to the flasks with reference solutions and the test solution. The contents of the flasks are heated to a boil, and when the first bubble appears, more ammonium persulfate is added at the tip of the scalpel. After boiling, the solutions are cooled to room temperature, brought to the mark with a solution of sulfuric acid, mass fractions of 5%, and stirred. The optical density of solutions is measured on a photoelectrocolorimeter relative to the first reference solution that does not contain manganese, in cuvettes with a translucent layer thickness of 10 mm at a wavelength of (540 ± 25) nm, using an appropriate light filter, or on a spectrophotometer at a wavelength of 535 nm. At the same time, a control experiment is carried out, excluding taking a sample of the premix.

For the preparation of a 0.01-normal solution of sulfuric acid, it is necessary to have data on its concentration.

The concentration of sulfuric acid can be determined by the specific gravity, which in turn is determined by the indicator of a hydrometer lowered into a cylinder filled with this acid.

Knowing the specific gravity of sulfuric acid, it is possible to establish with the help of an auxiliary table and its concentration (see appendices). In other words, it is possible to determine how much chemically pure acid is contained in a particular volume of the mixture, as well as what percentage this amount corresponds to (industry produces sulfuric acid with an admixture of a small amount of water and some other substances).

The molecular weight of sulfuric acid is 98.06, and the equivalent is 49.03 g. Therefore, 1 liter of a 0.01 normal sulfuric acid solution should contain 0.4903 g of pure acid.

Having found out the required amount of pure sulfuric acid for the preparation of a centinormal solution, one can also determine the amount of strong sulfuric acid (with a predetermined concentration) to be taken to prepare the specified solution. So, for example, selling strong (concentrated) sulfuric acid, which usually has a specific gravity of 1.84 and contains 96% pure sulfuric acid, you need to take 0.5107 g (100 x 0.4902: 96), or 0.28 ml ( 0.5107:1.84).

The amount of concentrated sulfuric acid established by such a calculation (in this case 0.28 ml), which will be used to prepare a given solution, is filtered from a microburet with a ground tap into a volumetric flask, where distilled water is then poured to the level of the liter mark.

Then, a centinormal sulfuric acid solution is poured from the flask into a bottle, closed with a rubber stopper, through which an outlet glass tube connected to a microburette is passed into the solution, and a correction for the accuracy of the prepared solution is determined, since it is rarely possible to prepare an exact solution with a given normality. In most cases, these solutions with this method of preparation are slightly stronger or weaker than santinormal.

The correction for the accuracy of a centinormal sulfuric acid solution is often determined by the storm (Na2 B4 O7 10 H2 O).

This definition goes like this:

1. Weigh out on an analytical balance 953 mg of chemically pure borax (The equivalent weight of borax is 190.6 g. Hence, to prepare a liter of 0.01-normal solution, you need to take 1.906 g of chemically pure borax (190.6: 100), and to prepare 500 ml of a solution with the indicated normality, it is necessary to take 953 mg of borax).

2. The resulting sample, intended for the preparation of a 0.01-normal solution of borax, carefully, trying not to spill, transfer through a funnel into a 500 ml volumetric flask.

3. Pour into the flask with distilled water the grains of borax remaining on the funnel.

4. Dissolve the contents of the flask by shaking, and then bring the solution level to the 500 ml mark with distilled water.

5. Close the flask with a clean stopper and mix thoroughly the prepared borax solution.

6. Pour 20 ml of a 0.01 normal solution of borax into a small conical flask from a microburette or pipette, add 2 ... 3 drops of a two-color indicator there and titrate with a 0.01 normal solution of sulfuric acid.

7. Calculate for a 0.01-normal solution of sulfuric acid an amendment to the accuracy, which is expressed as a quotient obtained from dividing milliliters of a 0.01-normal solution of borax taken for titration by the number of milliliters of a 0.01-normal solution of sulfuric acid that went to neutralization. Let us explain what has been said with a specific example.

Suppose that 22 ml of sulfuric acid solution went to neutralize 20 ml of borax solution. This means that the prepared acid solution is weaker than 0.01 normal. If this solution corresponded to 0.01-normal, then an equal amount of acid solution would be used to neutralize each milliliter of borax solution.

In our example, as already mentioned, 22 ml of acid solution was used to neutralize 20 ml of borax solution, and hence the correction to the prepared acid solution:

The operation to establish the amendment is repeated 2-3 times. The results of parallel determinations must necessarily converge to within 0.001. The final value of the correction factor is taken as the arithmetic mean obtained from two or three determinations.

To recalculate the prepared sulfuric acid solution to an exact 0.01-normal solution, one or another of its quantities taken for analysis should be multiplied by the correction factor. Usually, the correction factor is written on a bottle with an acid solution and is periodically updated, since during long-term work with this solution or its long-term storage, it can change its strength.

approximate solutions. In most cases, the laboratory has to use hydrochloric, sulfuric and nitric acids. Acids are commercially available in the form of concentrated solutions, the percentage of which is determined by their density.

The acids used in the laboratory are technical and pure. Technical acids contain impurities, and therefore are not used in analytical work.

Concentrated hydrochloric acid smokes in air, so you need to work with it in a fume hood. The most concentrated hydrochloric acid has a density of 1.2 g/cm3 and contains 39.11% hydrogen chloride.

