The gram equivalent of sulfuric acid is 49.04 (98.08:2), hydrochloric acid is 36.465. Therefore, to prepare normal solutions it is necessary to take sulfuric or hydrochloric acid in quantities corresponding to these values.

Sulfuric and hydrochloric acids are prepared from chemically pure concentrated solutions of these acids. The required amount of acids is calculated as follows. Suppose there is sulfuric acid with a relative density of 1.84 (95.6%), it is necessary to prepare 1 liter of 1 N. acid solution, for this you should take concentrated acid:

The required amount of hydrochloric acid is calculated in the same way. If the relative density of the concentrated acid is 1.185 (37.3%), then to prepare 1 liter of 1 N. you need to take the solution:

The required amount of acid is measured by volume, poured into water, cooled, then transferred to a 1-liter volumetric flask and the volume is adjusted to the mark.

The titer of acids is determined using chemically pure reagents: sodium carbonate, borax, or a titrated solution of sodium hydroxide.

Setting the sodium carbonate titer

Three portions of sodium carbonate of 0.15-0.20 g each (for a 0.1 N solution) are taken into separate bottles with an accuracy of 0.0001 g and dried at 150 ° C to constant mass (weight). After this, the samples are transferred into 250 ml conical flasks and dissolved in 25 ml of distilled water. The bottle is weighed again and the mass (weight) of a sample of the dried reagent is determined by the difference. An indicator - 1-2 drops of methyl orange - is added to the solution in the flask and titrated with the prepared acid solution until the color changes from yellow to orange-yellow. The correction factor is calculated using the formula (for 0.1 N solution)

where g is the weight of salt, g; V is the amount of acid consumed for titration, ml; 0.0053 - the amount of sodium carbonate corresponding to 1 ml of exactly 0.1 N. acid solution, g.

Setting the acid titer for borax

The borax is pre-dried between sheets of filter paper until individual crystals no longer stick to the glass rod. It is best to dry borax in a desiccator filled with a saturated solution of sodium chloride and sugar or a saturated solution of sodium bromide.

Take, with an accuracy of 0.0001 g, three portions of borax into flasks in the amount of 0.5 g (for a 0.1 N solution) and transfer them to conical flasks with a capacity of 250 ml, weigh the flasks and set them according to the difference exact mass(weight) of the hitch. Then add 30-60 ml of warm water to the samples, shaking vigorously. Then, adding 1-2 drops of methyl red solution, titrate the borax solution with the prepared acid solution until the color changes from yellow to red. The correction factor is calculated using the following formula:

where the meaning of the letters is the same as in the previous formula; 0.019072 - the amount of borax corresponding to 1 ml exactly 0.1 n. acid solution, g.

In neutralization assays, 0.1 N is used. and 0.5 n. accurate solutions of sulfuric and hydrochloric acids, and in other methods of analysis, for example, redox, 2 N is often used. approximate solutions of these acids.

To quickly prepare accurate solutions, it is convenient to use fixals, 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 l. solution from fixanal is obtained 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 check the concentration of other solutions. Fixanal acids can be stored for a long time.

To prepare an accurate solution from fixanal, the ampoule is washed 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 liter volumetric flask. insert a glass funnel, and into it a glass striker, the sharp end of which should be directed upward. After this, the ampoule with fixanal is lightly struck with its thin bottom against the tip of the striker or allowed to fall freely so that the bottom breaks when it hits the tip. Then, using a glass pin with a pointed end, they break the thin wall of the recess in the upper part of the ampoule and allow the liquid contained in the ampoule to flow out. Then the ampoule located in the funnel is thoroughly washed with distilled water from the wash, 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, capped and mixed.

When preparing solutions from dry fixinals (for example, from oxalic acid fixanal), take a dry funnel so that the contents of the ampoule can be poured into the flask with gentle shaking. After the substance is transferred to the flask, wash the ampoule and funnel, dissolve the substance in the water in the flask, and bring the volume of the solution to the mark with distilled water.

Large quantities 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, the density of the concentrated acid is determined using a hydrometer or densimeter.

Based on 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 are used to 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 for preparing the solution is determined by the difference in the volumes of the solution and concentrated acid.

A solution of hydrochloric acid is prepared by pouring half the required amount of distilled water into a vessel for preparing the solution, and then concentrated acid; After mixing, the solution is added to the full volume with the remaining amount of water. Use part of the second portion of water to rinse the beaker used to measure the acid.

A solution of sulfuric acid is prepared by slowly pouring concentrated acid with constant stirring (to prevent heating) to water poured into a heat-resistant glass vessel. In this case, a small amount of water is left to rinse the beaker with which the acid was measured, 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 mass of a sample of acid is calculated using 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, water that does not contain carbon dioxide is used.

