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A subgroup of oxygen, or chalcogens - the 6th group of the periodic system of D.I. Mendellev, including the following elements: O; S; Se; Te; Po. The group number indicates the maximum valency of the elements in this group. The general electronic formula of chalcogens is: ns2np4 - at the outer valence level, all elements have 6 electrons, which rarely give up and more often accept 2 missing electrons before the completion of the electron level. The presence of the same valence level determines the chemical similarity of chalcogens. Typical oxidation states: -1; -2; 0; +1; +2; +4; +6. Oxygen shows only -1 - in peroxides; -2 - in oxides; 0 - in a free state; +1 and +2 - in fluorides - O2F2, OF2 because it does not have a d-sub-level and electrons cannot be separated, and the valency is always 2; S - everything except +1 and -1. Sulfur has a d-sublevel and electrons with 3p and 3s in the excited state can separate and go to the d-sublevel. In the unexcited state, the valency of sulfur is 2 in SO, 4 in SO2, and 6 in SO3. Se+2; +4; +6, Te +4; +6, Po +2; -2. The valencies of selenium, tellurium and polonium are also 2, 4, 6. The values ​​of the oxidation states are reflected in the electronic structure of the elements: O - 2s22p4; S, 3s23p4; Se—4s24p4; Te—5s25p4; Po - 6s26p4. From top to bottom, with an increase in the external energy level, the physical and chemical properties of chalcogens naturally change: the radius of the atom of the elements increases, the ionization energy and electron affinity, as well as the electronegativity decrease; non-metallic properties decrease, metal properties increase (oxygen, sulfur, selenium, tellurium are non-metals), polonium has a metallic luster and electrical conductivity. Hydrogen compounds of chalcogens correspond to the formula: H2R: H2O, H2S, H2Se, H2Te are hydrogen chalcogens. Hydrogen in these compounds can be replaced by metal ions. The oxidation state of all chalcogens in combination with hydrogen is -2 and the valency is also 2. When hydrogen chalcogens are dissolved in water, the corresponding acids are formed. These acids are reducing agents. The strength of these acids increases from top to bottom, since the binding energy decreases and promotes active dissociation. Oxygen compounds of chalcogens correspond to the formula: RO2 and RO3 are acid oxides. When these oxides are dissolved in water, they form the corresponding acids: H2RO3 and H2RO4. In the direction from top to bottom, the strength of these acids decreases. H2RO3 are reducing acids, H2RO4 are oxidizing agents.

Oxygen is the most abundant element on earth. It makes up 47.0% of the mass of the earth's crust. Its content in the air is 20.95% by volume or 23.10% by mass. Oxygen is found in water, rocks, many minerals, salts, and is found in proteins, fats, and carbohydrates that make up living organisms. In the laboratory, oxygen is obtained: - decomposition by heating bertolet salt (potassium chlorate) in the presence of a catalyst MnO2: 2KClO3 = 2KCl + 3O2 - decomposition by heating potassium permanganate: 2KMnO4 = K2MnO4 + MnO2 + O2 In this case, very pure oxygen is obtained. oxygen can also be obtained by electrolysis of an aqueous solution of sodium hydroxide (electrodes are nickel); The main source of industrial production of oxygen is air, which is liquefied and then fractionated. First, nitrogen is released (tboil = -195°C), and almost pure oxygen remains in the liquid state, since its boiling point is higher (-183°C). A method of obtaining oxygen based on the electrolysis of water is widespread. Under normal conditions, oxygen is a colorless, tasteless and odorless gas, slightly heavier than air. It is slightly soluble in water (31 ml of oxygen dissolves in 1 liter of water at 20°C). At a temperature of -183°C and a pressure of 101.325 kPa, oxygen passes into a liquid state. Liquid oxygen has a bluish color and is drawn into a magnetic field. Natural oxygen contains three stable isotopes 168O (99.76%), 178O (0.04%) and 188O (0.20%). Three unstable isotopes - 148O, 158O, 198O were artificially obtained. To complete the external electronic level, the oxygen atom lacks two electrons. Taking them vigorously, oxygen exhibits an oxidation state of -2. However, in compounds with fluorine (OF2 and O2F2), the common electron pairs are shifted towards fluorine, as a more electronegative element. In this case, the oxidation states of oxygen are respectively +2 and +1, and of fluorine -1. The oxygen molecule consists of two O2 atoms. The chemical bond is covalent non-polar. Oxygen forms compounds with all chemical elements, except for helium, neon and argon. It interacts directly with most elements, except for halogens, gold and platinum. The rate of reaction of oxygen with both simple and complex substances depends on the nature of the substances, temperature, and other conditions. Such an active metal as cesium ignites spontaneously in atmospheric oxygen even at room temperature. Oxygen actively reacts with phosphorus when heated to 60 ° C, with sulfur - up to 250 ° C, with hydrogen - more than 300 ° C, with carbon (in the form of coal and graphite) - at 700-800 ° С. =2CO2+3H2OCH4+2O2=CO2+2H20 4FeS2+11O2=2Fe2O3+8SO2 The above reactions are accompanied by the release of both heat and light. Such processes involving oxygen are called combustion. In terms of relative electronegativity, oxygen is the second element. Therefore, in chemical reactions with both simple and complex substances, it is an oxidizing agent, tk. accepts electrons. Combustion, rusting, rotting and breathing proceed with the participation of oxygen. These are redox processes. To accelerate the oxidation processes, oxygen or air enriched with oxygen is used instead of ordinary air. Oxygen is used to intensify oxidative processes in the chemical industry (production of nitric acid, sulfuric acid, artificial liquid fuel, lubricating oils and other substances). The metallurgical industry consumes quite a lot of oxygen. Oxygen is used to produce high temperatures. The temperature of an oxygen-acetylene flame reaches 3500°C, an oxygen-hydrogen flame reaches 3000°C In medicine, oxygen is used to facilitate breathing. It is used in oxygen devices when working in an atmosphere difficult to breathe.

