Chlorine forms 4 oxygen-containing acids: hypochlorous, chlorous, hypochlorous and perchloric.

Hypochlorous acid (HClO) is formed by the interaction of chlorine with water, as well as its salts with strong mineral acids. It is a weak acid and is very unstable. The composition of the products of its decomposition reaction depends on the conditions. With strong illumination of hypochlorous acid, the presence of a reducing agent in the solution, as well as prolonged standing, it decomposes with the release of atomic oxygen:

HClO = HCl + O.

In the presence of water-removing substances, chlorine oxide (I) is formed:

2 HClO = 2 H 2 O + Cl 2 O.

3 HClO = 2 HCl + HClO 3.

Therefore, when chlorine interacts with a hot alkali solution, salts are formed not of hydrochloric and hypochlorous acids, but of hydrochloric and hypochlorous acids:

6 NaOH + 3 Cl 2 = 5 NaCl + NaClO 3 + 3 H 2 O.

Salts of hypochlorous acid - hypochlorites - are very strong oxidizing agents. They are formed when chlorine reacts with alkalis in the cold. At the same time, salts of hydrochloric acid are formed. Of these mixtures, the most widely used are bleach and javel water.

Hypochlorous acid (HClO 3) is formed by the action of its salts - chlorates - with sulfuric acid. This is a very unstable acid, a very strong oxidizing agent. Can only exist in dilute solutions.

By evaporating a solution of HClO 3 at low temperature in a vacuum, you can obtain a viscous solution containing about 40% perchloric acid. At higher acid contents, the solution decomposes explosively.

Explosive decomposition also occurs at lower concentrations in the presence of reducing agents. In dilute solutions, perchloric acid exhibits oxidizing properties, and the reactions proceed quite calmly:

HClO 3 + 6 HBr = HCl + 3 Br 2 + 3 H 2 O.

Salts of perchloric acid are formed during the electrolysis of chloride solutions in the absence of a diaphragm between the cathode and anode spaces, as well as during the dissolution of chlorine in a hot alkali solution, as shown above. Potassium chlorate (Berthollet salt) formed during electrolysis is slightly soluble in water and is easily separated from other salts in the form of a white precipitate. Like acid, chlorates are fairly strong oxidizing agents:

KClO 3 + 6 HCl = KCl + 3 Cl 2 + 3 H 2 O.

Chlorates are used for the production of explosives, as well as for the production of oxygen in the laboratory and salts of perchloric acid - perchlorates. When Berthollet salt is heated in the presence of manganese dioxide (MnO2), which plays the role of a catalyst, oxygen is released. If you heat potassium chlorate without a catalyst, it decomposes to form potassium salts of hydrochloric and perchloric acids:

2 KClO 3 = 2 KCl + 3 O 2;

4 KClO 3 = KCl + 3 KClO 4.

By treating perchlorates with concentrated sulfuric acid, perchloric acid can be obtained:

KClO 4 + H 2 SO 4 = KHSO 4 + HClO 4.

This is the strongest acid. It is the most stable of all oxygen-containing chlorine acids, however, anhydrous acid can decompose explosively when heated, shaken or in contact with reducing agents. Dilute solutions of perchloric acid are quite stable and safe to use. Chlorates of potassium, rubidium, cesium, ammonium and most organic bases are poorly soluble in water.

In industry, potassium perchlorate is obtained by electrolytic oxidation of Berthollet salt:

2 H + + 2 e - = H 2 (at the cathode);

ClO 3 - - 2 e - + H 2 O = ClO4 - + 2 H + (at the anode).

Chlorous acid (HClO 2) is formed by the action of concentrated sulfuric acid on alkali metal chlorites, which are obtained as intermediate products during the electrolysis of solutions of alkali metal chlorides in the absence of a diaphragm between the cathode and anode spaces. It is a weak, unstable acid, a very strong oxidizing agent in an acidic environment. When it interacts with hydrochloric acid, chlorine is released:

HClO 2 + 3 HCl = Cl 2 + 2 H 2 O.

Sodium chlorites are used to produce chlorine dioxide, for water disinfection, and also as a bleaching agent.

Chloric, or bleaching, lime (CaOCl 2), or CaCl (ClO), is formed by the interaction of chlorine with powdered calcium hydroxide - fluff:

Ca(OH) 2 + Cl 2 = Cl-O-Ca-Cl + H 2 O,

2 Ca(OH) 2 + 2 Cl 2 = CaCl 2 + Ca(OCl) 2 + 2 H 2 O.

The quality of bleach is determined by the hypochlorite content in it. It has very strong oxidizing properties and can even oxidize manganese salts to permanganate:

5 CaOCl 2 + 2 Mn(NO 3) 2 + 3 Ca(OH) 2 = Ca(MnO 4) 2 + 5 CaCl 2 + 2 Ca(NO 3) 2 + 3 H 2 O.

Under the influence of carbon dioxide contained in the air, it decomposes with the release of chlorine:

CaOCl 2 + CO 2 = CaCO 3 + Cl 2,

CaCl 2 + Ca(OCl) 2 + 2 CO 2 = 2 CaCO 3 + 2 Cl 2.

Bleach is used as a bleaching and disinfectant.

Chemistry tutor

Continuation. See in No. 22/2005; 1, 2, 3, 5, 6, 8, 9, 11, 13, 15, 16, 18, 22/2006;
3, 4, 7, 10, 11, 21/2007;
2, 7, 11, 18, 19, 21/2008;
1, 3/2009

LESSON 29

10th grade(first year of study)

Halogens and their most important compounds

1. Position in the table of D.I. Mendeleev, structure of the atom.

2. Origin of names.

3. Physical properties.

4. Chemical properties (using the example of chlorine).

5. Being in nature.

6. Basic methods of production (using the example of chlorine).

7. Hydrogen chloride and chlorides.

8. Oxygen-containing chlorine acids and their salts.

The halogens (“solenols”) are located in the VIIa subgroup of the periodic table. These include fluorine, chlorine, bromine, iodine and astatine. All halogens are R-elements, have the configuration of the external energy level ns 2 p 5 . Since at the outer level of the halogen atoms there is 1 unpaired R-electron, the characteristic valence is I. In addition to fluorine, the number of unpaired electrons in the atoms of all halogens in an excited state can increase, so valences III, V and VII are possible.

