Titration methods in analytical chemistry briefly. Neutralization method in the titrimetric method of analysis

Goal of the work : acquiring skills in using one of the methods of quantitative analysis - titrimetric, and learning basic techniques for statistical processing of measurement results.

Theoretical part

Titrimetric analysis is a method of quantitative chemical analysis based on measuring the volume of a reagent solution with a precisely known concentration consumed to react with the substance being determined.

Titrimetric determination of a substance is carried out by titration - adding one of the solutions to another in small portions and separate drops while constantly recording (monitoring) the result.

One of the two solutions contains a substance in an unknown concentration and represents the analyzed solution.

The second solution contains a reagent of precisely known concentration and is called a working solution, standard solution, or titrant.

Requirements for reactions used in titrimetric analysis:

1. The ability to fix the equivalence point, the most widely used is observation of its color, which can change under the following conditions:

One of the reactants is colored, and the colored reagent changes color during the reaction;

The substances used - indicators - change color depending on the properties of the solution (for example, depending on the reaction of the environment).

2. Quantitative course of the reaction, up to equilibrium, characterized by the corresponding value of the equilibrium constant

3. Sufficient rate of chemical reaction, because It is extremely difficult to fix the equivalence point in slow reactions.

4. Absence of side reactions in which accurate calculations are impossible.

Methods of titrimetric analysis can be classified according to the nature of the chemical reaction underlying the determination of substances: acid-base titration (neutralization), precipitation, complexation, oxidation-reduction.

Working with solutions.

Volumetric flasks designed to measure the exact volume of liquid. They are round, flat-bottomed vessels with a narrow, long neck, on which there is a mark to which the flask should be filled (Fig. 1).

Fig.1 Volumetric flasks

Technique for preparing solutions in volumetric flasks from fixanals.

To prepare a solution from fixanal, the ampoule is broken over a funnel inserted into a volumetric flask, the contents of the ampoule are washed off with distilled water; then dissolve it in a volumetric flask. The solution in the volumetric flask is brought to the mark. After bringing the liquid level to the mark, the solution in the flask is mixed well.

Burettes They are thin glass tubes graduated in milliliters (Fig. 2). A glass tap is soldered to the lower, slightly narrowed end of the burette or a rubber hose with a ball valve and a glass spout is attached. A burette is selected for work depending on the volume of solution used in the analysis.

Fig.2. Burettes

How to use a burette

1. The burette is washed with distilled water.

2. The burette prepared for work is fixed vertically in a stand; using a funnel, the solution is poured into the burette so that its level is above the zero mark.

3. Remove air bubbles from the lower extended end of the burette. To do this, bend it upward and release the liquid until all the air is removed. Then the capillary is lowered down.

4. The liquid level in the burette is set to zero division.

5. When performing a titration, press the rubber tube on the side of the ball and drain the liquid from the burette into the flask, rotating the latter. First, the titrant in the burette is poured out in a thin stream. When the color of the indicator at the point where the titrant drops fall begins to change, the solution is added carefully, drop by drop. The titration is stopped when a sharp change in the color of the indicator occurs due to the addition of one drop of titrant, and the volume of solution consumed is recorded.

6. At the end of the work, the titrant is drained from the burette, the burette is washed with distilled water.

Acid-base titration (neutralization) method

The acid-base titration method is based on the reaction between acids and bases, i.e. for neutralization reactions:

H + + OH¯ = H 2 O

When performing this task, the acid-base titration method is used, based on the use of a neutralization reaction:

2NaOH + H 2 SO 4 = Na 2 SO 4 + 2H 2 O

The method consists in gradually adding a solution of sulfuric acid of known concentration to a solution of the substance being determined - sodium hydroxide. The addition of the acid solution is continued until its amount becomes equivalent to the amount of sodium hydroxide reacting with it, i.e. until the alkali is neutralized. The moment of neutralization is determined by the change in color of the indicator added to the titrated solution. According to the law of equivalents according to the equation:

C n (k-you) · V (k-you) = C n (alkalis) · V (alkalis)

Cn(k-ty) and Cn(alkali) – molar concentrations of equivalents of reacting solutions, mol/l;

V (totals) and V (alkalis) – volumes of reacting solutions, l (ml).

C (NaOH) and - molar concentrations of equivalent NaOH and H 2 SO 4 in reacting solutions, mol/l;

V(NaOH) and ) - volumes of reacting solutions of alkali and acid, ml.

Examples of problem solving.

1. To neutralize 0.05 l of acid solution, 20 cm 3 of 0.5 N alkali solution was used. What is the normality of acid?

2. How much and what substance will remain in excess if 120 cm 3 of 0.3 N solution of potassium hydroxide is added to 60 cm 3 of 0.4 N sulfuric acid solution?

The solution to problems of determining the pH of a solution and concentrations of various types is presented in the methodological manual.

EXPERIMENTAL PART

Receive a flask with an alkali solution of unknown concentration from the laboratory assistant. Measure 10 ml samples of the analyzed solution into three conical titration flasks using a graduated cylinder. Add 2-3 drops of methyl orange indicator to each of them. The solution will turn yellow (methyl orange is yellow in an alkaline environment and orange-red in an acidic environment).

Prepare the titration installation for work (Fig. 3). Rinse the burette with distilled water, and then fill it with a solution of sulfuric acid of a precisely known concentration (the molar concentration of the equivalent of H 2 SO 4 is indicated on the bottle) above zero division. Bend the rubber tube with the glass tip up and, pulling the rubber away from the glass olive that covers the exit from the burette, slowly release the liquid so that after filling the tip there are no air bubbles left in it. Release the excess acid solution from the burette into a substitute glass, while the lower meniscus of the liquid in the burette should be set to zero.

Place one of the flasks of the alkali solution under the tip of the burette on a sheet of white paper and proceed directly to the titration: with one hand, slowly feed the acid from the burette, and with the other, continuously stir the solution in a circular motion of the flask in a horizontal plane. At the end of the titration, the acid solution should be fed dropwise from the burette until one drop turns the solution into a permanent orange color.

Determine the volume of acid used for titration, accurate to 0.01 ml. Count the divisions of the burette along the lower meniscus, while the eye should be at the level of the meniscus.

Repeat the titration 2 more times, each time starting from the zero division of the burette. Record the titration results in Table 1.

Calculate the concentration of the alkali solution using the formula:

Table 1

Results of titration of sodium hydroxide solution

Carry out statistical processing of the titration results according to the method described in the appendix. Summarize the results of statistical processing of experimental data in Table 2.

table 2

Results of statistical processing of experimental data from titration of sodium hydroxide solution. Confidence probability α = 0.95.

n			Sx

Write down the result of determining the molar concentration of NaOH equivalent in the analyzed solution as a confidence interval.

QUESTIONS FOR SELF-CONTROL

1. Potassium hydroxide solution has pH = 12. The concentration of the base in the solution at 100% dissociation is ... mol/l.

1) 0.005; 2) 0.01; 3) 0.001; 4) 1·10 -12; 5) 0.05.

2. To neutralize 0.05 l of acid solution, 20 cm3 of 0.5 N alkali solution was used. What is the normality of acid?

1) 0.2 n; 2) 0.5 n; 3) 1.0 n; 4) 0.02 n; 5) 1.25 n.

3. How much and what substance will remain in excess if 125 cm 3 of 0.2 N solution of potassium hydroxide is added to 75 cm 3 of 0.3 N solution of sulfuric acid?

1) 0.0025 g of alkali; 2) 0.0025 g acid; 3) 0.28 g of alkali; 4) 0.14 g of alkali; 5) 0.28 g of acid.

4. An analysis method based on determining the increase in boiling point is called...

1) spectrophotometric; 2) potentiometric; 3) ebullioscopic; 4) radiometric; 5) conductometric.