Sulfur- one of the few chemical elements that have been used by humans for several millennia. It is widely distributed in nature and occurs both in the free state (native sulfur) and in compounds. Minerals containing sulfur can be divided into two groups - sulfides (pyrites, shines, blendes) and sulfates. Native sulfur is found in large quantities in Italy (the island of Sicily) and the USA. In the CIS, there are deposits of native sulfur in the Volga region, in the states of Central Asia, in the Crimea and other regions. The minerals of the first group include lead luster PbS, copper luster Cu2S, silver luster - Ag2S, zinc blende - ZnS, cadmium blende - CdS, pyrite or iron pyrites - FeS2, chalcopyrite - CuFeS2, cinnabar - HgS. The minerals of the second group include gypsum CaSO4 2H2O, mirabilite (Glauber's salt) - Na2SO4 10H2O, kieserite - MgSO4 H2O. Sulfur is found in organisms of animals and plants, as it is part of protein molecules. Organic sulfur compounds are found in oil. Receipt 1. When obtaining sulfur from natural compounds, for example, from sulfur pyrite, it is heated to high temperatures. Sulfur pyrite decomposes with the formation of iron (II) sulfide and sulfur: FeS2=FeS+S 2. Sulfur can be obtained by oxidation of hydrogen sulfide with a lack of oxygen according to the reaction: 2H2S+O2=2S+2H2O3. Currently, it is common to obtain sulfur by carbon reduction of sulfur dioxide SO2 - a by-product in the smelting of metals from sulfur ores: SO2 + C \u003d CO2 + S4. Off-gases from metallurgical and coke ovens contain a mixture of sulfur dioxide and hydrogen sulfide. This mixture is passed at high temperature over a catalyst: H2S+SO2=2H2O+3S Sulfur is a lemon yellow brittle solid. It is practically insoluble in water, but highly soluble in carbon disulfide CS2 aniline and some other solvents. It conducts heat and electric current poorly. Sulfur forms several allotropic modifications: Natural sulfur consists of a mixture of four stable isotopes: 3216S, 3316S, 3416S, 3616S. Chemical properties A sulfur atom, having an incomplete external energy level, can attach two electrons and show an oxidation state of -2. Sulfur exhibits this oxidation state in compounds with metals and hydrogen (Na2S, H2S). When electrons are donated or pulled to an atom of a more electronegative element, the oxidation state of sulfur can be +2, +4, +6. In the cold, sulfur is relatively inert, but its reactivity increases with increasing temperature. 1. With metals, sulfur exhibits oxidizing properties. During these reactions, sulfides are formed (does not react with gold, platinum and iridium): Fe + S = FeS
2. Under normal conditions, sulfur does not interact with hydrogen, and at 150-200 ° C a reversible reaction occurs: H2 + S "H2S properties. S+3F2=SF6 (does not react with iodine)4. The combustion of sulfur in oxygen proceeds at 280°C, and in air at 360°C. This produces a mixture of SO2 and SO3:S+O2=SO2 2S+3O2=2SO35. When heated without air access, sulfur directly combines with phosphorus, carbon, showing oxidizing properties: 2P + 3S = P2S3 2S + C = CS26. When interacting with complex substances, sulfur behaves mainly as a reducing agent:

7. Sulfur is capable of disproportionation reactions. So, when sulfur powder is boiled with alkalis, sulfites and sulfides are formed: Sulfur is widely apply in industry and agriculture. About half of its production is used to produce sulfuric acid. Sulfur is used to vulcanize rubber: in this case, rubber turns into rubber. In the form of a sulfur color (fine powder), sulfur is used to combat diseases of the vineyard and cotton. It is used to obtain gunpowder, matches, luminous compositions. In medicine, sulfur ointments are prepared for the treatment of skin diseases.