Cl: 1 s 2 2s 2 2p 6 3s 2 3p 5 3d 0 (valence I),

Cl*: 1 s 2 2s 2 2p 6 3s 2 3p 4 3d 1 (valence III),

Cl**: 1 s 2 2s 2 2p 6 3s 2 3p 3 3d 2 (valence V),

Cl***: 1 s 2 2s 2 2p 6 3s 1 3p 3 3d 3 (valency VII).

Halogens are typical nonmetals and exhibit oxidizing properties. The oxidation state of halogens in compounds with metals and hydrogen is –1; in all oxygen-containing compounds, halogens (except fluorine) exhibit oxidation states +1, +3, +5, +7, for example:

Down the subgroup, the aggregative state of halogens changes, solubility in water decreases, the radius of the atom increases, electronegativity, non-metallic properties and oxidizing ability decrease (fluorine is the strongest oxidizing agent). For halogen compounds: from Cl – to I – the reducing ability of halide ions increases. In the series of oxygen-free and oxygen-containing acids, the acidic properties increase:

The name fluorine comes from the Greek word - destructive, since hydrofluoric acid, from which they tried to obtain fluorine, corrodes glass. Chlorine gets its name from the Greek word yellow-green, the color of fading leaves. Bromine is named for the smell of liquid bromine from the Greek word - foul. The name iodine comes from the Greek word - violet - for the color of iodine vapor. Radioactive astatine is named from the Greek word - unstable.

According to the physical properties, fluorine is a difficult to liquefy gas of light green color, chlorine is an easily liquefiable gas of yellow-green color, bromine is a heavy liquid of red-brown color, iodine is a solid crystalline substance of dark purple color with a metallic luster, easily sublimated (sublimation). All halogens, except iodine, have a strong suffocating odor and are toxic.

Chemical properties

All halogens exhibit high chemical activity, which decreases when moving from fluorine to iodine. Let's look at the chemical properties of halogens using chlorine as an example:

(F 2 - with explosion; Br 2, I 2 - in light and at elevated temperature.)

Metals (+):

2Na + Cl 2 = 2NaCl;

2Fe + 3Cl 2 2FeCl 3 .

Non-metals (+/–):*

N 2 + Cl 2 reaction does not occur.

Basic oxides (–).

Acidic oxides (–).

Bases (+/–):

Acids (+/–):

2HBr + Cl 2 = 2HCl + Br 2,

HCl + Br 2 reaction does not occur.

Salts (+/–):

2KBr + Cl 2 = 2KCl + Br 2,

KCl + Br 2 reaction does not occur.

In nature, halogens are not found in free form due to their high chemical activity. Among the most common chlorine compounds are rock or table salt (NaCl), sylvinite (KCl NaCl), carnallite (KCl MgCl 2). Large amounts of chlorides are found in sea water. Chlorine is part of chlorophyll. Natural chlorine consists of two isotopes 35 Cl and 37 Cl. We emphasize that in the case of chlorine, the number of neutrons in an atom can only be calculated for each isotope separately:

35 Cl, p = 17, e = 17, n = 35 – 17 = 18;

37 Cl, p = 17,e = 17, n = 37 – 17 = 20.

Industrially, chlorine is obtained by electrolysis of an aqueous solution or melt of chloride:

Laboratory methods for obtaining (the effect of concentrated hydrochloric acid on various oxidizing agents):

MnO 2 + 4HCl (conc.) = MnCl 2 + Cl 2 + 2H 2 O,

2KMnO 4 + 16HCl (conc.) = 2MnCl 2 + 5Cl 2 + 2KCl + 8H 2 O,

KClO 3 + 6HCl (conc.) = KCl + 3Cl 2 + 3H 2 O,

K 2 Cr 2 O 7 + 14HCl (conc.) = 2CrCl 3 + 3Cl 2 + 2KCl + 7H 2 O,

Ca(ClO) 2 + 4HCl (conc.) = CaCl 2 + 2Cl 2 + 2H 2 O.

C h o l o r d i d e s

Hydrogen chloride(HCl) is a colorless gas with a pungent odor, heavier than air, highly soluble in water (450 volumes of hydrogen chloride dissolve in 1 volume of water). The molecule is formed according to the type of covalent polar bond. An aqueous solution of hydrogen chloride is called hydrochloric acid. Concentrated hydrochloric acid “smoke” in air; the maximum concentration of hydrogen chloride in the solution is 35–36%. It is a strong acid that exhibits all the characteristic properties of acids:

HCl H + + Cl – ,

2HCl + Zn = ZnCl 2 + H 2,

HCl + Cu reaction does not occur,

2HCl + CaO = CaCl 2 + H 2 O,

HCl + NaOH = NaCl + H 2 O,

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

A qualitative reaction to hydrochloric acid and its salts (chlorides) is the reaction with a solution of silver nitrate:

Ag + + Cl – -> AgCl,

AgNO 3 + NaCl -> AgCl + NaNO 3.

Hydrogen chloride can be obtained:

Direct synthesis from hydrogen and chlorine (synthetic method):

The action of concentrated sulfuric acid on solid chlorides - the sulfate method (HF can be obtained similarly, but HBr and HI cannot be obtained):

NaCl (solid) + H 2 SO 4 (conc.) = HCl + NaHSO 4.

As the oxidation state of chlorine increases, the strength of the acids increases sharply. Thus, hypochlorous acid is very weak (weaker than carbonic acid), and perchloric acid is the strongest of all known acids.

A c l o d e r s a l ts

Acid oxides	Cl2O	Cl2O3	Cl2O5	Cl2O7
Acids	HClO Hypochlorous	HClO2 Chloride	HClO 3 Chloric	HClO 4 Chloric
Graphic formulas acids	H–O–Cl	H–O–Cl=O
Names and examples of salts	Sodium hypochlorite NaClO	Sodium chlorite NaClO2	Sodium chlorate NaClO3	Sodium perchlorate NaClO4

Hypochlorous acid(HClO) – weak, very unstable.

Salts of this acid (hypochlorites) are very strong oxidizing agents. The most widely used is a mixed salt of hydrochloric and hypochlorous acids - calcium chloride-hypochlorite (bleach):

Chloric acid(HClO 3) - exists only in dilute solutions. The acid itself and its salts (chlorates) are strong oxidizing agents. The best known salt of this acid is potassium chlorate (Berthollet's salt).