5. Determine the percentage concentration, molarity and normality of a solution of sulfuric acid obtained by dissolving 36 g of acid in 114 g of water, if the density of the solution is 1.031 g/cm3.

1) 31,6 ; 3,77; 7,54 ; 2) 31,6; 0,00377; 0,00377 ;

3) 24,0 ; 2,87; 2,87 ; 4) 24,0 ; 0,00287; 0,00287;

5) 24,0; 2,87; 5,74.

The titrimetric method of analysis (titration) allows for volumetric quantitative analysis and is widely used in chemistry. Its main advantage is the variety of methods and methods, thanks to which it can be used to solve a variety of analytical problems.

Principle of analysis

The titrimetric method of analysis is based on measuring the volume of a solution of known concentration (titrant) that reacted with the test substance.

For the analysis, you will need special equipment, namely a burette - a thin glass tube with graduations applied. The upper end of this tube is open, and at the lower end there is a shut-off valve. Using a funnel, the calibrated burette is filled with titrant to the zero mark. The analysis is carried out to the titration end point (ETP) by adding a small amount of solution from the burette to the test substance. The end point of the titration is identified by a change in the color of the indicator or some physicochemical property.

The final result is calculated based on the volume of titrant expended and is expressed in titer (T) - the mass of the substance per 1 ml of solution (g/ml).

Rationale for the process

The titrimetric method of quantitative analysis gives accurate results because the substances react with each other in equivalent quantities. This means that the product of their volume and quantity are identical to each other: C 1 V 1 = C 2 V 2. From this equation it is easy to find the unknown value of C 2 if the remaining parameters are set independently (C 1, V 2) and are established during the analysis (V 1).

Titration end point detection

Since timely recording of the end of the titration is the most important part of the analysis, it is necessary to choose the right methods. The most convenient is the use of color or fluorescent indicators, but instrumental methods can also be used - potentiometry, amperometry, photometry.

The final choice of a method for detecting CFTs depends on the required accuracy and selectivity of the determination, as well as its speed and the possibility of automation. This is especially true for cloudy and colored solutions, as well as aggressive environments.

Titration Reaction Requirements

In order for the titrimetric method of analysis to give the correct result, you need to correctly select the reaction that will underlie it. The requirements for it are as follows:

stoichiometry;
high flow rate;
high equilibrium constant;
the presence of a reliable method for recording the experimental end of the titration.

Suitable reactions may be of any type.

Types of analysis

The classification of titrimetric analysis methods is based on the type of reaction. Based on this feature, the following titration methods are distinguished:

acid-base;
redox;
complexometric;
precipitative.

Each type is based on its own type of reaction; specific titrants are selected, depending on which subgroups of methods are distinguished in the analysis.

Acid-base titration

The titrimetric method of analysis using the reaction of hydronium with hydroxide ion (H 3 O + + OH - = H 2 O) is called acid-base. If a known substance in solution forms a proton, which is typical for acids, the method belongs to the acidimetry subgroup. Here, stable hydrochloric acid HCl is usually used as a titrant.

If the titrant produces a hydroxide ion, the method is called alkalimetry. The substances used are alkalis, such as NaOH, or salts obtained by reacting a strong base with a weak acid, such as Na 2 CO 3.

In this case, colored indicators are used. They are weak organic compounds - acids and bases, which have different structures and colors of protonated and non-protonated forms. The most commonly used indicators in acid-base titration are single-color phenolphthalein (a clear solution turns crimson in an alkaline solution) and two-color methyl orange (a red substance turns yellow in an acidic solution).

Their widespread use is due to their high light absorption, due to which their color is clearly visible to the naked eye, and their contrast and narrow color transition region.

Redox titration

Redox titrimetric analysis is a method of quantitative analysis based on changing the ratio of concentrations of oxidized and reduced forms: aOx 1 + bRed 2 = aRed 1 + bOx 2.

The method is divided into the following subgroups:

permanganatometry (titrant - KMnO 4);
iodometry (I 2);
dichromatometry (K 2 Cr 2 O 7);
bromatometry (KBrO 3);
iodometry (KIO 3);
cerimetry (Ce(SO 4) 2);
vanadatometry (NH 4 VO 3);
titanometry (TiCl 3);
chromometry (CrCl 2);
ascorbinometry (C 6 H 8 OH).

In some cases, the role of an indicator can be played by a reagent that participates in the reaction and changes its color to acquire an oxidized or reduced form. But specific indicators are also used, for example:

when determining iodine, starch is used, which forms a dark blue compound with I 3 - ions;
When titrating ferric iron, thiocyanate ions are used, which form complexes with the metal, colored bright red.

In addition, there are special redox indicators - organic compounds that have different colors in their oxidized and reduced forms.

Complexometric titration

In short, the titrimetric method of analysis, called complexometric, is based on the interaction of two substances to form a complex: M + L = ML. If mercury salts are used, for example, Hg(NO 3) 2, the method is called mercurimetry, if ethylenediaminetetraacetic acid (EDTA) is called complexometry. In particular, using the latter method, a titrimetric method is used to analyze water, namely its hardness.

In complexometry, transparent metal indicators are used that acquire color when they form complexes with metal ions. For example, when titrating ferric salts with EDTA, clear sulfosalicylic acid is used as an indicator. It turns the solution red when it forms a complex with iron.

However, more often metal indicators have their own color, which changes depending on the concentration of the metal ion. Polybasic acids are used as such indicators, forming fairly stable complexes with metals, which are quickly destroyed when exposed to EDTA with a contrasting color change.

Precipitation titration

The titrimetric method of analysis, which is based on the reaction of interaction of two substances with the formation of a solid compound that precipitates (M + X = MX↓), is precipitation. It is of limited significance, since deposition processes are usually non-quantitative and non-stoichiometric. But sometimes it is still used and has two subgroups. If the method uses silver salts, for example, AgNO 3, it is called argentometry, if mercury salts, Hg 2 (NO 3) 2, then mercurometry.

The following methods are used to detect the titration end point:

Mohr's method, in which the indicator is the chromate ion, which forms a brick-red precipitate with silver;
Volhard's method, based on the titration of a solution of silver ions with potassium thiocyanate in the presence of ferric iron, which forms a red complex with the titrant in an acidic environment;
Faience method, which involves titration with adsorption indicators;
Gay-Lussac method, in which the CTT is determined by the clearing or turbidity of the solution.

The latter method has hardly been used recently.

Titration methods

Titration is classified not only by the underlying reaction, but also by the method of execution. Based on this feature, the following types are distinguished:

direct;
reverse;
titration of the substituent.

The first case is used only under ideal reaction conditions. The titrant is added directly to the substance being determined. Thus, magnesium, calcium, copper, iron and about 25 other metals are determined using EDTA. But in other cases, more complex methods are often used.

Back titration

It is not always possible to find the ideal reaction. Most often, it proceeds slowly, or it is difficult to find a method for fixing the end point of the titration, or volatile compounds are formed among the products, due to which the analyte is partially lost. These disadvantages can be overcome by using the back titration method. To do this, a large amount of titrant is added to the substance to be determined so that the reaction proceeds to completion, and then it is determined how much of the solution remains unreacted. To do this, the remaining titrant from the first reaction (T 1) is titrated with another solution (T 2), and its amount is determined by the difference in the products of volumes and concentrations in two reactions: C T1 V T 1 -C T 2 V T 2.

The use of the inverse titrimetric method of analysis underlies the determination of manganese dioxide. Its reaction with ferrous sulfate proceeds very slowly, so the salt is taken in excess and the reaction is accelerated by heating. The unreacted amount of iron ion is titrated with potassium dichromate.

Titration of the substituent

Substituent titration is used in the case of non-stoichiometric or slow reactions. Its essence is that a stoichiometric reaction with an auxiliary compound is selected for the substance being determined, after which the reaction product is subjected to titration.