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

We calculate the volume of a concentrated hydrochloric acid solution:

V k = 36.5N V / T k = 36.5 0.1 10000 / 315 = 315 ml.

Amount of water required to prepare the solution:

V H 2 O = 10000 - 115 = 9885 ml.

Weight of a sample of oxalic acid H2C2O4 2H2O:

63.03N V / 1000 = 63.03 0.1 3000 / 1000 = 12.6 g.

Establishing the concentration of working acid solutions can be carried out with sodium carbonate, borax, precise alkali solution (titrated or prepared from fixanal). When establishing the concentration of solutions of hydrochloric or sulfuric acids using sodium carbonate or borax, they use the titration method of weighed portions or (less often) the pipetting method. When using the titration method, burettes with a capacity of 50 or 25 ml are used.

When establishing the concentration of acids great importance has a choice of indicator. Titration is performed in the presence of an indicator whose color transition occurs in the pH range corresponding to the equivalence point for chemical reaction occurring during titration. When a strong acid interacts with a strong base, methyl orange, methyl red, phenolphthalein and others, in which the color transition occurs at pH = 4?10, can be used as indicators.

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

When establishing the concentration of hydrochloric or sulfuric acids based on sodium carbonate On an analytical balance in separate bottles, take three or four weighed portions of anhydrous chemically pure sodium carbonate 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 mass of the sample should be about 0.15 g. By drying in an oven at 150 ° C, the samples are brought 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 bottles with carbonate residues are weighed and the exact mass of each sample is determined from the difference in mass.

Titration of a solution of sodium carbonate with an 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 “witness” solution, for the preparation of which one drop of acid from a burette and as many drops of 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 preparing the “witness” solution should be approximately equal to the volume of the solution in the flask at the end of the titration.

The normal acid concentration is calculated from the titration results:

N = 1000m N/E Na 2 CO 3 V = 1000m N/52.99V

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

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

The average convergent concentration value is taken from several experiments.

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

Weight of soda sample:

52.99 0.1 20 / 1000 = 0.1 g.

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

Normal concentration of hydrochloric acid:

1000 0.1482 / 52.99 28.2 = 0.1012 n.

When determining the concentration of an acid solution with respect to sodium carbonate by pipetting, a sample of chemically pure sodium carbonate, previously brought to a constant mass by drying in an oven and weighed with an accuracy of 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 to 0.1 N. the acid solution should be about 0.5 g (to obtain approximately 0.1 N solution when dissolved). For titration, pipet 10-25 ml of sodium carbonate solution (depending on the capacity of the burette) and 1-2 drops of 0.1% methyl orange solution.

The pipetting method is often used to determine the concentration of solutions using 10 ml semi-microburettes with 0.02 ml divisions.

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

N = 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 k is the volume of the volumetric flask in which the carbonate sample was dissolved;

V 2 is the volume of acid solution consumed for titration.

Example 5. Determine the concentration of a sulfuric acid solution if, to establish it, 0.5122 g of sodium carbonate was dissolved in a 100.00 ml volumetric flask and 14.70 ml of an acid solution was used to titrate 15.00 ml of a 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 using sodium tetraborate (borax) Usually the titration method is used. Borax crystalline hydrate Na 2 B 4 O 7 10H 2 O must be chemically pure and before determining the acid concentration, it is subjected to recrystallization. For recrystallization, 50 g of borax are dissolved in 275 ml of water at 50-60°C; the solution is filtered and cooled to 25-30°C. Stirring the solution vigorously causes crystallization. The crystals are filtered on a Buchner funnel, dissolved again and recrystallized. After filtering, 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 sodium chloride solution. The dried crystals should not stick to the glass rod.

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

Titration of borax solutions with acid is carried out 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 using the formula:

N = 1000m n / 190.69V,

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

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

It is assumed that 15 ml of acid solution will be used for titration.

Weight of borax sample:

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 was used to titrate a 0.4952 g sample of borax.

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

Determination of acid concentration using sodium hydroxide solution or caustic potassium is carried out by titrating an alkali solution with an acid 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 above. It is usually used in control tests of acid concentrations. An alkali solution prepared from fixanal is often used as a starting solution.

The normal concentration of acid solution N2 is calculated using the formula:

N 2 = N 1 V 1 / V 2,

where N 1 is the normal concentration of the alkali solution;

V 1 - volume of alkali solution taken for titration;

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

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

Acid solution concentration.