5KClO 3 + 6P = 3P 2 O 5 + 5KCl,

KClO 3 + 3MnO 2 + 6KOH = KCl + 3K 2 MnO 4 + 3H 2 O,

4KClO 3 + 3K 2 S = 4KCl + 3K 2 SO 4.

Many salts of oxygen-containing chlorine acids are thermally unstable, for example:

2KClO 3 2KCl + 3O 2,

4KClO 3 3KClO 4 + KCl (without catalyst),

3KClO KClO 3 + 2KCl,

KClO 4 KCl + 2O 2.

Test on the topic “Halogens and their most important compounds”

1. The gas has a density of 3.485 g/l at a pressure of 1.2 atm and a temperature of 25 °C. Determine the formula of the gas.

a) Fluorine; b) chlorine;

c) hydrogen bromide;

d) hydrogen chloride.

2. The phenomenon of transition of a substance from a solid to a gaseous state, bypassing the liquid state, is called:

a) condensation; b) sublimation;

c) sublimation; d) distillation.

3. Natural chlorine is a mixture of isotopes with mass numbers 35 and 37. Calculate the isotopic composition of chlorine, taking its relative atomic mass to be 35.5.

a) 75% and 25%;

b) 24.4% and 75.8%;

c) 50% and 50%;

d) there is not enough data to solve the problem.

4. Chlorine can be obtained by electrolysis:

a) molten potassium chloride;

b) potassium chloride solution;

c) molten copper chloride;

d) copper chloride solution.

5. A solution of hydrogen fluoride in water is called:

a) Javel water;

b) hydrofluoric acid;

c) bleaching lime;

d) hydrofluoric acid.

6. Chlorine(V) oxide is the anhydride of the following acid:

a) hypochlorous; b) chloric;

c) chloride; d) chlorine.

7. When calcining Berthollet salt in the presence of manganese dioxide as a catalyst, the following are formed:

a) potassium chloride and oxygen;

b) potassium perchlorate and potassium chloride;

c) potassium perchlorate and ozone;

d) potassium hypochlorite and chlorine.

8. A solution of sodium iodide was added to an acidified solution containing 0.543 g of a certain salt, which contains lithium, chlorine and oxygen, until the release of iodine ceased. The mass of released iodine was 4.57 g. Name of the original salt:

a) lithium hypochlorite; b) lithium chlorite;

c) lithium chlorate; d) lithium perchlorate.

9. In halogen molecules the chemical bond is:

a) covalent polar;

b) covalent nonpolar;

c) ionic;

d) donor-acceptor.

10. Chlorine, unlike fluorine, under certain conditions may react with:

a) water; b) hydrogen;

c) copper; d) sodium hydroxide.

Key to the test

1	2	3	4	5	6	7	8	9	10
b	b, c	A	a B C D	b, d	b	A	V	b	G

Problems and exercises on halogens and their compounds

Chain of transformation

1. Potassium chloride -> chlorine -> hydrogen chloride -> calcium chloride -> hydrogen chloride -> chlorine -> potassium chlorate.

2. Chlorine -> bertholet's salt -> potassium chloride -> hydrochloric acid + manganese dioxide + water -> chlorine -> copper(II) chloride -> chlorine.

3. Potassium chloride -> chlorine -> potassium chlorate -> potassium chloride -> potassium.

4. Potassium chloride -> chlorine -> hydrogen chloride -> chlorine -> potassium hypochlorite.

5. Sodium chloride -> hydrogen chloride -> chlorine -> bertholet's salt -> potassium chloride -> potassium hydroxide -> potassium hypochlorite.

6. Potassium chlorate -> A -> B-> C -> A -> potassium nitrate (substances A, B, C contain chlorine, the first three transformations are redox reactions).

7. Calcium oxide -> calcium hydroxide -> bleaching lime -> calcium chloride -> calcium.

8. Sodium bromide -> sodium chloride -> chlorine -> bleach -> calcium carbonate -> calcium bicarbonate -> carbon dioxide.

9. Sodium iodide -> iodine -> potassium iodide -> silver iodide.

10. Potassium hypochlorite -> potassium chlorate -> potassium perchlorate -> potassium chloride.

Level A

1. A vessel with 200 g of chlorine water was kept in direct sunlight and the released gas was collected, the volume of which at normal conditions was amounted to 0.18 l. Determine the composition of chlorine water (mass fraction of chlorine).

Answer. 0,57 %.

2. The gas obtained by calcination of 9.8 g of berthollet salt is mixed with the gas obtained at the anode as a result of complete electrolysis of the melt of 22.2 g of calcium chloride. The resulting gas mixture was passed through 400 g of a 2% hot sodium hydroxide solution. Determine the composition of the resulting solution.

Answer. 2.38% NaCl; 0.84% ​​NaClO 3 .

3. Calculate the mass of salt and volume of gas (no.s.) formed during the decomposition of 17 g of salt, which colors the burner flame yellow and contains 27.06% metal, 16.47% nitrogen and 56.47% oxygen. What mass of bertholite salt will be required to produce the same amount of gas?

Answer. 13.8 g NaNO 2; 2.24 l O 2; 8.13 g KClO 3 .

4. What volume of chlorine (no.) can be obtained from 1 m 3 of solution (density 1.23 g/cm 3) containing 20.7% sodium chloride and 4.3% magnesium chloride?

Answer. 61.2 m3.

5. The gas released at the anode during the electrolysis of 200 g of a 20% sodium chloride solution was passed through 400 g of a 30% potassium bromide solution. An excess of silver nitrate solution was added to the resulting solution. Determine the quantitative composition of the precipitate.

Answer. 59.4 g AgBr; 98.154 g AgCl.

Level B

1. 1.3 liters of chlorine were passed through a tube containing a powdered mixture of sodium chloride and iodide weighing 3 g at a temperature of 42 °C and a pressure of 101.3 kPa. The substance obtained in the tube was calcined at 300 °C, leaving 2 g of the substance. Determine the mass fractions of salts in the original mixture.

Answer. 45.3% NaCl; 54.6% NaI.

2. A mixture of magnesium iodide and zinc iodide was treated with excess bromine water, and the resulting solution was evaporated. The mass of the dry residue turned out to be 1.445 times less than the mass of the original mixture. How many times will the mass of the precipitate obtained after treating the same mixture with excess sodium carbonate be less than the mass of the original mixture?