This is exactly what is done when determining dichromate. Potassium iodide is added to it, resulting in the release of an amount of iodine equivalent to the substance being determined, which is then titrated with sodium thiosulfate.

Thus, titrimetric analysis makes it possible to determine the quantitative content of a wide range of substances. Knowing their properties and the characteristics of the reactions, you can choose the optimal method and titration method that will give results with a high degree of accuracy.

Introduction

The laboratory workshop is carried out after studying the theoretical course “Analytical chemistry and physical chemical analysis” and serves to consolidate and deepen the acquired knowledge.

The task of quantitative analysis is to determine the amount (content) of elements (ions), radicals, functional groups, compounds or phases in the analyzed object. This course covers the basic methods of titrimetric (volumetric) analysis, titration methods and their practical applications.

Before starting laboratory work, students undergo safety instructions. Before completing each work, the student must pass a colloquium on the sections specified by the teacher, as well as on the analysis methodology. To do this you need:

1) repeat the corresponding section of the course;

2) become familiar with the work methodology in detail;

3) draw up equations of chemical reactions that form the basis of the chemical analysis being carried out;

4) study the features of the analysis from a safety point of view.

Based on the results of their work, students draw up a report, which should indicate:

· job title;

· Objective;

· theoretical foundations of the method: essence of the method, basic equation, calculations and construction of titration curves, choice of indicator;

· reagents and equipment used during the work;

· analysis technique:

Preparation of primary standards;

Preparation and standardization of working solution;

Determination of the content of the test substance in solution;

· experimental data;

· statistical processing of analysis results;

· conclusions.

TITRIMETRIC ANALYSIS METHODS

Titrimetric method of analysis is based on measuring the volume of a reagent of precisely known concentration (titrant) spent on a chemical reaction with the substance being determined.

The determination procedure (titration) consists of adding a titrant dropwise from a burette to a precisely known volume of a solution of the analyte with an unknown concentration until the equivalence point is reached.

Where X– analyte; R– titrant, P– reaction product.

Equivalence point (i.e.)- this is the theoretical state of the solution that occurs at the moment of adding an equivalent amount of titrant R to the analyte X. In practice, the titrant is added to the analyte until it reaches the end point of titration (e.t.t.), which is understood in the visual indication of the equivalence point as the moment the color of the indicator added to the solution changes. In addition to visual indication, the equivalence point can be registered by instrumental means. In this case, the end point of titration (end point of titration) is understood as the moment of a sharp change in a physical quantity measured during the titration process (current strength, potential, electrical conductivity, etc.).

The titrimetric method of analysis uses the following types of chemical reactions: neutralization reactions, oxidation-reduction reactions, precipitation reactions and complexation reactions.

Depending on the type of chemical reaction used, the following are distinguished: titrimetric analysis methods:

– acid-base titration;

– precipitation titration;

– complexometric titration or complexometry;

– redox titration or redoximetry.

The reactions used in the titrimetric method of analysis require the following: requirements:

· the reaction must proceed in stoichiometric ratios, without side reactions;

· the reaction must proceed almost irreversibly (≥ 99.9%), the equilibrium constant of the reaction K p >10 6, the resulting precipitates must have solubility S < 10 -5 моль/дм 3 , а образующиеся комплексы – К уст > 10 -6 ;

· the reaction must proceed at a sufficiently high speed;

· the reaction must take place at room temperature;

· the equivalence point must be fixed clearly and reliably in some way.

Titration methods

In any titrimetric analysis method, there are several titration methods. Distinguish forward titration, back titration and displacement titration .

Direct titration– the titrant is added dropwise to the solution of the analyte until the equivalence point is reached.

Titration scheme: X + R = P.

Law of equivalents for direct titration:

C (1/ z) X V X = C (1/ z) R V R . (2)

The amount (mass) of the analyte contained in the test solution is calculated using the law of equivalents (for direct titration)

m X = C (1/z)R V R M (1/z) X٠10 -3 , (3)

Where C (1/ z) R– molar concentration of titrant equivalent, mol/dm 3 ;

V R– titrant volume, cm3;

M ( 1/ z) X– molar mass of the equivalent of the substance being determined;

C (1/ z) X– molar concentration of the equivalent of the analyte, mol/dm 3 ;

V X– volume of the substance being determined, cm3.

Back titration– two titrants are used. At first
The exact volume of the first titrant is added to the solution being analyzed ( R 1), taken in excess. The remainder of the unreacted titrant R1 is titrated with a second titrant ( R 2). Titrant quantity R 1, spent
for interaction with the analyte ( X) is determined by the difference between the added volume of titrant R 1 (V 1) and titrant volume R 2 (V 2) spent on titration of the remaining titrant R 1.

Titration scheme: X + R 1 fixed excess = P 1 (R 1 remainder).

R 1 remainder + R 2 = P2.

When using back titration, the law of equivalents is written as follows:

The mass of the analyte in the case of back titration is calculated using the formula

The reverse titration method is used in cases where it is impossible to select a suitable indicator for a direct reaction or it proceeds with kinetic difficulties (low rate of chemical reaction).

Titration by substitution (indirect titration)– used in cases where direct or reverse titration of the analyte is impossible or difficult, or when a suitable indicator is not available.

To the analyte X add some reagent A in excess, upon interaction with which an equivalent amount of the substance is released R. Then the reaction product R titrate with a suitable titrant R.

Titration scheme: X + A excess = P1.

P 1 + R = P2.

The law of equivalents for titration by substitution is written as follows:

Since the number of equivalents of the analyte is X and reaction product R are the same, the calculation of the mass of the analyte in the case of indirect titration is calculated using the formula

m X = C (1/z) R V R M (1/z) X٠10 -3 . (7)

Reagents

1. Succinic acid H 2 C 4 H 4 O 4 (reagent grade) – primary standard.

2. Sodium hydroxide NaOH solution with molar concentration
~2.5 mol/dm 3

3. H 2 O distilled.

Equipment students describe on their own.

Work progress:

1. Preparation of the primary standard of succinic acid HOOCCH 2 CH 2 COOH.

Succinic acid is prepared in a volume of 200.00 cm 3 with a molar concentration of the equivalent mol/dm 3 .

g/mol.

Reaction equation:

Taking a sample (weighing):

Hitch weight

Weighed quantitatively transferred to a volumetric flask ( cm 3), add 50 - 70 cm 3 of distilled water, stir until succinic acid is completely dissolved, adjust to the mark with distilled water
and mix thoroughly.

count on
according to the formula

Reagents

1. Sodium carbonate Na 2 CO 3 (reagent grade) – primary standard.

2. H 2 O distilled.

3. Hydrochloric acid HCl concentration 1:1 (r=1.095 g/cm3).

4. Acid-base indicator (selected according to the titration curve).

5. Mixed indicator - methyl orange and methylene blue.

Work progress:

1. Preparation of primary standard sodium carbonate (Na 2 CO 3).

A sodium carbonate solution is prepared with a volume of 200.00 cm 3 with a molar concentration of the equivalent mol/dm 3 .

Calculation of sample mass, g: (mass is taken accurate to the fourth decimal place).

Reaction equations:

1) Na 2 CO 3 + HCl = NaHCO 3 + NaCl

2) NaHCO 3 + HCl = NaCl + H 2 O + CO 2

_____________________________________

Na 2 CO 3 + 2HCl = 2NaCl + H 2 O + CO 2

H 2 CO 3 – weak acid (K a1= 10 -6.35 , K a2 = 10 -10,32).

Taking a sample (weighing):

Weight of watch glass (glass)

Weight of watch glass (glass) with weight

Hitch weight

Weighed quantitatively transferred to a volumetric flask ( cm 3), add 50 - 70 cm 3 of distilled water, mix until sodium carbonate is completely dissolved, adjust to the mark with distilled water
and mix thoroughly.