31 Elements of IV A subgroup.

Carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb) - elements of group 4 of the main subgroup of PSE. On the outer electron layer, the atoms of these elements have 4 electrons: ns2np2. In the subgroup, with an increase in the ordinal number of the element, the atomic radius increases, non-metallic properties weaken, and metallic properties increase: carbon and silicon are non-metals, germanium, tin, lead are metals. Elements of this subgroup exhibit both positive and negative oxidation states: -4; +2; +4.

Element	Electric formula	rad nm	OEO	S.O.
C	2s 2 2p 2	0.077	2.5	-4; 0; +3; +4
14 Si	3s 2 3p 2	0.118	1.74	-4; 0; +3; +4
32ge	4s 2 4p 2	0.122	2.02	-4; 0; +3; +4
50 sn	5s 2 5p 2	0.141	1.72	0; +3; +4
82Pb	6s 2 6p 2	0.147	1.55	0; +3; +4

---------------------->(metallic properties increase)

GROUP VIA ELEMENTS

general characteristics

Oxygen, dioxygen, trioxygen

Oxygen compounds

Sulfur

Hydrogen sulfide. Sulfides

Sulfur oxygen compounds

Sulphuric acid

Other sulfur compounds

Selenium, tellurium, polonium and their compounds

GENERAL CHARACTERISTICS

Elements and their symbols: oxygen O, sulfur S, selenium Se, tellurium Te, polonium Po. Group name of the elements of the VIA group - chalcogens.

The degree of oxidation. For oxygen, the oxidation state is characteristic (-2), for the remaining elements (except polonium) - (+6), (+4) and (-2), polonium in compounds exhibits the oxidation state (+4), (+2) and ( -2). Oxidation stability (+V1) decreases from S to Te, oxidation state stability (+4) increases from S to Po, and oxidation state stability (-2) decreases from O to Po.

Properties(Table 1). The metallic properties increase from oxygen to polonium. In general, the elements O and S are non-metals; Se and Te show an increase in the metallic nature, for example, Se exists in free form in metallic and non-metallic modifications, and Te - only in metallic, Po - metal.

Hydroxides of elements of group VIA in the highest oxidation state correspond to acids H 2 SO 4, H 2 SeO 4 (strong acids) and H 6 TeO 6 (weak acid). The hydroxides of these elements in the oxidation state (+4) correspond to weak acids SO 2 * nH 2 O, H 2 SeO 3 and H 2 TeO 3, the strength of which decreases with an increase in the serial number of the acid-forming element, PoO (OH) 2 is an amphoteric hydroxide. Compared to the elements of the VA group, all of these hydroxides are more acidic, and compared to the elements of the VIIA group, they are more basic.

The stability of hydrogen compounds - hydrogen chalcogens H 2 O, H 2 S, H 2 Se, H 2 Te and H 2 Po - decreases from O to Po, their acidity in an aqueous solution, on the contrary, increases in this order. Water H 2 O is considered neutral, the strength of H 2 Te approximately corresponds to the strength of orthophosphoric acid, Chalcogen hydrogens exhibit respectively greater and less acidic properties than hydrogen compounds of the VA group elements and hydrogen halides.

OXYGEN, Dioxygen, trioxygen

Opening. Oxygen was first obtained in free form by heating saltpeter in 1770 (Scheele, Sweden) and in 1774 by decomposition of HgO oxide and red lead (Pb 2 II Pb IV)O 4 (Priestley, England). The role of oxygen in the combustion reactions of many substances in air was explained in 1775 (Lavoisier, France), which undermined the foundations of the theory of phlogiston put forward in 1697 (Stahl, Germany).

distribution in nature. Oxygen is the most abundant element on Earth. Its content in the earth's crust is 55.1% at. Free oxygen is found in the air (» 1.1 * 10 15 tons) and in natural waters (biochemical self-purification of river and sea water comes with oxygen consumption). Bound oxygen is found in water, silicates, quartz and other minerals, as well as in living organisms.

Atmospheric air composition: Nitrogen 78.09% (vol) 75.51% (mass); Oxygen 20.95 23.15; Argon 0.93 1.28; Carbon dioxide 0.03 0.046; Water vapor (25 °С)<3 <0,27.