Answer. 2.74 times.

3. To oxidize 2.17 g of alkaline earth metal sulfite, chlorine water containing 1.42 g of chlorine was added. An excess of potassium bromide was added to the resulting mixture, and 1.6 g of bromine was released. Determine the composition of the sediment contained in the mixture and calculate its mass.

(BaSO 4) = (BaSO 3) = 0.01 mol,

m(BaSO 4) = (BaSO 4) M(BaSO 4) = 0.01 233 = 2.33 g.

Answer. 2.33 g BaSO 4 .

4. A current was passed through 800 g of a 10% aqueous solution of sodium chloride. After the salt electrolysis process was completed, all the gas released at the anode was absorbed by the hot solution resulting from electrolysis. Determine the composition of the solution obtained after gas absorption.

Answer. In a solution of 8.35% NaCl and
3.03% NaClO 3 .

5. The density of a mixture of chlorine and hydrogen at a pressure of 0.2 atm and a temperature of 27 °C is 0.0894 g/l. Hydrogen chloride, obtained by explosion of 100 l (n.s.) of such a mixture, was dissolved in 500 g of 10% hydrochloric acid. Find the mass fraction of hydrogen chloride in the resulting solution.

Answer. 17 %.

HONEST TASKS

1. Name substances A, B and C if they are known to undergo reactions described in the schemes below; write the complete reaction equations for these schemes:

A + H 2 -> B,

A + H 2 O B + C,

A + H 2 O + SO 2 -> B + ...,

C -> B + … .

Answer. Substances: A – Cl 2,
B – HCl; C – HClO.

2. Gas A, under the influence of concentrated sulfuric acid, is converted into a simple substance B, which reacts with hydrosulfide acid to form a simple substance C and a solution of the starting substance A. Identify the substances, write the reaction equations.

Answer. Substances: A – HBr; B – Br 2 ; C–S.

3. When chlorine is passed through a solution of strong acid A, a simple substance B is released and the solution becomes dark in color. With further passage of chlorine, substance B is converted into acid C and the solution becomes discolored. Name the substances A, B and C, write the reaction equations.

Answer. Substances: A – HI; B – I 2, C – HIO 3.

4. Give examples of reactions during which complete reduction of free bromine occurs: a) in an acidic aqueous solution; b) in an alkaline aqueous solution; c) in the gas phase.

Answer. Reaction equations:

5. What substances reacted and under what conditions, if as a result the following substances were formed (all products are indicated without coefficients): a) barium chloride and potassium hydroxide; b) calcium bromide and hydrogen bromide; c) potassium chloride and phosphorus pentoxide. Write complete reaction equations.

Answer. Reaction equations:

a) Ba(ClO) 2 + 2KH = BaCl 2 + 2KOH;

b) CaH 2 + 2Br 2 = CaBr 2 + 2HBr;

c) 5KClO 3 + 6P 5KCl + 3P 2 O 5.

6. For degassing, 254 g of bleach is required. The laboratory contains: calcium, manganese dioxide, sodium, zinc, sodium chloride, sulfuric acid, water, phosphorus, sulfur, barium sulfate. What reagents and in what quantity will be required? Write complete reaction equations.

Answer. 142 g Ca; 830.7 g NaCl; 308.85 g MnO 2;
1391.6 g H 2 SO 4.

Reaction equations:

Ca + 2H 2 O = Ca(OH) 2 + H 2,

NaCl (solid) + H 2 SO 4 (conc.) = HCl + NaHSO 4,

MnO 2 + 4HCl = Cl 2 + MnCl 2 + 2H 2 O,

2Cl 2 + 2Ca(OH) 2 Ca(ClO) 2 + CaCl 2 + 2H 2 O.

7. Freshly prepared chlorine water is added dropwise to an aqueous solution of potassium iodide. Explain why the initially appearing color of the solution then disappears. Support your answer with reaction equations.

Answer. Reaction equations:

2KI + Cl 2 = 2KCl + I 2,

I 2 + 5Cl 2 + 6H 2 O = 2HIO 3 + 10HCl.

* The +/– sign means that this reaction does not occur with all reagents or under specific conditions.

To be continued

Chlorine forms four oxygen-containing acids: hypochlorous, chlorous, hypochlorous and perchloric.

Hypochlorous acid HClO is formed by the interaction of chlorine with water, as well as its salts with strong mineral acids. It is a weak acid and is very unstable. The composition of the products of its decomposition reaction depends on the conditions. With strong illumination of hypochlorous acid, the presence of a reducing agent in the solution, and also long-term standing, it decomposes with the release of atomic oxygen: HClO = HCl + O

In the presence of water-removing substances, chlorine oxide (I) is formed: 2 HClO = 2 H2O + Cl2O

Therefore, when chlorine interacts with a hot alkali solution, salts are formed not of hydrochloric and hypochlorous acids, but of hydrochloric and hypochlorous acids: 6 NaOH + 3 Cl2 = 5 NaCl + NaClO3 + 3 H2O

Hypochlorous acid salts- very strong oxidizing agents. They are formed when chlorine reacts with alkalis in the cold. At the same time, salts of hydrochloric acid are formed. Of these mixtures, the most widely used are bleach and javel water.

Chlorous acid HClO2 is formed by the action of concentrated sulfuric acid on alkali metal chlorites, which are obtained as intermediate products during the electrolysis of solutions of alkali metal chlorides in the absence of a diaphragm between the cathode and anode spaces. It is a weak, unstable acid, a very strong oxidizing agent in an acidic environment. When it interacts with hydrochloric acid, chlorine is released: HClO2 + 3 HCl = Cl2 + 2 H2O

Hypochlorous acid HClO3 is formed by the action of its salts - chlorates- sulfuric acid. This is a very unstable acid, a very strong oxidizing agent. Can only exist in dilute solutions. By evaporating a solution of HClO3 at low temperature in a vacuum, you can obtain a viscous solution containing about 40% perchloric acid. At higher acid contents, the solution decomposes explosively. Explosive decomposition also occurs at lower concentrations in the presence of reducing agents. In dilute solutions, perchloric acid exhibits oxidizing properties, and the reactions proceed quite calmly:

HClO3 + 6 HBr = HCl + 3 Br2 + 3 H2O

Salts of perchloric acid - chlorates - are formed during the electrolysis of chloride solutions in the absence of a diaphragm between the cathode and anode spaces, as well as when chlorine is dissolved in a hot alkali solution, as shown above. Potassium chlorate (Berthollet salt) formed during electrolysis is slightly soluble in water and is easily separated from other salts in the form of a white precipitate. Like acid, chlorates are fairly strong oxidizing agents:

KClO3 + 6 HCl = KCl + 3 Cl2 + 3 H2O

Chlorates are used for the production of explosives, as well as for the production of oxygen in laboratory conditions and salts of perchloric acid - perchlorates. When Berthollet salt is heated in the presence of manganese dioxide MnO2, which plays the role of a catalyst, oxygen is released. If you heat potassium chlorate without a catalyst, it decomposes to form potassium salts of hydrochloric and perchloric acids:

2 KClO3 = 2 KCl + 3 O2

4 KClO3 = KCl + 3 KClO4

By treating perchlorates with concentrated sulfuric acid, perchloric acid can be obtained:

KClO4 + H2SO4 = KHSO4 + HClO4

This is the strongest acid. It is the most stable of all oxygen-containing chlorine acids, but anhydrous acid can decompose explosively when heated, shaken, or in contact with reducing agents. Dilute solutions of perchloric acid are quite stable and safe to use. Chlorates of potassium, rubidium, cesium, ammonium and most organic bases are poorly soluble in water.

In industry, potassium perchlorate is obtained by electrolytic oxidation of Berthollet salt:

2 H+ + 2 e- = H2 (at the cathode)

ClO3- - 2 e- + H2O = ClO4- + 2 H+ (at the anode)

Biological role.

It belongs to the vital irreplaceable elements. In the human body 100 g.

Chlorine ions play a very important biological role. Entering together with ions K+, Mg2+, Ca2+, HCO~, H3PO4 and proteins, they play a dominant role in creating a certain level of osmotic pressure (osmotic homeostasis) of blood plasma, lymph, cerebrospinal fluid, etc.

Chlorine ion is involved in the regulation of water-salt metabolism and the volume of fluid retained by tissues, maintaining the pH of intracellular fluid and the membrane potential created by the operation of the sodium-potassium pump, which is explained (as in the case of its participation in osmosis) by the ability to diffuse through cell membranes like the way Na+ and K+ ions do this. Chlorine ion is a necessary component (together with H2PO4, HSO4 ions, enzymes, etc.) of gastric juice, which is part of hydrochloric acid.

By promoting digestion, hydrochloric acid also destroys a variety of pathogenic bacteria.

15.1. General characteristics of halogens and chalcogens

Halogens (“generating salts”) are elements of group VIIA. These include fluorine, chlorine, bromine and iodine. This group also includes unstable, and therefore not found in nature, astatine. Sometimes hydrogen is also included in this group.
Chalcogens (“copper-producing”) are elements of the VIA group. These include oxygen, sulfur, selenium, tellurium and polonium, which is practically not found in nature.
Of the eight atoms existing in nature elements of these two groups the most common oxygen atoms ( w= 49.5%), followed by chlorine atoms in abundance ( w= 0.19%), then – sulfur ( w= 0.048%), then fluorine ( w= 0.028%). The atoms of other elements are hundreds and thousands of times smaller. You already studied oxygen in eighth grade (Chapter 10); of the other elements, the most important are chlorine and sulfur - you will get acquainted with them in this chapter.
The orbital radii of the atoms of halogens and chalcogens are small and only the fourth atoms of each group approach one angstrom. This leads to the fact that all of these elements are non-metal forming elements and only tellurium and iodine show some signs of amphotericity.
The general valence electronic formula of halogens is ns 2 n.p. 5 , and chalcogens – ns 2 n.p. 4 . The small size of atoms does not allow them to give up electrons; on the contrary, the atoms of these elements tend to accept them, forming singly charged (for halogens) and doubly charged (for chalcogens) anions. By combining with small atoms, the atoms of these elements form covalent bonds. Seven valence electrons enable halogen atoms (except fluorine) to form up to seven covalent bonds, and six valence electrons of chalcogen atoms - up to six covalent bonds.
In fluorine compounds, the most electronegative element, only one oxidation state is possible, namely –I. Oxygen, as you know, has a maximum oxidation state of +II. For atoms of other elements, the highest oxidation state is equal to the group number.

The simple substances of group VIIA elements are of the same type in structure. They consist of diatomic molecules. Under normal conditions, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. According to their chemical properties, these substances are strong oxidizing agents. Due to the increase in the size of atoms with increasing atomic number, their oxidative activity decreases.
Of the simple substances of group VIA elements, under normal conditions only oxygen and ozone are gaseous, consisting of diatomic and triatomic molecules, respectively; the rest are solids. Sulfur consists of eight-atom cyclic molecules S 8, selenium and tellurium from polymer molecules Se n and Te n. In terms of their oxidative activity, chalcogens are inferior to halogens: only oxygen is a strong oxidizing agent, while the rest exhibit oxidizing properties to a much lesser extent.

Compound hydrogen compounds halogens (HE) fully complies with the general rule, and chalcogens, in addition to ordinary hydrogen compounds of the composition H 2 E, can also form more complex hydrogen compounds of the composition H 2 E n chain structure. In aqueous solutions, both hydrogen halides and other chalcogen hydrogens exhibit acidic properties. Their molecules are acid particles. Of these, only HCl, HBr and HI are strong acids.
For halogens formation oxides uncharacteristic, most of them are unstable, but higher oxides of the composition E 2 O 7 are known for all halogens (except for fluorine, the oxygen compounds of which are not oxides). All halogen oxides are molecular substances; their chemical properties are acidic oxides.
In accordance with their valence capabilities, chalcogens form two series of oxides: EO 2 and EO 3. All these oxides are acidic.

Hydroxides of halogens and chalcogens are oxoacids.

Make abbreviated electronic formulas and energy diagrams of atoms of elements of groups VIA and VIIA. Indicate the outer and valence electrons.

Chlorine is the most common and therefore the most important of the halogens.
In the earth's crust, chlorine is found in the minerals: halite (rock salt) NaCl, sylvite KCl, carnallite KCl MgCl 2 6H 2 O and many others. The main industrial method of production is electrolysis of sodium or potassium chlorides.