Actual concentration of the primary standard count on
according to the formula

2. Preparation and standardization of titrant (HCl solution)

A solution of hydrochloric acid is prepared with a volume of approximately 500 cm3
with a molar concentration equivalent of approximately 0.05÷0.06 mol/dm 3)

Titrant - a solution of hydrochloric acid with an approximate concentration of 0.05 mol/dm 3 is prepared from hydrochloric acid diluted 1:1 (r = 1.095 g/cm 3).

Standardization of the solution HCl is carried out according to the primary standard Na 2 CO 3 by direct titration, using the pipetting method.

The indicator is selected according to the titration curve of sodium carbonate with hydrochloric acid (Fig. 4).

Rice. 4. Titration curve of 100.00 cm 3 Na 2 CO 3 solution with WITH= 0.1000 mol/dm 3 HCl solution with C 1/ z= 0.1000 mol/dm 3

When titrating to the second equivalence point, use the indicator methyl orange, 0.1% aqueous solution (pT = 4.0). Change in color from yellow to orange (tea rose color). Transition interval
(pH = 3.1 – 4.4).

Scheme 3. Standardization of HCl solution

Place a 25.00 cm 3 aliquot of a standard Na 2 CO 3 solution (with a pipette) into a conical titration flask with a capacity of 250 cm 3, add 2–3 drops of methyl orange, dilute with water to 50–75 cm 3 and titrate with a solution of hydrochloric acid until the color changes. from yellow to “tea rose” color with one drop of titrant. Titration is carried out in the presence of a “witness” (a stock solution of Na 2 CO 3 with an indicator). The titration results are recorded in the table. 4. The concentration of hydrochloric acid is determined according to the law of equivalents: .

Table 4

Results of standardization of hydrochloric acid solution

Tasks

1. Formulate the concept of equivalent in acid-base reactions. Calculate the equivalents of soda and phosphoric acid in the following reactions:

Na 2 CO 3 + HCl = NaHCO 3 + NaCl

Na 2 CO 3 + 2HCl = 2NaCl + CO 2 + H 2 O

H 3 PO 4 + NaOH = NaH 2 PO 4 + H 2 O

H 3 PO 4 + 2NaOH = Na 2 HPO 4 + H 2 O

H 3 PO 4 + 3NaOH = Na 3 PO 4 + 3H 2 O

2. Write the reaction equations between hydrochloric acid, sulfuric acid, sodium hydroxide, aluminum hydroxide, sodium carbonate, potassium bicarbonate and calculate the equivalent mass of these substances.

3. Plot a titration curve for 100.00 cm 3 of hydrochloric acid with a molar concentration equivalent to 0.1 mol/dm 3 with sodium hydroxide with a molar concentration equivalent to 0.1 mol/dm 3. Select possible indicators

4. Plot a titration curve for 100.00 cm 3 acrylic acid (CH 2 =CHCOOH, pK a= 4.26) with molar concentration equivalent
0.1 mol/dm 3 sodium hydroxide with molar concentration equivalent
0.1 mol/dm3. How does the composition of a solution change during titration? Select possible indicators and calculate the indicator error of the titration.

5. Plot a titration curve for hydrazine (N 2 H 4 + H 2 O, pK b= 6,03)
with a molar concentration equivalent to 0.1 mol/dm 3 hydrochloric acid
with a molar concentration equivalent of 0.1 mol/dm 3 . What are the similarities
and the difference in pH calculations and titration curve compared to the titration curve of a weak acid with alkali? Select possible indicators
and calculate the indicator error of titration.

6. Calculate activity coefficients and active ion concentrations
in 0.001 M solution of aluminum sulfate, 0.05 M sodium carbonate, 0.1 M potassium chloride.

7. Calculate the pH of a 0.20 M solution of methylamine if its ionization in an aqueous solution is described by the equation

B + H 2 O = BH + + OH - , K b= 4.6 ×10 - 3, where B is the base.

8. Calculate the dissociation constant of hypochlorous acid HOCl if a 1.99 × 10 - 2 M solution has pH = 4.5.

9. Calculate the pH of a solution containing 6.1 g/mol glycolic acid (CH 2 (OH)COOH, K A= 1.5 × 10 - 4).

10. Calculate the pH of the solution obtained by mixing 40 ml of 0.015 M hydrochloric acid solution with:

a) 40 ml of water;

b) 20 ml of 0.02 M sodium hydroxide solution;

c) 20 ml of 0.02 M barium hydroxide solution;

d) 40 ml of 0.01 M solution of hypochlorous acid, K A=5.0 × 10 - 8.

11. Calculate the concentration of acetate ion in a solution of acetic acid
with a mass fraction of 0.1%.

12. Calculate the concentration of ammonium ion in an ammonia solution with a mass fraction of 0.1%.

13. Calculate the mass of a sample of sodium carbonate required to prepare 250.00 ml of a 0.5000 M solution.

14. Calculate the volume of a solution of hydrochloric acid with a molar concentration equivalent to 11 mol/l and the volume of water that must be taken to prepare 500 ml of a 0.5 M solution of hydrochloric acid.

15. 0.15 g of metallic magnesium was dissolved in 300 ml of a 0.3% solution of hydrochloric acid. Calculate the molar concentration of hydrogen, magnesium and chlorine ions in the resulting solution.

16. When 25.00 ml of sulfuric acid solution is mixed with a barium chloride solution, 0.2917 g of barium sulfate is obtained. Determine the titer of the sulfuric acid solution.

17. Calculate the mass of calcium carbonate that reacted
with 80.5 mmol hydrochloric acid.

18. How many grams of monosodium phosphate should be added?
to 25.0 ml of 0.15 M sodium hydroxide solution to obtain a solution with pH = 7? For phosphoric acid pK a1= 2.15; pK a2= 7.21; pK a3 = 12,36.

19. To titrate 1.0000 g of fuming sulfuric acid, thoroughly diluted with water, 43.70 ml of 0.4982 M sodium hydroxide solution is consumed. Fuming sulfuric acid is known to contain sulfuric anhydride dissolved in anhydrous sulfuric acid. Calculate the mass fraction of sulfuric anhydride in fuming sulfuric acid.

20. The absolute error in measuring volume using a burette is 0.05 ml. Calculate the relative error of measuring volumes in 1; 10 and 20 ml.

21. A solution is prepared in a volumetric flask with a capacity of 500.00 ml
from a sample of 2.5000 g of sodium carbonate. Calculate:

a) molar concentration of the solution;

b) molar concentration of the equivalent (½ Na 2 CO 3);

c) solution titer;

d) titer for hydrochloric acid.

22. What is the volume of 10% sodium carbonate solution with the density
1.105 g/cm 3 needs to be taken for preparation:

a) 1 liter of solution with a titer of TNa 2 CO 3 = 0.005000 g/cm 3 ;

b) 1 liter of solution with TNa 2 CO 3 /HCl = 0.003000 g/cm 3?

23. What volume of hydrochloric acid with a mass fraction of 38.32% and a density of 1.19 g/cm3 should be taken to prepare 1500 ml of a 0.2 M solution?

24. What volume of water must be added to 1.2 L of 0.25 M HCl to prepare a 0.2 M solution?

25. From 100 g of technical sodium hydroxide containing 3% sodium carbonate and 7% indifferent impurities, 1 liter of solution was prepared. Calculate the molar concentration and hydrochloric acid titer of the resulting alkaline solution, assuming that sodium carbonate is titrated to carbonic acid.

