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-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.

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 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, 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.

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 the preparation of which one drop of acid (if an alkali is being titrated) or alkali (if an acid is being titrated) and as many drops of an 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 of acids is carried out by calculation using the formulas given for these and the cases described above.

Preparation of solutions. A solution is a homogeneous mixture of two or more substances. The concentration of a solution is expressed in different ways:

in weight percent, i.e. by the number of grams of substance contained in 100 g of solution;

in volume percentage, i.e. by the number of volume units (ml) of the substance in 100 ml of solution;

molarity, i.e. the number of gram-moles of a substance contained in 1 liter of solution (molar solutions);

normality, i.e. the number of gram equivalents of the dissolved substance in 1 liter of solution.

Solutions of percentage concentration. Percentage solutions are prepared as approximate solutions, while a sample of the substance is weighed on a technochemical balance, and volumes are measured using measuring cylinders.

To prepare percentage solutions, several methods are used.

Example. It is necessary to prepare 1 kg of 15% sodium chloride solution. How much salt do you need to take for this? The calculation is carried out according to the proportion:

Therefore, for this you need to take 1000-150 = 850 g of water.

In cases where it is necessary to prepare 1 liter of 15% sodium chloride solution, the required amount of salt is calculated in a different way. Using the reference book, find the density of this solution and, multiplying it by the given volume, obtain the mass of the required amount of solution: 1000-1.184 = 1184 g.

Then it follows:

Therefore, the required amount of sodium chloride is different for preparing 1 kg and 1 liter of solution. In cases where solutions are prepared from reagents containing water of crystallization, it should be taken into account when calculating the required amount of reagent.

Example. It is necessary to prepare 1000 ml of a 5% solution of Na2CO3 with a density of 1.050 from a salt containing water of crystallization (Na2CO3-10H2O)

The molecular weight (weight) of Na2CO3 is 106 g, the molecular weight (weight) of Na2CO3-10H2O is 286 g, from here the required amount of Na2CO3-10H2O is calculated to prepare a 5% solution:

Solutions are prepared using the dilution method as follows.

Example. It is necessary to prepare 1 liter of 10% HCl solution from an acid solution with a relative density of 1.185 (37.3%). The relative density of a 10% solution is 1.047 (according to the reference table), therefore, the mass (weight) of 1 liter of such a solution is 1000X1.047 = 1047 g. This amount of solution should contain pure hydrogen chloride

To determine how much 37.3% acid needs to be taken, we make up the proportion:

When preparing solutions by diluting or mixing two solutions, the diagonal scheme method or the “rule of the cross” is used to simplify calculations. At the intersection of two lines, the given concentration is written, and at both ends on the left - the concentration of the initial solutions; for the solvent it is equal to zero.

Chemistry is a fascinating science. Those who are interested not only in theory, but also try their skills in practice, know exactly what we are talking about. Every schoolchild is familiar with most of the elements from the periodic table. But has everyone been able to try mixing reagents and conducting chemical tests first-hand? Even today, not all modern schools have the necessary equipment and reagents, so chemistry remains a science open to independent study. Many seek to understand it more deeply by conducting research at home.

Not a single home-made worker can do without nitric acid - an extremely important thing in the household. It is difficult to obtain the substance: it can only be purchased in a specialized store, where purchases are made using documents confirming the peaceful use of the substance. Therefore, if you are a DIYer, you most likely will not be able to get this component. This is where the question arises of how to make nitric acid at home. The process does not seem to be complicated, however, the output should be a substance of a sufficient level of purity and the required concentration. There is no way to do this without the skills of an experimental chemist.

Where is the substance used?

It is reasonable to use nitric acid for safe purposes. The substance is used in the following areas of human activity:

• creation of coloring pigments;
• developing photographic films;
• preparation of medicines;
• recycling of plastic products;
• use in chemistry;
• fertilization of garden and vegetable crops;
• dynamite production.

Pure nitric acid in its unchanged form appears as a liquid substance, which upon contact with air begins to release white vapors. It freezes already at -42 o C, and boils at +80 o C. How to remove a substance such as nitric acid with your own hands at home?

Method 1

The fuming substance is obtained by exposing the concentrate to sodium (potassium) nitrate (sodium (potassium) nitrate). As a result of the reaction, the desired substance and sodium (potassium) hydrogen sulfate are obtained. The reaction scheme looks like this: NaNO 3 + H 2 SO 4 => HNO 3 + NaHSO 4. Remember that the concentration of the resulting substance depends on before entering into the reaction.

Method 2

Obtaining nitric acid at home with a lower concentration of the substance occurs in the same way, you only need to replace sodium nitrate with ammonium nitrate. The chemical equation looks like this: N.H. 4 NO 3 + H 2 SO 4 =>(N.H.4) 2 SO 4 + HNO 3 . Please note that ammonium nitrate is more accessible than potassium or sodium nitrate, which is why most researchers carry out the reaction based on it.