The simple substance chlorine is a greenish gas with a pungent, suffocating odor. At –101 °C it condenses into a yellow-green liquid. Chlorine is very poisonous; during the First World War they even tried to use it as a chemical warfare agent.
Chlorine is one of the most powerful oxidizing agents. It reacts with most simple substances (exceptions: noble gases, oxygen, nitrogen, graphite, diamond and some others). As a result, halides are formed:
Cl 2 + H 2 = 2HCl (when heated or exposed to light);
5Cl 2 + 2P = 2PCl 5 (when burned in excess of chlorine);
Cl 2 + 2Na = 2NaCl (at room temperature);
3Cl 2 + 2Sb = 2SbCl 3 (at room temperature);
3Cl 2 + 2Fe = 2FeCl 3 (when heated).
In addition, chlorine can oxidize many complex substances, for example:
Cl 2 + 2HBr = Br 2 + 2HCl (in the gas phase and in solution);
Cl 2 + 2HI = I 2 + 2HCl (in the gas phase and in solution);
Cl 2 + H 2 S = 2HCl + S (in solution);
Cl 2 + 2KBr = Br 2 + 2KCl (in solution);
Cl 2 + 3H 2 O 2 = 2HCl + 2H 2 O + O 2 (in concentrated solution);
Cl 2 + CO = CCl 2 O (in the gas phase);
Cl 2 + C 2 H 4 = C 2 H 4 Cl 2 (in the gas phase).
In water, chlorine is partially dissolved (physically), and partially reacts reversibly with it (see § 11.4 c). With a cold solution of potassium hydroxide (and any other alkali), a similar reaction occurs irreversibly:

Cl 2 + 2OH = Cl + ClO + H 2 O.

As a result, a solution of potassium chloride and hypochlorite is formed. When reacted with calcium hydroxide, a mixture of CaCl 2 and Ca(ClO) 2 is formed, called bleach.

With hot concentrated solutions of alkalis, the reaction proceeds differently:

3Cl 2 + 6OH = 5Cl + ClO 3 + 3H 2 O.

When reacted with KOH, this produces potassium chlorate, called Berthollet salt.
Hydrogen chloride is the only hydrogen connection chlorine This colorless gas with a suffocating odor is highly soluble in water (it completely reacts with it, forming oxonium ions and chloride ions (see § 11.4). Its solution in water is called hydrochloric or hydrochloric acid. This is one of the most important products of chemical technology, since Hydrochloric acid is consumed in many industries. It is also of great importance for humans, in particular because it is contained in gastric juice, facilitating the digestion of food.
Hydrogen chloride was previously produced industrially by burning chlorine in hydrogen. Currently, the need for hydrochloric acid is almost completely satisfied through the use of hydrogen chloride, formed as a by-product during the chlorination of various organic substances, for example, methane:

CH 4 + Cl 2 = CH 3 + HCl

And laboratories produce hydrogen chloride from sodium chloride by treating it with concentrated sulfuric acid:
NaCl + H 2 SO 4 = HCl + NaHSO 4 (at room temperature);
2NaCl + 2H 2 SO 4 = 2HCl + Na 2 S 2 O 7 + H 2 O (when heated).
Higher oxide chlorine Cl 2 O 7 – colorless oily liquid, molecular substance, acidic oxide. As a result of reaction with water, it forms perchloric acid HClO 4, the only chlorine oxoacid that exists as an individual substance; the remaining chlorine oxoacids are known only in aqueous solutions. Information about these chlorine acids is given in Table 35.

Table 35. Chlorine acids and their salts

C/O chlorine	Formula acids	Name acids	Force acids	Name salts
		hydrochloric
		hypochlorous		hypochlorites
		chloride
		hypochlorous
				perchlorates

Most chlorides are soluble in water. The exceptions are AgCl, PbCl 2, TlCl and Hg 2 Cl 2. Formation of a colorless precipitate of silver chloride when silver nitrate solution is added to the test solution – qualitative reaction for chloride ion:

Ag + Cl = AgCl

Chlorine can be obtained from sodium or potassium chlorides in the laboratory:

2NaCl + 3H 2 SO 4 + MnO 2 = 2NaHSO 4 + MnSO 4 + 2H 2 O + Cl 2

As an oxidizing agent when producing chlorine using this method, you can use not only manganese dioxide, but also KMnO 4, K 2 Cr 2 O 7, KClO 3.
Sodium and potassium hypochlorites are included in various household and industrial bleaches. Bleach is also used as a bleach and is also used as a disinfectant.
Potassium chlorate is used in the production of matches, explosives and pyrotechnic compositions. When heated, it decomposes:
4KClO 3 = KCl + 3KClO 4;
2KClO 3 = 2KCl + O 2 (in the presence of MnO 2).
Potassium perchlorate also decomposes, but at a higher temperature: KClO 4 = KCl + 2O 2.

1. Compose molecular equations for reactions for which ionic equations are given in the text of the paragraph.
2. Write down equations for the reactions given in the text of the paragraph descriptively.
3. Make up reaction equations characterizing the chemical properties of a) chlorine, b) hydrogen chloride (and hydrochloric acid), c) potassium chloride and d) barium chloride.
Chemical properties of chlorine compounds

Various allotropic modifications are stable under different conditions element sulfur. Under normal conditions simple substance sulfur is a yellow, brittle crystalline substance consisting of eight-atomic molecules:

This is the so-called orthorhombic sulfur (or -sulfur) S 8. (The name comes from a crystallographic term characterizing the symmetry of the crystals of this substance). When heated, it melts (113 ° C), turning into a mobile yellow liquid consisting of the same molecules. With further heating, cycles are broken and very long polymer molecules are formed - the melt darkens and becomes very viscous. This is the so-called -sulfur S n. Sulfur boils (445 °C) in the form of diatomic molecules S 2, similar in structure to oxygen molecules. The structure of these molecules, like that of oxygen molecules, cannot be described within the framework of the covalent bond model. In addition, there are other allotropic modifications of sulfur.
In nature there are deposits of native sulfur, from which it is extracted. Most of the mined sulfur is used to produce sulfuric acid. Some of the sulfur is used in agriculture to protect plants. Purified sulfur is used in medicine to treat skin diseases.
From hydrogen compounds sulfur, the most important is hydrogen sulfide (monosulfan) H 2 S. It is a colorless poisonous gas with the smell of rotten eggs. It is slightly soluble in water. Dissolution is physical. To a small extent, protolysis of hydrogen sulfide molecules occurs in an aqueous solution and, to an even lesser extent, the resulting hydrosulfide ions (see Appendix 13). However, a solution of hydrogen sulfide in water is called hydrogen sulfide acid (or hydrogen sulfide water).