26. There is a sample that may contain NaOH, Na 2 CO 3, NaHCO 3 or a mixture of these compounds weighing 0.2800 g. The sample was dissolved in water.
To titrate the resulting solution in the presence of phenolphthalein, 5.15 ml is consumed, and in the presence of methyl orange - 21.45 ml of hydrochloric acid with a molar concentration equivalent of 0.1520 mol/l. Determine the composition of the sample and the mass fractions of components in the sample.

27. Plot a titration curve for a 100.00 cm 3 0.1000 M ammonia solution with a 0.1000 M hydrochloric acid solution, justify the choice of indicator.

28. Calculate the pH of the equivalence point, beginning and end of the titration of 100.00 cm 3 0.1000 M malonic acid solution (HOOCCH 2 COOH) with 0.1000 M sodium hydroxide solution (pK a 1=1.38; rK a 2=5,68).

29. The titration of 25.00 cm 3 of sodium carbonate solution with a molar concentration equivalent of 0.05123 mol/dm 3 required 32.10 cm 3 of hydrochloric acid. Calculate the molar concentration of hydrochloric acid equivalent.

30. How many ml of 0.1 M ammonium chloride solution must be added
to 50.00 ml of 0.1 M ammonia solution to form a buffer solution
with pH=9.3.

31. A mixture of sulfuric and phosphoric acids was transferred to a 250.00 cm 3 volumetric flask. For titration, two samples of 20.00 cm 3 were taken, one was titrated with a solution of sodium hydroxide with a molar concentration of the equivalent
0.09940 mol/dm 3 with methyl orange indicator, and the second with phenolphthalein. The sodium hydroxide consumption in the first case was 20.50 cm 3 , and in the second case 36.85 cm 3 . Determine the masses of sulfuric and phosphoric acids in the mixture.

In complexometry

Up to the equivalence point =( C M V M – C EDTA V EDTA)/( V M+ V EDTA). (21)

At the equivalence point = . (22)

After the equivalence point = . (23)

In Fig. Figure 9 shows the titration curves of calcium ion in buffer solutions with different pH values. It can be seen that titration of Ca 2+ is possible only at pH ³ 8.

Reagents

2. H 2 O distilled.

3. Standard solution of Mg(II) with molar concentration
0.0250 mol/dm3.

4. Ammonia buffer with pH = 9.5.

5. Solution of potassium hydroxide KOH with a mass fraction of 5%.

6. Eriochrome black T, indicator mixture.

7. Kalcon, indicator mixture.

Theoretical foundations of the method:

The method is based on the interaction of Ca 2+ and Mg 2+ ions with the disodium salt of ethylenediaminetetraacetic acid (Na 2 H 2 Y 2 or Na-EDTA) with the formation of stable complexes in the molar ratio M:L=1:1 in a certain pH range.

To fix the equivalence point when determining Ca 2+ and Mg 2+, calcon and eriochrome black T are used.

Determination of Ca 2+ is carried out at pH ≈ 12, while Mg 2+ is
in solution in the form of a precipitate of magnesium hydroxide and is not titrated with EDTA.

Mg 2+ + 2OH - = Mg(OH) 2 ↓

Ca 2+ + Y 4- « CaY 2-

At pH ≈ 10 (ammonia buffer solution), Mg 2+ and Ca 2+ are
in solution in the form of ions and upon addition of EDTA are titrated together.

Ca 2+ + HY 3- « CaY 2- + H +

Mg 2+ + HY 3- « MgY 2- +H +

To determine the volume of EDTA spent on the titration of Mg 2+,
from the total volume used for titrating the mixture at pH ≈ 10, subtract the volume used for titration of Ca 2+ at pH ≈ 12.

To create a pH ≈ 12, use a 5% KOH solution to create
pH ≈ 10 use an ammonia buffer solution (NH 3 × H 2 O + NH 4 Cl).

Work progress:

1. Standardization of titrant - EDTA solution (Na 2 H 2 Y)

An EDTA solution is prepared with an approximate concentration of 0.025 M
from ≈ 0.05 M solution, diluting it with distilled water 2 times. To standardize EDTA, use a standard solution of MgSO 4
with a concentration of 0.02500 mol/dm3.

Scheme 5. Standardization of titrant - EDTA solution

In a conical titration flask with a capacity of 250 cm 3, place 20.00 cm 3 of a standard MgSO 4 solution with a concentration of 0.02500 mol/dm 3, add ~ 70 cm 3 of distilled water, ~ 10 cm 3 of ammonia buffer solution with pH ~ 9.5 – 10 and add the indicator eriochrome black T about 0.05 g
(at the tip of the spatula). In this case, the solution turns wine red. The solution in the flask is slowly titrated with EDTA solution until the color changes from wine red to green. The titration results are recorded in the table. 6. The concentration of EDTA is determined according to the law of equivalents: .

Table 6

Results of standardization of EDTA solution

2. Determination of Ca 2+ content

Titration curves of Ca 2+ with EDTA solution at pH=10 and pH=12 are constructed independently.

The solution of the problem in a volumetric flask is brought to the mark with distilled water and mixed thoroughly.

Scheme 6. Determination of Ca 2+ content in solution

An aliquot of the test solution 25.00 cm 3 containing calcium and magnesium is placed in a conical titration flask with a capacity of 250 cm 3, ~ 60 cm 3 of water, ~ 10 cm 3 of a 5% KOH solution are added. After an amorphous precipitate of Mg(OH) 2 ↓ has formed, a calcone indicator of about 0.05 g is added to the solution (at the tip of a spatula) and slowly titrated with an EDTA solution until the color changes from pink to pale blue. Titration results ( V 1) are entered in Table 7.

Table 7

Experience no.		Volume of EDTA, cm 3	Ca 2+ content in solution, g
	25,00
	25,00
	25,00
	25,00
	25,00

3. Determination of Mg 2+ content

The titration curve of Mg 2+ with EDTA solution at pH=10 is constructed independently.

Scheme 7. Determination of Mg 2+ content in solution

An aliquot of 25.00 cm 3 of the test solution containing calcium and magnesium is placed in a conical titration flask with a capacity of 250 cm 3, ~ 60 cm 3 of distilled water, ~ 10 cm 3 of ammonia buffer solution with pH ~ 9.5–10 are added, and an indicator is added. eriochrome black T about 0.05 g
(at the tip of the spatula). In this case, the solution turns wine red. The solution in the flask is slowly titrated with EDTA solution until the color changes from wine red to green. Titration results ( V 2) entered into the table. 8.

Table 8

Results of titration of a solution containing calcium and magnesium

Experience no.	Volume of the test solution, cm 3	Volume of EDTA, V∑, cm 3	Mg 2+ content in solution, g
	25,00
	25,00
	25,00
	25,00
	25,00

Reagents

1. EDTA solution with a molar concentration of ~ 0.05 mol/dm 3.

2. Standard solution of Cu(II) with a titer of 2.00×10 -3 g/dm 3 .

3. H 2 O distilled.

4. Ammonia buffer with pH ~ 8 – 8.5.

5. Murexide, indicator mixture.

Tasks

1. Calculate α 4 for EDTA at pH=5, if the ionization constants of EDTA are as follows: K 1 =1.0·10 -2, K 2 =2.1·10 -3, K 3 =6.9·10 -7 , K 4 =5.5·10 -11.

2. Plot a titration curve for 25.00 ml of 0.020 M nickel solution with 0.010 M EDTA solution at pH = 10, if the stability constant
K NiY = 10 18.62. Calculate p after adding 0.00; 10.00; 25.00; 40.00; 50.00 and 55.00 ml titrant.

3. For titration of 50.00 ml of solution containing calcium ions
and magnesium, it took 13.70 ml of 0.12 M EDTA solution at pH=12 and 29.60 ml at pH=10. Express the concentrations of calcium and magnesium in solution in mg/ml.