The higher the concentration of H 2 SO 4, the more concentrated the nitric acid will be. To obtain a balanced substance, it is necessary to increase the volume of electrolyte required for the reaction. To achieve the desired result, in practice they use the evaporation method, which consists of gradually reducing the volume of the electrolyte by about 4 times the original.

Features of the evaporation method

Sifted sand is poured into the bottom of the dish and a reservoir with electrolyte is placed. In this case, the boiling process is regulated by the valve of the gas stove, turning up or reducing the fire. The process takes a long time, so patience is important in this matter. Experts recommend using boilers - glass or ceramic tubes designed for chemical experiments, including evaporation. They neutralize the formation of bubbles and reduce the boiling force, preventing splashing of the substance. Under such conditions, it is permissible to obtain nitric acid at home with a concentration of about 93%.

Tools and reagents for practical preparation of the substance

To carry out the reaction you will need:

• concentrated H 2 SO 4 (>95%) - 50 ml;
• ammonium nitrate, potassium, sodium;
• 100 ml container;
• 1000 ml container;
• glass funnel;
• elastic bands;
• water bath;
• crushed ice (can be replaced with snow or cold water);
• thermometer.

Obtaining nitric acid at home, like carrying out any other chemical reaction, requires the following precautions:

• In the process of producing nitric acid at home, it is necessary to maintain the temperature within 60-70 o C. If these limits are exceeded, the acid will begin to disintegrate.
• During the reaction, vapors and gases may be released, so when working with acids, be sure to use a protective mask. Hands must be protected from sudden contact of the substance with the skin, so chemists work in rubber gloves. In large chemical plants, where people come into contact with substances hazardous to health, workers generally work in special protective suits.

Now you know how to get nitric acid in a simple reaction. Be careful when using such a substance and use it only for peaceful purposes.

To prepare the solution, it is necessary to mix the calculated amounts of acid of known concentration and distilled water.

Example.

It is necessary to prepare 1 liter of HCL solution with a concentration of 6% wt. from hydrochloric acid with a concentration of 36% wt.(this solution is used in KM carbonatometers produced by NPP Geosphere LLC) .
By table 2Determine the molar concentration of an acid with a weight fraction of 6% wt. (1.692 mol/l) and 36% wt. (11.643 mol/l).
Calculate the volume of concentrated acid containing the same amount of HCl (1.692 g-eq.) as in the prepared solution:

1.692 / 11.643 = 0.1453 l.

Therefore, adding 145 ml of acid (36% wt.) to 853 ml of distilled water will obtain a solution of the given weight concentration.

Experiment 5. Preparation of aqueous solutions of hydrochloric acid of a given molar concentration.

To prepare a solution with the required molar concentration (Mp), it is necessary to pour one volume of concentrated acid (V) into the volume (Vв) of distilled water, calculated according to the ratio

Vв = V(M/Mp – 1)

where M is the molar concentration of the starting acid.
If the acid concentration is not known, determine it by density usingtable 2.

Example.

The weight concentration of the acid used is 36.3% wt. It is necessary to prepare 1 liter of an aqueous solution of HCL with a molar concentration of 2.35 mol/l.
By table 1find by interpolating the values ​​of 12.011 mol/l and 11.643 mol/l the molar concentration of the acid used:

11.643 + (12.011 – 11.643)·(36.3 – 36.0) = 11.753 mol/l

Using the above formula, calculate the volume of water:

Vв = V (11.753 / 2.35 – 1) = 4 V

Taking Vв + V = 1 l, obtain the volume values: Vв = 0.2 l and V = 0.8 l.

Therefore, to prepare a solution with a molar concentration of 2.35 mol/L, you need to pour 200 ml of HCL (36.3% wt.) into 800 ml of distilled water.

Questions and tasks:

1. 
   What is the concentration of a solution?
2. 
   What is the normality of a solution?
3. 
   How many grams of sulfuric acid are contained in the solution if 20 ml are used for neutralization? sodium hydroxide solution whose titer is 0.004614?

LPZ No. 5: Determination of residual active chlorine.

Materials and equipment:

Progress:

Iodometric method

Reagents:

1. Potassium iodide is chemically pure, crystalline, and does not contain free iodine.

Examination. Take 0.5 g of potassium iodide, dissolve in 10 ml of distilled water, add 6 ml of buffer mixture and 1 ml of 0.5% starch solution. The reagent should not turn blue.

2. Buffer mixture: pH = 4.6. Mix 102 ml of a molar solution of acetic acid (60 g of 100% acid in 1 liter of water) and 98 ml of a molar solution of sodium acetate (136.1 g of crystalline salt in 1 liter of water) and bring to 1 liter with distilled water, previously boiled.

3. 0.01 N sodium hyposulfite solution.