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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 of 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 when preparing solutions and features of preparing solutions of different concentrations
2.1 Calculations for preparing solutions of normal concentration
2.2 Calculations when preparing solutions, the concentration of which is expressed in grams per 1 l
2.3 Calculations when preparing solutions of a certain percentage concentration

1. Preparation of acid solutions

To quickly prepare accurate solutions, it is convenient to use fixanals, which are weighed portions (0.1 g-equiv or 0.01 g-equiv) of chemically pure substances, weighed with an accuracy of four to five significant figures, located in sealed glass ampoules. When preparing 1 l. solution from fixanal is obtained 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 check the concentration of other solutions. Fixanal acids can be stored for a long time.

To prepare an accurate solution from fixanal, wash the ampoule 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 liter volumetric flask. insert a glass funnel, and into it a glass striker, the sharp end of which should be directed upward. After this, the ampoule with fixanal is lightly struck with its thin bottom against the tip of the striker or allowed to fall freely so that the bottom breaks when it hits the tip. Then, using a glass pin with a pointed end, they break the thin wall of the recess in the upper part of the ampoule and allow the liquid contained in the ampoule to flow out. Then the ampoule located in the funnel is thoroughly washed with distilled water from the wash, 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, capped and mixed.

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

Density g/cm 3		Density g/cm 3

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

Density g/cm 3

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.

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

We calculate the volume of a concentrated hydrochloric acid solution:

V c = 36.5N*V / T c = 36.5*0.1*10000 / 315 = 315 ml.

Amount of water required to prepare the solution:

V H2O = 10000 - 115 = 9885 ml.

Weight of a sample of oxalic acid H2C2O4*2H2O:

63.03N*V / 1000 = 63.03*0.1*3000 / 1000 = 12.6 g.

When determining the concentration of acids, the choice of indicator is of great importance. Titration is performed in the presence of an indicator in which the color transition occurs in the pH range corresponding to the equivalence point for the chemical reaction occurring during 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 = 4:10.

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

The normal acid concentration is calculated from the titration results:

N = 1000m N/E Na2CO3 V = 1000m N/52.99V

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

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

The average convergent concentration value is taken from several experiments.

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

Weight of soda sample:

52.99 * 0.1 * 20 / 1000 = 0.1 g.

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

Normal concentration of hydrochloric acid:

1000 * 0.1482 / 52.99 * 28.2 = 0.1012 n.

The pipetting method is often used to determine the concentration of solutions using 10 ml semi-microburettes with 0.02 ml divisions.

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

N = 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 k is the volume of the volumetric flask in which the carbonate sample was dissolved;

V 2 is the volume of acid solution consumed for titration.

Example 5. Determine the concentration of a sulfuric acid solution if, to establish it, 0.5122 g of sodium carbonate was dissolved in a 100.00 ml volumetric flask and 14.70 ml of an acid solution was used to titrate 15.00 ml of a 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 using sodium tetraborate (borax) Usually the titration method is used. Borax crystalline hydrate Na 2 B 4 O 7 *10H 2 O must be chemically pure and before determining the acid concentration, it is subjected to recrystallization. For recrystallization, 50 g of borax are dissolved in 275 ml of water at 50-60°C; the solution is filtered and cooled to 25-30°C. Stirring the solution vigorously causes crystallization. The crystals are filtered on a Buchner funnel, dissolved again and recrystallized. After filtering, 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 sodium chloride solution. The dried crystals should not stick to the glass rod.

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

N = 1000m n / 190.69V,

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

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

It is assumed that 15 ml of acid solution will be used for titration.

Weight of borax sample:

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 was used to titrate a 0.4952 g sample of borax.

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

The normal concentration of acid solution N2 is calculated using the formula:

N 2 = N 1 V 1 / V 2,

where N 1 is the normal concentration of the alkali solution;

V 1 - volume of alkali solution taken for titration;

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

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

Acid solution concentration:

0.1 * 25 / 25.43 = 0.09828 n.

2. Calculations when preparing solutions and features of preparing solutions of different concentrations

solution acid concentration beaker

The accuracy of calculations when preparing solutions depends on the type of solution being 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 to be 35.5 instead of 35.453, the atomic mass of hydrogen is 1.0 instead of 1.00797, etc. Rounding is usually done upward.

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

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

2.1 Calculations for preparing solutions of normal concentration

The weighed amount of a substance (g) for preparing a solution of a certain normality is calculated using the formula:

m n =ENV/1000,

where E is the chemical equivalent of the soluble substance;

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

V - volume of 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 using a graduated cylinder or beaker.

If a solution is prepared from a sample of 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 the concentrated solution (ml) is calculated using the formula:

V k =ENV/T k,

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

where Nk is the normality of the concentrated solution, or:

V to =ENV/10 p to d to,

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

dk is the density of the concentrated solution, g/cm3.