4. When analyzing 1 liter of water, 0.2173 g of calcium oxide and 0.0927 g of magnesium oxide were found. Calculate what volume of EDTA with a concentration of 0.0500 mol/l was spent on titration.

5. To titrate 25.00 ml of a standard solution containing 0.3840 g of magnesium sulfate, 21.40 ml of Trilon B solution was consumed. Calculate the titer of this solution for calcium carbonate and its molar concentration.

6. Based on the formation constants (stability) of metal complexonates given below, evaluate the possibility of complexometric titration of metal ions at pH = 2; 5; 10; 12.

7. When titrating a 0.01 M solution of Ca 2+ with a 0.01 M solution of EDTA at pH = 10, the stability constant K CaY = 10 10.6. Calculate what the conditional stability constant of the metal complex with the indicator should be at pH=10 if = at the end point of titration.

8. The acid ionization constant of the indicator used in complexometric titration is 4.8·10 -6. Calculate the content of acidic and alkaline forms of the indicator at pH = 4.9, if its total concentration in the solution is 8.0·10 -5 mol/l. Determine the possibility of using this indicator when titrating a solution
with pH=4.9, if the color of its acid form matches the color of the complex.

9. To determine the aluminum content in the sample, a 550 mg sample was dissolved and 50.00 ml of a 0.05100 M solution of complexone III was added. The excess of the latter was titrated with 14.40 ml of 0.04800 M zinc(II) solution. Calculate the mass fraction of aluminum in the sample.

10. When destroying a complex containing bismuth and iodide ions, the latter are titrated with a solution of Ag(I), and bismuth with complexone III.
To titrate a solution containing 550 mg of sample, 14.50 ml of 0.05000 M solution of complexone III is required, and to titrate the iodide ion contained in 440 mg of sample, 23.25 ml of 0.1000 M Ag(I) solution is required. Calculate the coordination number of bismuth in the complex if iodide ions are the ligand.

11. A sample weighing 0.3280 g containing Pb, Zn, Cu was dissolved
and transferred to a 500.00 cm 3 volumetric flask. The determination was carried out in three stages:
a) for the titration of the first portion of a solution with a volume of 10.00 cm 3 containing Pb, Zn, Cu, 37.50 cm 3 of 0.0025 M EDTA solution was spent; b) in the second portion with a volume of 25.00 cm 3, Cu was masked, and 27.60 cm 3 EDTA was used for titration of Pb and Zn; c) in the third portion with a volume of 100.00 cm 3 Zn was masked
and Cu, 10.80 cm 3 EDTA was spent on the titration of Pb. Determine the mass fraction of Pb, Zn, Cu in the sample.

Titration curves

In redoxmetry, titration curves are plotted in coordinates E = f(C R),
they illustrate graphically the change in system potential during the titration process. Before the equivalence point, the potential of the system is calculated by the ratio of the concentrations of the oxidized and reduced forms of the analyte (because before the equivalence point, one of the titrant forms is practically absent), after the equivalence point - by the ratio of the concentrations of the oxidized and reduced forms of the titrant (because after the equivalence point, the analyte is titrated almost completely).

The potential at the equivalence point is determined by the formula

, (26)

where is the number of electrons participating in half-reactions;

– standard electrode potentials of half-reactions.

In Fig. Figure 10 shows the titration curve of a solution of oxalic acid H 2 C 2 O 4 with a solution of potassium permanganate KMnO 4 in an acidic medium
( = 1 mol/dm3).

Rice. 10. Titration curve for 100.00 cm 3 oxalic solution

acids H 2 C 2 O 4 s C 1/ z= 0.1000 mol/dm 3 permanganate solution

potassium KMnO 4 s C 1/ z= 0.1000 mol/dm 3 at = 1 mol/dm 3

Half-reaction potential MnO 4 - + 5 e+ 8H + → Mn 2+ + 4H 2 O depends on the pH of the medium, since hydrogen ions participate in the half-reaction.

Permanganatometry

The titrant is a solution of potassium permanganate KMnO 4, which is a strong oxidizing agent. Basic equation:

MnO 4 - +8H + + 5e = Mn 2+ + 4H 2 O, =+1.51 V.

M 1/ z (KMnO 4)= g/mol.

In slightly acidic, neutral and slightly alkaline environments, due to the lower redox potential, the permanganate ion is reduced to Mn +4.

MnO 4 - +2H 2 O + 3e = MnO 2 ¯ + 4OH - , = +0.60 V.

M 1/ z (KMnO 4) = 158.03/3 = 52.68 g/mol.

In an alkaline environment, a solution of potassium permanganate is reduced
up to Mn +6.

MnO 4 - + 1e = MnO 4 2-, = +0.558 V.

M 1/ z (KMnO 4) = 158.03 g/mol.

To eliminate side reactions, titration with potassium permanganate is carried out in an acidic environment, which is created with sulfuric acid. It is not recommended to use hydrochloric acid to create a medium, since potassium permanganate can oxidize the chloride ion.

2Cl - – 2e = Cl 2 , = +1.359 V.

Potassium permanganate is most often used in the form of a solution
with a molar equivalent concentration of ~ 0.05 – 0.1 mol/dm 3 . It is not a primary standard due to the fact that aqueous solutions of potassium permanganate are capable of oxidizing water and organic impurities in it:

4MnO 4- + 2H 2 O = 4MnО 2 ¯+ 3O 2 + 4OH -

The decomposition of potassium permanganate solutions is accelerated in the presence of manganese dioxide. Since manganese dioxide is a product of the decomposition of permanganate, this precipitate has autocatalytic effect to the decomposition process.

Solid potassium permanganate used to prepare solutions is contaminated with manganese dioxide, so it is impossible to prepare a solution from an accurate sample. In order to obtain a sufficiently stable solution of potassium permanganate, after dissolving a sample of KMnO 4 in water, it is left in a dark bottle for several days (or boiled), and then the MnO 2 is separated by filtration through glass filter (a paper filter cannot be used, as it reacts with potassium permanganate to form manganese dioxide).

The color of the potassium permanganate solution is so intense that
that an indicator is not required in this method. In order to give a noticeable pink color to 100 cm 3 of water, 0.02 - 0.05 cm 3 of KMnO 4 solution is sufficient
with a molar concentration equivalent of 0.1 mol/dm 3 (0.02 M). The color of potassium permanganate at the end point of titration is unstable and gradually discolors as a result of the interaction of excess permanganate
with manganese(II) ions present at the end point in relatively large quantities:

2MnO 4 - + 3Mn 2+ + 2H 2 O « 5MnО 2 ¯ + 4H +

Standardization of working solution KMnO 4 is carried out with sodium oxalate or oxalic acid (freshly recrystallized and dried at 105°C).

Use solutions of primary standards with a molar concentration equivalent WITH(½ Na 2 C 2 O 4) = 0.1000 or 0.05000 mol/l.

C 2 O 4 2- – 2e ® 2CO 2 , = -0.49 V

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Plan

1. The essence of precipitation titration

2. Argentometric titration

3. Thiocyanatometric titration

4. Application of precipitation titration

4.1 Preparation of a standardized solution of silver nitrate

4.2 Preparation of standardized ammonium thiocyanate solution

4.3 Determination of chlorine content in a sample according to Volhard

4.4 Determination of sodium trichloroacetate content in a technical preparation

1. The essence of precipitationtitration

The method combines titrimetric determinations based on the formation reactions of poorly soluble compounds. Only certain reactions that satisfy certain conditions are suitable for these purposes. The reaction must proceed strictly according to the equation and without side processes. The resulting precipitate must be practically insoluble and fall out fairly quickly, without the formation of supersaturated solutions. In addition, it is necessary to be able to determine the end point of the titration using an indicator. Finally, the phenomena of adsorption (co-precipitation) must be expressed during titration so weakly that the result of the determination is not distorted.