4. 0.5% starch solution.

5. 0.01 N solution of potassium dichromate. Setting the titer of a 0.01 N hyposulfite solution is carried out as follows: pour 0.5 g of pure potassium iodide into a flask, dissolve it in 2 ml of water, add first 5 ml of hydrochloric acid (1:5), then 10 ml of 0.01 N dichromate solution potassium and 50 ml of distilled water. The released iodine is titrated with sodium hyposulfite in the presence of 1 ml of starch solution, added at the end of the titration. The correction factor to the sodium hyposulfite titer is calculated using the following formula: K = 10/a, where a is the number of milliliters of sodium hyposulfite used for titration.

Analysis progress:

a) add 0.5 g of potassium iodide into a conical flask;

b) add 2 ml of distilled water;

c) stir the contents of the flask until the potassium iodide dissolves;

d) add 10 ml of buffer solution if the alkalinity of the water being tested is not higher than 7 mg/eq. If the alkalinity of the test water is higher than 7 mg/eq, then the number of milliliters of the buffer solution should be 1.5 times greater than the alkalinity of the test water;

e) add 100 ml of test water;

f) titrate with hyposulfite until the solution turns pale yellow;

g) add 1 ml of starch;

h) titrate with hyposulfite until the blue color disappears.

X = 3.55  N  K

where H is the number of ml of hyposulfite spent on titration,

K - correction factor to the titer of sodium hyposulfite.

Questions and tasks:

1. 
   What is the iodometric method?
2. 
   What is pH?

LPZ No. 6: Determination of chloride ion

Goal of the work:

Materials and equipment: drinking water, litmus paper, ash-free filter, potassium chromate, silver nitrate, titrated sodium chloride solution,

Progress:

Depending on the results of the qualitative determination, 100 cm 3 of the test water or a smaller volume (10-50 cm 3) is selected and adjusted to 100 cm 3 with distilled water. Chlorides are determined at concentrations up to 100 mg/dm 3 without dilution. The pH of the titrated sample should be in the range of 6-10. If the water is cloudy, it is filtered through an ashless filter washed with hot water. If the water has a color value above 30°, the sample is decolorized by adding aluminum hydroxide. To do this, add 6 cm3 of aluminum hydroxide suspension to 200 cm 3 of sample, and the mixture is shaken until the liquid becomes discolored. The sample is then filtered through an ashless filter. The first portions of the filtrate are discarded. A measured volume of water is added to two conical flasks and 1 cm 3 of potassium chromate solution is added. One sample is titrated with a solution of silver nitrate until a faint orange tint appears, the second sample is used as a control sample. If the chloride content is significant, a precipitate of AgCl is formed, which interferes with the determination. In this case, add 2-3 drops of titrated NaCl solution to the titrated first sample until the orange tint disappears, then titrate the second sample, using the first as a control sample.

The following interfere with the determination: orthophosphates in concentrations exceeding 25 mg/dm 3 ; iron in a concentration of more than 10 mg/dm3. Bromides and iodides are determined in concentrations equivalent to Cl - . When normally present in tap water, they do not interfere with determination.

2.5. Processing the results.

where v is the amount of silver nitrate spent on titration, cm 3;

K is the correction factor to the titer of the silver nitrate solution;

g is the amount of chlorine ion corresponding to 1 cm 3 solution of silver nitrate, mg;

V is the sample volume taken for determination, cm3.

Questions and tasks:

1. 
   Methods for determining chloride ions?
2. 
   Conductometric method for determining chloride ions?
3. 
   Argentometry.

LPZ No. 7 “Determination of total water hardness”

Goal of the work:

Materials and equipment:

Experiment 1. Determination of the total hardness of tap water

Measure 50 ml of tap water with a measuring cylinder and pour it into a 250 ml flask, add 5 ml of ammonia buffer solution and an indicator - eriochrome black T - until a pink color appears (a few drops or a few crystals). Fill the burette with 0.04 N EDTA solution (synonyms: Trilon B, Complexon III) to the zero mark.

Titrate the prepared sample slowly with constant stirring with a solution of complexone III until the pink color changes to blue. Record the titration result. Repeat the titration one more time.

If the difference in titration results exceeds 0.1 ml, then titrate the water sample a third time. Determine the average volume of complexone III (V K, CP) consumed for titration of water, and from it calculate the total hardness of the water.

F TOTAL = , (20) where V 1 – volume of analyzed water, ml; V K,SR – average volume of complexone III solution, ml; N K – normal concentration of complexone III solution, mol/l; 1000 – conversion factor mol/l to mmol/l.

Write the results of the experiment in the table:

V K,SR	N K	V 1	F GEN

Example 1. Calculate the hardness of water, knowing that 500 liters contain 202.5 g of Ca(HCO 3) 2.