Dilution of the acid is carried out according to the calculation described above.

Example. It is necessary to prepare 1 liter of a 5% solution of hydrochloric acid, using its solution with a density of 1.19 g / cm3. According to the reference book, we learn that a 5% solution has a density of 1.024 g / cm3; therefore, 1 liter of it will weigh 1.024 * 1000 \u003d 1024 g. This amount should contain pure hydrogen chloride:

Acid with a density of 1.19 g/cm3 contains 37.23% HCl (we also find it in the reference book). To find out how much this acid should be taken, make up the proportion:

or 137.5 / 1.19 \u003d 115.5 acids with a density of 1.19 g / cm3. Having measured 116 ml of an acid solution, bring its volume to 1 liter.

Sulfuric acid is also diluted. When diluting it, remember that you need to add acid to water ~, and not vice versa. When diluted, strong heating occurs, and if water is added to the acid, then splashing is possible, which is dangerous, since sulfuric acid causes severe burns. If acid gets on clothes or shoes, quickly wash the spilled area with plenty of water, and then neutralize the acid with sodium carbonate or ammonia solution. In case of contact with the skin of the hands or face, immediately wash the area with plenty of water.

Special care must be taken when handling oleum, which is sulfuric acid monohydrate saturated with sulfuric anhydride SO3. According to the content of the latter, oleum can be of several concentrations.

It should be remembered that with a slight cooling, the oleum crystallizes and is in a liquid state only at room temperature. In air, it smokes with the release of SO3, which forms sulfuric acid vapors when interacting with air moisture.

Great difficulties are caused by the transfusion of oleum from a large container into a small one. This operation should be carried out either under draft or in air, but where the resulting sulfuric acid and SO3 cannot have any harmful effect on people and surrounding objects.

If the oleum has hardened, it should first be heated by placing the container with it in a warm room. When the oleum melts and turns into an oily liquid, it must be taken out into the air and poured into smaller dishes, using the method of squeezing with the help of air (dry) or an inert gas (nitrogen).

When mixed with water, nitric acid also heats up (although not as strong as in the case of sulfuric acid), and therefore precautions must be taken when working with it.

In laboratory practice, solid organic acids are used. Handling them is much easier and more convenient than liquid ones. In this case, care should only be taken to ensure that the acids are not contaminated by anything foreign. If necessary, solid organic acids are purified by recrystallization (see Ch. 15 "Crystallization"),

precise solutions. Accurate acid solutions they are prepared in the same way as the approximate ones, with the only difference that at first they strive to obtain a solution of a slightly higher concentration, so that after that it can be diluted accurately, according to calculation. For precise solutions, only chemically pure preparations are taken.

The required amount of concentrated acids is usually taken by volume, calculated from the density.

Example. It is necessary to prepare 0.1 and. H2SO4 solution. This means that 1 liter of solution should contain:

Acid with a density of 1.84 g / cmg contains 95.6% H2SO4 n for the preparation of 1 l of 0.1 n. solution, you need to take the following amount (x) of it (in g):

The corresponding volume of acid will be:

Having measured exactly 2.8 ml of acid from a burette, dilute it to 1 liter in a volumetric flask and then titrate with an alkali solution and establish the normality of the resulting solution. If the solution turns out to be more concentrated), the calculated amount of water is added to it from the burette. For example, during titration, it was found that 1 ml of 6.1 N. H2SO4 solution contains not 0.0049 g H2SO4, but 0.0051 g. To calculate the amount of water that is needed to prepare exactly 0.1 N. solution, make up the proportion:

The calculation shows that this volume is equal to 1041 ml. the solution must be added 1041 - 1000 = 41 ml of water. It should also take into account the amount of solution that is taken for titration. Let 20 ml be taken, which is 20/1000 = 0.02 of the available volume. Therefore, water should be added not 41 ml, but less: 41 - (41 * 0.02) \u003d \u003d 41 -0.8 \u003d 40.2 ml.

* To measure acid, use a carefully dried burette with a ground stopcock. .

The corrected solution should be checked again for the content of the substance taken for dissolution. Accurate solutions of hydrochloric acid are also prepared by the ion-exchange method, based on the exact calculated sample of sodium chloride. The sample calculated and weighed on an analytical balance is dissolved in distilled or demineralized water, the resulting solution is passed through a chromatographic column filled with a cation exchanger in the H-form. The solution flowing from the column will contain an equivalent amount of HCl.

As a rule, exact (or titrated) solutions should be stored in tightly closed flasks. It is imperative to insert a calcium chloride tube into the cork of the vessel, filled in the case of an alkali solution with soda lime or ascarite, and in the case of an acid, with calcium chloride or simply cotton wool.

To check the normality of acids, calcined sodium carbonate Na2COs is often used. However, it is hygroscopic and therefore does not fully meet the requirements of analysts. It is much more convenient to use for these purposes acidic potassium carbonate KHCO3, dried in a desiccator over CaCl2.

When titrating, it is useful to use a “witness”, for the preparation of which one drop of acid (if titrating alkali) or alkali (if titrating acid) and as many drops of indicator solution as added to the titrated solution are added to distilled or demineralized water.

The preparation of empirical, according to the substance being determined, and standard solutions, acids is carried out according to the calculation using the formulas given for these and the cases described above.