Hydrogen sulfide burns in air:

2H 2 S + 3O 2 = 2H 2 O + SO 2 (with excess oxygen).

A qualitative reaction to the presence of hydrogen sulfide in the air is the formation of black lead sulfide (blackening of filter paper moistened with a solution of lead nitrate:

H 2 S + Pb 2 + 2H 2 O = PbS + 2H 3O

The reaction proceeds in this direction due to the very low solubility of lead sulfide.

In addition to hydrogen sulfide, sulfur also forms other sulfanes H 2 S n, for example, disulfan H 2 S 2, similar in structure to hydrogen peroxide. It is also a very weak acid; its salt is pyrite FeS 2.

In accordance with the valence capabilities of its atoms, sulfur forms two oxide: SO 2 and SO 3 . Sulfur dioxide (commonly known as sulfur dioxide) is a colorless gas with a pungent odor that causes coughing. Sulfur trioxide (the old name is sulfuric anhydride) is a solid, extremely hygroscopic, non-molecular substance that turns into a molecular substance when heated. Both oxides are acidic. When reacting with water, they form sulfur dioxide and sulfur dioxide, respectively. acids.
In dilute solutions, sulfuric acid is a typical strong acid with all its characteristic properties.
Pure sulfuric acid, as well as its concentrated solutions, are very strong oxidizing agents, and the oxidizing atoms here are not hydrogen atoms, but sulfur atoms, moving from the +VI oxidation state to the +IV oxidation state. As a result, when reacting with concentrated sulfuric acid, sulfur dioxide is usually formed, for example:

Cu + 2H 2 SO 4 = CuSO 4 + SO 2 + 2H 2 O;
2KBr + 3H 2 SO 4 = 2KHSO 4 + Br 2 + SO 2 + 2H 2 O.

Thus, even metals that are in the voltage series to the right of hydrogen (Cu, Ag, Hg) react with concentrated sulfuric acid. At the same time, some fairly active metals (Fe, Cr, Al, etc.) do not react with concentrated sulfuric acid, this is due to the fact that a dense protective film is formed on the surface of such metals under the influence of sulfuric acid, preventing further oxidation. This phenomenon is called passivation.
Being a dibasic acid, sulfuric acid forms two rows salts: medium and sour. Acid salts are isolated only for alkaline elements and ammonium; the existence of other acid salts is questionable.
Most medium sulfates are soluble in water and, since the sulfate ion is practically not an anionic base, do not undergo anion hydrolysis.
Modern industrial methods for the production of sulfuric acid are based on the production of sulfur dioxide (1st stage), its oxidation into trioxide (2nd stage) and the interaction of sulfur trioxide with water (3rd stage).

Sulfur dioxide is produced by burning sulfur or various sulfides in oxygen:

S + O 2 = SO 2;
4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2.

The process of roasting sulfide ores in non-ferrous metallurgy is always accompanied by the formation of sulfur dioxide, which is used to produce sulfuric acid.
Under normal conditions, it is impossible to oxidize sulfur dioxide with oxygen. Oxidation is carried out by heating in the presence of a catalyst - vanadium(V) or platinum oxide. Even though the reaction

2SO 2 + O 2 2SO 3 + Q

reversible, yield reaches 99%.
If the resulting gas mixture of sulfur trioxide and air is passed through clean water, most of the sulfur trioxide is not absorbed. To prevent losses, the gas mixture is passed through sulfuric acid or its concentrated solutions. This produces disulfuric acid:

SO 3 + H 2 SO 4 = H 2 S 2 O 7.

A solution of disulfuric acid in sulfuric acid is called oleum and is often represented as a solution of sulfur trioxide in sulfuric acid.
By diluting oleum with water, you can obtain both pure sulfuric acid and its solutions.

1.Create structural formulas
a) sulfur dioxide, b) sulfur trioxide,
c) sulfuric acid, d) disulfuric acid.

At increasing st.OK . chlorine acid stability too growing .

The increase in stability is explained by:

A) hardening bonds in anions due to a decrease in the number of NEPs in chlorine,

b) increasing ratio the number of π-overlaps to the number of σ-bonds from 0/1 in ClO - to 3/4 in ClO - 4. Compare the graphic formulas of acids:

H – O - Cl, H - O - Cl = O, H – O – Cl = O H – O – Cl = O

c) increases from HClO to HClO 4 symmetry anion (both due to an increase

number of oxygen atoms, and as a result of a decrease polarizing actions

hydrogen due to weakening of its bond with the anion).

d) decreases attack angle chlorine atom (i.e. its spatial accessibility for interaction).

Acid properties of halogen hydroxides. Acid-base properties

of any hydroxide depend on the ratio of the bond strengths H - O and O - E in

fragment H - O - E. Obviously, the greater the electronegativity of the element, the more the electron density from the H - O bond is shifted to the O - E bond

(H − O − E) and especially the hydroxide exhibits acidic properties.

Therefore, an important factor is nature halogen. Thus, when moving from chlorine to iodine, in accordance with a decrease in the value of E.O. the acidic properties of hydroxides are reduced. Moreover, so much so that hypoiodic acid dissociates according to acidic type in less degree НIO → Н + + IO - (K d = 4 ∙10 − 13),

than according to the main one: IOH → I + + OH − (K d = 3 ∙10 − 10).

Even a neutralization reaction is possible (but reversible): IOH + HNO 3 → INO 3 + H 2 O.

Salts of chlorine acids, as more stable (than acids) compounds, all

isolated in a free state, but also their activity increases with decreasing temperature. Cl. Thus, KClO 3 (Berthollet salt) oxidizes iodide ions only in an acidic environment, and KClO - in a neutral environment.

2.8.1. Hypochlorous acid HCl +1 O H–O–Cl (hypochlorites)

Physical properties. Exists only in the form of dilute aqueous solutions.

Receipt.

Cl 2 + H 2 O ↔ HCl + HClO

Chemical properties.