Solution. 1 liter of water contains 202.5:500 = 0.405 g Ca(HCO 3) 2. The equivalent mass of Ca(HCO 3) 2 is 162:2 = 81 g/mol. Therefore, 0.405 g is 0.405:81 = 0.005 equivalent masses or 5 mmol eq/L.

Example 2. How many grams of CaSO 4 are contained in one cubic meter of water if the hardness due to the presence of this salt is 4 mmol eq

CONTROL QUESTIONS

1. What cations are called hardness ions?

2. What technological indicator of water quality is called hardness?

3. Why can’t hard water be used for steam recovery in thermal and nuclear power plants?

4. Which softening method is called thermal? What chemical reactions occur when water is softened using this method?

5. How is water softened using the sedimentation method? What reagents are used? What reactions take place?

6. Is it possible to soften water using ion exchange?

LPZ No. 8 “Photocolorimetric determination of element content in solution”

Purpose of the work: to study the design and operating principle of the KFK-2 photocolorimeter

PHOTOELECTROCOLORIMETERS. A photoelectric colorimeter is an optical device in which monochromatization of the radiation flux is carried out using light filters. Photoelectric concentration colorimeter KFK – 2.

Purpose and technical data. Single-beam photocolorimeter KFK - 2

designed for measuring transmittance, optical density and concentration of colored solutions, scattering suspensions, emulsions and colloidal solutions in the spectral region of 315–980 nm. The entire spectral range is divided into spectral intervals, separated using light filters. Transmission measurement limits from 100 to 5% (optical density from 0 to 1.3). The basic absolute error of transmittance measurement is no more than 1%. Rice. General view of KFK-2. 1 - illuminator; 2 - handle for inserting color filters; 3 - cuvette compartment; 4 - handle for moving cuvettes; 5 - handle (introducing photodetectors into the light flux) “Sensitivity”; 6 - handle for setting the device to 100% transmission; 7 - microammeter. Light filters. In order to isolate rays of certain wavelengths from the entire visible region of the spectrum, selective light absorbers - light filters - are installed in photocolorimeters on the path of light fluxes in front of the absorbing solutions. Operating procedure

1. Turn on the colorimeter 15 minutes before starting measurements. During heating, the cuvette compartment must be open (in this case, the curtain in front of the photodetector blocks the light beam).

2. Enter a working filter.

3. Set the colorimeter sensitivity to minimum. To do this, set the “SENSITIVITY” knob to position “1”, the “SETTING 100 ROUGH” knob to the extreme left position.

4. Set the colorimeter needle to zero using the “ZERO” potentiometer.

5. Place the cuvette with the control solution into the light beam.

6. Close the cuvette compartment lid

7. Using the “SENSITIVITY” and “SETTING 100 ROUGH” and “FINE” knobs, set the microammeter needle to the “100” division of the transmittance scale.

8. By turning the handle of the cuvette chamber, place the cuvette with the test solution into the light stream.

9. Take readings on the colorimeter scale in the appropriate units (T% or D).

10. After finishing work, unplug the colorimeter, clean and wipe dry the cuvette chamber. Determination of the concentration of a substance in a solution using KFK-2. When determining the concentration of a substance in a solution using a calibration graph, the following sequence should be observed:

examine three samples of potassium permanganate solution of different concentrations and record the results in a journal.

Questions and tasks:

1. 
   Design and principle of operation of KFK - 2

5. Information support for training(list of recommended educational publications, Internet resources, additional literature)

Basic literature for students:

1. Course of basic notes according to the program OP.06 Fundamentals of Analytical Chemistry.-Manual / A.G. Bekmukhamedova - teacher of general professional disciplines ASHT - Branch of the Federal State Budgetary Educational Institution of Higher Professional Education OGAU; 2014

Additional literature for students:

1. Klyukvina E.Yu. Fundamentals of general and inorganic chemistry: textbook / E.Yu. Klyukvina, S.G. Bezryadin. - 2nd ed. - Orenburg. Publishing center OSAU, 2011 - 508 pages.

Basic literature for teachers:

1. 1.Klyukvina E.Yu. Fundamentals of general and inorganic chemistry: textbook / E.Yu. Klyukvina, S.G. Bezryadin. - 2nd ed. - Orenburg. Publishing center OSAU, 2011 - 508 pages.

2. Klyukvina E.Yu. Laboratory notebook on analytical chemistry. - Orenburg: OSAU Publishing Center, 2012 - 68 pages

Additional reading for teachers:

1. 1.Klyukvina E.Yu. Fundamentals of general and inorganic chemistry: textbook / E.Yu. Klyukvina, S.G. Bezryadin. - 2nd ed. - Orenburg. Publishing center OSAU, 2011 - 508 pages.

2. Klyukvina E.Yu. Laboratory notebook on analytical chemistry. - Orenburg: OSAU Publishing Center, 2012 - 68 pages