Concentrated solutions are diluted in volumetric flasks. When preparing exact 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 the concentrated solution with a volume of water equal to the difference between the volumes of the dilute and concentrated solutions:

2.2 Calculations when preparing solutions, the concentration of which is expressed in grams per 1 l

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

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

V - volume of solution, ml.

The dissolution of the 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 a solution is prepared from a sample of crystalline hydrate, and the concentration of the solution is expressed based on an anhydrous substance, the sample size of crystalline hydrate is calculated using the formula:

m n =TVM k /1000M,

where M k - molecular mass crystalline hydrate;

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

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

V k =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 =VT/EN k,

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

Solutions are prepared in the same way as when preparing 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 differences in the concentrations of the mixed and the solutions obtained by mixing. This is expressed by diagrams:

where N 1, T 1, N 3, T 3 are the concentrations of mixed solutions;

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

V 1, V 3 - volumes of mixed solutions.

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

2.3 Calculations when preparing solutions of a certain percentage concentration

The mass of the sample (g) is calculated using the formula:

where p is the percentage concentration of the solution;

Q is the mass of the solution, g.

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

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

The mass of the sample at a given volume of solution is calculated:

The mass of water for dissolving the sample is determined:

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

If a solution is prepared by dissolving the crystalline hydrate of a 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 using the formula:

m n =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 the concentrated solution is calculated using the formula:

V k =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 samples of substances with an accuracy of two or three significant figures are weighed on technical scales, and beakers or graduated 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 original solutions can be determined using the dilution rule (the “rule of the cross”), which for solutions of a certain percentage concentration states: the masses of the mixed solutions are inversely proportional differences in the percentage concentrations of the mixed and resulting solutions. This rule is expressed by the diagram:

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

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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.

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.

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

Example. You need to prepare 1 liter of a 5% solution of hydrochloric acid, using a solution with a density of 1.19 g/cm3. From the reference book we find out that a 5% solution has a density of 1.024 g/cm3; therefore, 1 liter of it will weigh 1.024 * 1000 = 1024 g. This amount should contain pure hydrogen chloride:

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

or 137.5/1.19 = 115.5 acid with a density of 1.19 g/cm3. Having measured out 116 ml of 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 you add water to the acid, it may splash, which is dangerous, since sulfuric acid causes severe burns. If acid gets on clothes or shoes, you should quickly wash the doused area with plenty of water, and then neutralize the acid with sodium carbonate or ammonia solution. In case of contact with the skin of your hands or face, immediately wash the area with plenty of water.

Particular care is required when handling oleum, which is a sulfuric acid monohydrate saturated with sulfuric anhydride SO3. According to the content of the latter, oleum comes in several concentrations.

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

It is very difficult to transfer oleum from large to small containers. 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 then poured into a smaller container, using the method of squeezing with air (dry) or an inert gas (nitrogen).

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

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

Precise solutions. Precise acid solutions They are prepared in the same way as approximate ones, with the only difference that at first they strive to obtain a solution of a slightly higher concentration, so that later it can be diluted precisely, according to calculations. For precise solutions, use only chemically pure preparations.

The required amount of concentrated acids is usually taken by volume calculated based on density.

Example. You need to prepare 0.1 and. H2SO4 solution. This means that 1 liter of solution should contain:

An acid with a density of 1.84 g/cmg contains 95.6% H2SO4 n to prepare 1 liter of 0.1 n. of the 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 the burette, dilute it to 1 liter in a volumetric flask and then titrate with an alkali solution to 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 a burette. For example, during titration it was found that 1 ml of 6.1 N. H2SO4 solution contains not 0.0049 g of H2SO4, but 0.0051 g. To calculate the amount of water needed to prepare exactly 0.1 N. solution, make up the proportion:

Calculation shows that this volume is 1041 ml; the solution needs to be added 1041 - 1000 = 41 ml of water. You should also take into account the amount of solution taken for titration. Let 20 ml be taken, which is 20/1000 = 0.02 of the available volume. Therefore, you need to add not 41 ml of water, but less: 41 - (41*0.02) = = 41 -0.8 = 40.2 ml.

* To measure the acid, use a thoroughly 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 using the ion exchange method, based on an accurately calculated sample of sodium chloride. The sample calculated and weighed on an analytical balance is dissolved in distilled or demineralized water, and 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, accurate (or titrated) solutions should be stored in tightly closed flasks. A calcium chloride tube must be inserted into the stopper of the vessel, filled with soda lime or ascarite in the case of an alkali solution, and with calcium chloride or simply cotton wool in the case of an acid.

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

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

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