HClO is a weak acid and a strong oxidizing agent:

1) Decomposes, releasing atomic oxygen

HClO – in the light → HCl + O HClO – vol. conventional → H 2 O + Cl 2 O НClO --- t → НCl + НClO 3

2) Gives salts with alkalis - hypochlorites

HClO + KOH → KClO + H 2 O CaOCl 2 – bleaching lime (bleach)

CaOCl 2 + CO 2 + H 2 O → CaCO 3 + CaCl 2 + HClO (HCl + O)

3) with a strong reducing agent HI

2HI + HClO → I 2 ↓ + HCl + H 2 O

2.8.2. Chlorous acid HCl +3 O 2 H–O–Cl=O (chlorites)

Physical properties. Exists only in aqueous solutions.

Receipt

It is formed by the interaction of hydrogen peroxide with chlorine oxide (IV), which is obtained from Berthollet salt and oxalic acid in an H 2 SO 4 environment:

2KClO 3 + H 2 C 2 O 4 + H 2 SO 4 → K 2 SO 4 + 2CO 2 + 2СlO 2 + 2H 2 O

2ClO 2 + H 2 O 2 → 2HClO 2 + O 2

Chemical properties

HClO 2 is a weak acid and a strong oxidizing agent.

1)HClO 2 + KOH → KClO 2 + H 2 O

KClO 2 + KI + H 2 SO 4 → I 2 + KCl + K 2 SO 4 + H 2 O

2) Unstable, decomposes during storage

4HClO 2 → HCl + HClO 3 + 2ClO 2 + H 2 O

5HClO 2 ---t→ 3HClO 3 + Cl 2 + H 2 O

2.8.3. Hypochlorous acid HCl +5 O 3 (chlorates)

Physical properties: Stable only in aqueous solutions.

Receipt: Ba (ClO 3) 2 + H 2 SO 4 → 2HClO 3 + BaSO 4 ↓

Chemical properties

HClO 3 - Strong acid and strong oxidizing agent; salts of perchloric acid -

chlorates:

6P + 5HClO 3 → 3P 2 O 5 + 5HCl HClO 3 + KOH → KClO3+H2O

- KClO 3 - Berthollet salt; it is obtained by passing chlorine through a heated (40°C) KOH solution: 3Cl 2 + 6KOH → 5KCl + KClO 3 + 3H 2 O

Berthollet's salt is used as an oxidizing agent; When heated, it decomposes:

4KClO 3 – without cat → KCl + 3KClO 4 2KClO 3 – MnO2 cat → 2KCl + 3O 2

2.8.4. Perchloric acid HCl +7 O 4 (perchlorates)

Physical properties: Colorless liquid, boiling point. = 25°C, temperature = -101°C.

Receipt: KClO 4 + H 2 SO 4 → KHSO 4 + HClO 4

Chemical properties:

HClO 4 is a very strong acid and a very strong oxidizing agent;

salts of perchloric acid - perchlorates .

1) HClO 4 + KOH → KClO 4 + H 2 O

2) When heated, perchloric acid and its salts decompose:

4HClO 4 – t° → 4ClO 2 + 3O 2 + 2H 2 O KClO 4 – t° → KCl + 2O 2

Hydrogen bromide HBr (BROMIDE)

Physical properties

Colorless gas, highly soluble in water; t°boil. = -67°C; t°pl. = -87°C.

Receipt

1) 2NaBr + H 3 PO 4 – t ° → Na 2 HPO 4 + 2HBr 2) PBr 3 + 3H 2 O → H 3 PO 3 + 3HBr

Chemical properties

An aqueous solution of hydrogen bromide is hydrobromic acid, which is even stronger than hydrochloric acid. She undergoes the same reactions as HCl

1) Dissociation: HBr ↔ H+ + Br -

2) With metals in the voltage series up to hydrogen:

Mg + 2HBr → MgBr 2 + H 2

3) with metal oxides:

CaO + 2HBr → CaBr 2 + H 2 O

4) with bases and ammonia:

NaOH + HBr → NaBr + H 2 O Fe(OH) 3 + 3HBr → FeBr 3 + 3H 2 O NH 3 + HBr → NH 4 Br

5) with salts

MgCO 3 + 2HBr → MgBr 2 + H 2 O + CO 2

Qualitative reaction: AgNO 3 + HBr → AgBr↓ + HNO 3

The formation of a yellow precipitate of silver bromide, insoluble in acids, serves to detect the Br - anion in solution.

6) restorative properties:

2HBr + H 2 SO 4 (conc.) → Br 2 + SO 2 + 2H 2 O 2HBr + Cl 2 → 2HCl + Br 2

Among the oxygen acids of bromine are known

Weak bromide HBr +1 O and

Strong brominated HBr +5 O 3 .

Hydrogen iodide (iodides)

Physical properties: Colorless gas with a pungent odor, highly soluble in water,

t°boil. = -35°C; t°pl. = -51°C.

Receipt:

1) I 2 + H 2 S → S + 2HI 2) 2P + 3I 2 + 6H 2 O → 2H 3 PO 3 + 6HI

Chemical properties

1) A solution of HI in water - strong hydroiodic acid:

HI ↔ H + + I - 2HI + Ba(OH) 2 → BaI 2 + 2H 2 O

Salts of hydroiodic acid - iodides (for other HI reactions, see the properties of HCl and HBr)

2) HI is a very strong reducing agent:

2HI + Cl 2 → 2HCl + I 2

8HI + H 2 SO 4 (conc.) → 4I 2 + H 2 S + 4H 2 O

5HI + 6KMnO 4 + 9H 2 SO 4 → 5HIO 3 + 6MnSO 4 + 3K 2 SO4 + 9H 2 O

3)Qualitative reaction: The formation of a dark yellow precipitate of silver iodide, insoluble in acids, serves to detect the iodine anion in solution.

NaI + AgNO 3 → AgI↓ + NaNO 3 HI + AgNO 3 → AgI↓ + HNO 3

3.0.1. Oxygen acids of iodine ( iodates )

a) Hydrous acid HI +5 O 3

Colorless crystalline substance, melting point = 110°C, highly soluble in water.

Receive: 3I 2 + 10HNO 3 → 6HIO 3 + 10NO + 2H 2 O

HIO 3 is a strong acid (salts - iodates) and a strong oxidizing agent.

b) Iodic acid H 5 I +7 O 6

Crystalline hygroscopic substance, highly soluble in water,

t°pl. = 130°C. Weak acid (salts - periodates); strong oxidizing agent.