The names of individual precipitation methods come from the names of the solutions used. The method using a solution of silver nitrate is called argentometry. This method determines the content of C1~ and Br~ ions in neutral or slightly alkaline media. Thiocyanatometry is based on the use of a solution of ammonium thiocyanate NH 4 SCN (or potassium KSCN) and serves to determine traces of C1- and Br~, but in highly alkaline and acidic solutions. It is also used to determine the silver content in ores or alloys.

The expensive argentometric method for determining halogens is gradually being replaced by mercurometric method. In the latter, a solution of mercury (I) nitrate Hg 2 (NO 3) 2 is used.

Let us consider in more detail argentometric and thiocyanatometric titrations.

2. Argentometric titration

The method is based on the reaction of precipitation of C1~ and Br~ ions by silver cations with the formation of poorly soluble halides:

Cl-+Ag+=AgClb Br^- + Ag+= AgBr

In this case, a solution of silver nitrate is used. If a substance is analyzed for silver content, then a solution of sodium (or potassium) chloride is used. titration solution drug

Titration curves are of great importance for understanding the argentometry method. As an example, consider the case of titration of 10.00 ml of 0.1 N. sodium chloride solution 0.1 N. solution of silver nitrite (without taking into account changes in the volume of the solution).

Before titration begins, the concentration of chloride ions in the solution is equal to the total concentration of sodium chloride, i.e. 0.1 mol/l or = --lg lO-i = 1.

When 9.00 ml of silver nitrate solution is added to a titrated solution of sodium chloride and 90% of chloride ions are precipitated, their concentration in the solution will decrease 10 times and become equal to N0~ 2 mol/l, and pCl will be equal to 2. Since the value nPAgci= IQ- 10, the concentration of silver ions will be:

10-yu/[C1-] = Yu-Yu/10-2 = 10-8 M ol/l, OR pAg= -- lg = -- IglO-s = 8.

All other points to construct the titration curve are calculated in a similar manner. At the equivalence point pCl=pAg= = 5 (see table).

Table Changes in pC\ and pAg during titration of 10.00 ml 0.1 N. sodium chloride solution 0.1 N. silver nitrate solution

AgNO 3 solution was added,


9.99 10.00 (eq. point) 10.01	yu-4 yu-5 yu-6.	yu- 6 yu- 5 yu-*

The jump interval during argentometric titration depends on the concentration of solutions and on the value of the solubility product of the precipitate. The smaller the PR value of the compound obtained as a result of titration, the wider the jump interval on the titration curve and the easier it is to record the end point of the titration using an indicator.

The most common argentometric determination of chlorine is the Mohr method. Its essence consists in direct titration of the liquid with a solution of silver nitrate with the indicator potassium chromate until the white precipitate turns brown.

The indicator of Mohr's method - a solution of K2CrO 4 with silver nitrate gives a red precipitate of silver chromate Ag 2 CrO 4, but the solubility of the precipitate (0.65-10~ 4 E/l) is much greater than the solubility of silver chloride (1.25X_X10~ 5 E/l ). Therefore, when titrating with a solution of silver nitrate in the presence of potassium chromate, a red precipitate of silver chromate appears only after adding an excess of Ag+ ions, when all chloride ions have already been precipitated. In this case, a solution of silver nitrate is always added to the liquid being analyzed, and not vice versa.

The possibilities for using argentometry are quite limited. It is used only when titrating neutral or slightly alkaline solutions (pH from 7 to 10). In an acidic environment, the silver chromate precipitate dissolves.

In strongly alkaline solutions, silver nitrate decomposes with the release of insoluble oxide Ag 2 O. The method is also unsuitable for analyzing solutions containing the NH^ ion, since in this case an ammonia complex + is formed with the Ag + cation - The analyzed solution should not contain Ba 2 +, Sr 2+, Pb 2+, Bi 2+ and other ions that precipitate with potassium chromate.Nevertheless, argentometry is convenient for the analysis of colorless solutions containing C1~ and Br_ ions.

3. Thiocyanatometric titration

Thiocyanatometric titration is based on the precipitation of Ag+ (or Hgl+) ions with thiocyanates:

Ag+ + SCN- = AgSCN|

For determination, a solution of NH 4 SCN (or KSCN) is required. Determine Ag+ or Hgi+ by direct titration with a thiocyanate solution.

Thiocyanatometric determination of halogens is performed using the so-called Volhard method. Its essence can be expressed in diagrams:

CI- + Ag+ (excess) -* AgCI + Ag+ (residue), Ag+ (residue) + SCN~-> AgSCN

In other words, an excess of a titrated solution of silver nitrate is added to the liquid containing C1~. Then the AgNO 3 residue is backtitrated with a thiocyanate solution and the result is calculated.

The indicator of the Volhard method is a saturated solution of NH 4 Fe(SO 4) 2 - 12H 2 O. While there are Ag+ ions in the titrated liquid, the added SCN~ anions are associated with the release of AgSCN precipitate, but do not interact with Fe 3+ ions. However, after the equivalence point, the slightest excess of NH 4 SCN (or KSCN) causes the formation of blood-red 2 + and + ions. Thanks to this, it is possible to determine the equivalent point.

Thiocyanatometric determinations are used more often than argentometric ones. The presence of acids does not interfere with titration using the Volhard method and even contributes to obtaining more accurate results, since the acidic environment suppresses the hydrolysis of the Fe salt**. The method makes it possible to determine the C1~ ion not only in alkalis, but also in acids. The determination is not hampered by the presence of Ba 2 +, Pb 2 +, Bi 3 + and some other ions. However, if the analyzed solution contains oxidizing agents or mercury salts, then the use of Volhard’s method becomes impossible: oxidizing agents destroy the SCN- ion, and the mercury cation precipitates it.

The alkaline test solution is neutralized before titration with nitric acid, otherwise the Fe 3 + ions included in the indicator will precipitate iron (III) hydroxide.

4. Applications of precipitation titration

4.1 Preparation of a standardized solution of silver nitrate

The primary standards for standardizing a silver nitrate solution are sodium or potassium chlorides. Prepare a standard solution of sodium chloride and approximately 0.02 N. silver nitrate solution, standardize the second solution to the first.

Preparation of a standard sodium chloride solution. A solution of sodium chloride (or potassium chloride) is prepared from chemically pure salt. The equivalent mass of sodium chloride is equal to its molar mass (58.45 g/mol). Theoretically, to prepare 0.1 l 0.02 n. solution requires 58.45-0.02-0.1 = 0.1169 g NaCl.

Take a sample of approximately 0.12 g of sodium chloride on an analytical balance, transfer it to a 100 ml volumetric flask, dissolve, bring the volume to the mark with water, and mix well. Calculate the titer and normal concentration of the original sodium chloride solution.

Preparation: 100 ml approximately 0.02 N. silver nitrate solution. Silver nitrate is a scarce reagent, and usually its solutions have a concentration of no higher than 0.05 N. 0.02 n is quite suitable for this work. solution.

During argentometric titration, the equivalent mass of AgN0 3 is equal to the molar mass, i.e. 169.9 g/mol. Therefore, 0.1 l 0.02 n. the solution should contain 169.9-0.02-0.1 = 0.3398 g AgNO 3. However, it makes no sense to take exactly this sample, since commercial silver nitrate always contains impurities. Weigh approximately 0.34 - 0.35 g of silver nitrate on a technochemical scale; weigh the solution in a small amount of water into a 100 ml volumetric flask and adjust the volume with water; store the solution in the flask, wrap it in black paper and pour into a dark glass flask. silver and prepare it for titration. Rinse the pipette with sodium chloride solution and transfer 10.00 ml of the solution into a conical flask. Add 2 drops of a saturated solution of potassium chromate and carefully, drop by drop, titrate with a solution of silver nitrate while stirring. Make sure that the color of the mixture changes from yellow to reddish due to one excess drop of silver nitrate. After repeating the titration 2-3 times, take the average of the convergent readings and calculate the normal concentration of the silver nitrate solution.

Let us assume that for titration of 10.00 ml 0.02097 n. sodium chloride solution, an average of 10.26 ml of silver nitrate solution was used. Then

A^ AgNOj. 10.26 = 0.02097. 10.00, AT AgNOs = 0.02097- 10.00/10.26 = 0.02043

If it is intended to determine the content of C1~ in the sample, then calculate, in addition, the titer of the silver nitrate solution with respect to chlorine: T, - = 35.46-0.02043/1000 = 0.0007244 g/ml, “l this means that 1 ml of silver nitrate solution corresponds to 0.0007244 g of titrated chlorine.

4.2 Preparation of a standardized ammonium thiocyanate solutionI

A solution of NH 4 SCN or KSCN with a precisely known titer cannot be prepared by dissolving a sample, since these salts are very hygroscopic. Therefore, a solution with an approximate normal concentration is prepared and adjusted to a standardized solution of silver nitrate. The indicator is a saturated solution of NH 4 Fe(SO 4) 2 - 12H 2 O. To prevent hydrolysis of the Fe salt, 6 N is added to the indicator itself and to the analyzed solution before titration. nitric acid.

Preparation: 100 ml approximately 0.05 N. ammonium thiocyanate solution. The equivalent mass of NH4SCN is equal to its molar mass, i.e. 76.12 g/mol. Therefore, 0.1 l 0.05 n. solution should contain 76.12.0.05-0.1=0.3806 g NH 4 SCN.

Take a sample of about 0.3-0.4 g on an analytical balance, transfer it to a 100 ml flask, dissolve, bring the volume of the solution to the mark with water and mix.

Standardization of ammonium thiocyanate solution with silver nitrate. Prepare a burette for titration with NH 4 SCN solution. Rinse the pipette with the silver nitrate solution and measure 10.00 ml of it into the conical flask. Add 1 ml of NH 4 Fe(SO 4)2 solution (indicator) and 3 ml. 6 n. nitric acid. Slowly, with continuous shaking, pour in the NH 4 SCN solution from the burette. Stop titration after the appearance of a brown-pink color 2+, which does not disappear with vigorous shaking.

Repeat the titration 2-3 times, take the average from the converging readings and calculate the normal concentration of NH 4 SCN.

Let us assume that for titration of 10.00 ml 0.02043 n. silver nitrate solution, an average of 4.10 ml of NH 4 SCN solution was used.

4.3 Definitioncontentchlorine in the sample according to Volhard

Volhard halogens are determined by back titration of the silver nitrate residue with a solution of NH 4 SCN. However, accurate titration is only possible here if measures are taken to prevent (or slow down) the reaction between silver chloride and excess ferric thiocyanate:

3AgCI + Fe (SCN) 3 = SAgSCNJ + FeCl 3

in which the color that appears first gradually disappears. It is best to filter off the AgCl precipitate before titrating the excess silver nitrate with NH 4 SCN solution. But sometimes, instead, some organic liquid is added to the solution, which is not mixed with water and, as it were, isolates the ApCl precipitate from excess nitrate.

Determination method. Take a test tube with a solution of the analyte containing sodium chloride. Dissolve a sample of the substance in a 100 ml volumetric flask and bring the volume of the solution to the mark with water (the chloride concentration in the solution should be no more than 0.05 N).

Pipette 10.00 ml of the test solution into a conical flask, add 3 ml of 6 N. nitric acid and pour in a known excess of AgNO 3 solution from the burette, for example 18.00 ml. Then filter off the silver chloride precipitate. Titrate the remaining silver nitrate with NH 4 SCN solution as described in the previous paragraph. After repeating the determination 2-3 times, take the average. If the silver chloride precipitate has been filtered, it should be washed and the washing water should be added to the filtrate.

Let us assume that the sample weight was 0.2254 g. To 10.00 ml of the analyzed solution, 18.00 ml of 0.02043 N was added. silver nitrate solution. To titrate the excess, 5.78 ml * 0.04982 N was used. NH 4 SCN solution.

First of all, let's calculate what volume is 0.02043 n. solution of silver nitrate corresponds to 5.78 ml of 0.04982 N spent on titration. NH 4 SCN solution:

therefore, 18.00 - 14.09 = 3.91 ml of 0.2043 N was used to precipitate the C1~ ion. silver nitrate solution. From here it is easy to find the normal concentration of sodium chloride solution.

Since the equivalent mass of chlorine is 35.46 g/mol,* the total mass of chlorine in the sample is:

772=0.007988-35.46-0.1 =0.02832 g.

0.2254 g C1-- 100%

x = 0.02832-100/0.2254 = 12.56%.:

0.02832 > C1 -- x%

The Volhard method is also used to determine the content of Br~ and I- ions. In this case, it is not necessary to filter out the precipitates of silver bromide or iodide. But it must be taken into account that the Fe 3 + ion oxidizes iodides to free iodine. Therefore, the indicator is added after all the I- ions have been precipitated by silver nitrate.

4.4 Determination of trichl contentOsodium racetate| in technical preparation (for chlorine)

Technical sodium trichloroacetate (TCA) is a herbicide for killing cereal weeds. It is a white or light brown crystalline substance, highly soluble in water. According to Volhard, the mass fraction of organic chloride compounds is first determined, and then after the destruction of chlorine. From the difference, the mass fraction (%) of sodium trichloroacetate chlorine is found.

Determination of the mass fraction (%) of chlorine inorganic compounds. Place an exact weighed portion of the drug (2-2.5 g) in a 250 ml volumetric flask, dissolve, bring the solution to the mark with water, and mix. Pipette 10 ml of solution into a conical flask and add 5-10 ml of concentrated nitric acid.

Add 5 or 10 ml of 0.05 N from the burette. silver nitrate solution and titrate the excess with 0.05 N. a solution of NH 4 SCN in the presence of NH 4 Fe(SO 4) 2 (indicator).

Calculate the mass fraction (%) of chlorine (x) of inorganic compounds using the formula

(V -- l/i) 0.001773-250x100

where V is the volume exactly 0.05 N. AgNO 3 solution taken for analysis; Vi -- volume exactly 0.05 N. NH 4 SCN solution, used for titration of excess AgNO 3; t—a sample of sodium trichloroacetate; 0.001773 -- mass of chlorine corresponding to 1 ml of 0.05 N. AgNO solution. Determination of mass fraction (%) of total chlorine. Take 10 ml of the previously prepared solution into a conical flask, add 10 ml of a solution with a mass fraction of NaOH 30% and 50 ml of water. Connect the flask to a reflux condenser and boil its contents for 2 hours. Allow the liquid to cool, rinse the condenser with water, collecting the washing water in the same flask. Add 20 ml of diluted (1:1) nitric acid to the solution and add 30 ml of 0.05 N from a burette. silver nitrate solution. Titrate excess silver nitrate to 0.05 N. a solution of NH 4 SCN in the presence of NH 4 Fe(SO 4)2. Calculate the mass fraction (%) of total chlorine (xi) using the above formula. Find the mass fraction (%) of sodium trichloroacetate in the preparation (x^) using the formula

x2 = (x1 -- x) (185.5/106.5),

where 185.5 is the molar mass of sodium trichloroacetate; 106.5 -- mass of chlorine contained in the molar mass of sodium trichloroacetate.

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Titration methods in analytical chemistry briefly. Neutralization method in the titrimetric method of analysis

Principle of analysis

Rationale for the process

Titration end point detection

Titration Reaction Requirements

Types of analysis

Acid-base titration

Redox titration

Complexometric titration

Precipitation titration

Titration methods

Back titration

Titration of the substituent

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