The air density is 1.293 g/l at 0°C and 101.33 kPa (1 atm). The air shell of the Earth absorbs and neutralizes the harmful ultraviolet radiation of the Sun and protects the earth's surface from overheating.

Table 1.

Properties of chalcogens

	Oxygen O	Sulfur S	Selenium Se	Tellurium Te	Polonium Po
Element number	8	16	34	52	84
Relative atomic mass	15,999	32,067	78,96	127,60	208,982
Melting point °C	-219	119	217	450	254
Boiling point °C	-183	445	685	1390	962
Density at 20 ° C, g / cm 3	1.27 (TV)	2,1	4.8 (meth)	6,2	9,4
Oxidation state
+6	–	increase in stability.
+4	–	increased sustainability ®
-2	increase in stability.
Hydroxides of elements (+6)	–	H2SO4	H 2 SeO 4	H 6 TeO 6	–
		Strong acids		Weak acid
Hydroxides of elements (+4)	–	SO 2 * nH 2 O,	H 2 SeO 3	H 2 TeO 3,	PoO(OH)
	Weak acids				Amphoteric hydro-d
Hydrogen compounds	H2O	H 2 S	H 2 Se	H 2 Te	H 2 Po
	neutral	Weak acids
	increase in stability.

physiological action. All organic substances are oxygen compounds, so oxygen is a vital element for almost all living organisms (anaerobic bacteria are an exception). Oxygen enters the blood through the lungs. In the blood, oxygen weakly binds to hemoglobin (the chromophore of red blood cells) to form oxyhemoglobin and in this form is supplied to the cells. Under the action of enzymes, oxygen oxidizes the grape sugar (glucose) also brought in by the blood, turning it into carbon dioxide and water; the energy released in this case is used for various life processes (muscle work, body heating, etc.).

allotropic modifications. In its free form, oxygen forms two modifications: dioxygen (regular and oxygen) O 2 and trioxygen (ozone) O 3.

Dioxygen O 2

Structure. The structure of the О 2 molecule, which has two unpaired electrons, is correctly transmitted only within the framework of the molecular orbital method. The traditional image of an oxygen molecule with a double bond (O = O) does not convey the features of its structure and is therefore not entirely correct.

Receipt.

1. From air by fractional condensation and distillation (Linde way), the method is used in industry.

2. Heating of oxygen-containing substances, namely chlorates in the presence of a catalyst - pyrolusite MnO 2 (reaction 1), nitrates (reaction 2), permanganates at moderate or very high temperatures (reactions 3 and 4, respectively), peroxides (reaction 5):

2KS1O 3 \u003d 2KS1 + 3O 2 (1)

2KNO 3 \u003d 2KNO 2 + O 2 (2)

2KMnO 4 \u003d K 2 MnO 4 + MnO 2 + O 2 (3)

4KMnO 4 \u003d 2K 2 O + 4MnO 2 + ZO 2 (4)

2ВаО 2 \u003d 2ВаО + О 2 (5)

3. Catalytic decomposition of hydrogen peroxide (catalyst - pyrolusite - MnO 2):

2H 2 O 2 \u003d 2H 2 O + O 2

4. Electrolysis of alkaline or sulfate solutions using insoluble (platinum) anodes, on which hydroxide ions are discharged or water is oxidized:

4OH - - 4e\u003d O 2 + 2H 2 O; 2H 2 O - 4e = About 2 + 4H+

5. Interaction of peroxides of alkaline elements with carbon dioxide:

2Na 2 O 2 + 2CO 2 \u003d 2Na 2 CO 3 + O 2

This reaction is carried out in oxygen insulating devices.

Physical Properties. Colorless gas, odorless and tasteless. Moderately soluble in water, but somewhat better than nitrogen; in dissolved air, the oxygen content is 36% (vol.). Liquid and solid dioxygen has a light blue color.

Chemical properties. At room temperature, it is relatively little reactive; at high temperatures, due to the weakening of the oxygen-oxygen bond, the activity of O 2 increases.

The chemical addition of oxygen is called oxidation, it is slow and fast. Slow oxidation is, for example, the processes of rust formation on iron objects, the absorption of food by the body, the decay of organic residues, the aging of rubber, and the curing of oil paints. Rapid oxidation, often accompanied by the appearance of a flame, is called burning. In pure (as well as in liquid) oxygen, substances burn more intensely than in air, for example, a wood splinter smoldering in air ignites. When substances are oxidized with oxygen, oxides, for example: 2H 2 S + 3O 2 \u003d 2H 2 O + 2SO 2.

Detection. By the bright ignition of a smoldering torch (with an oxygen content of more than 30%); by brown staining of an alkaline solution of pyrogallol.

Application. Oxygen is stored and transported in steel cylinders under an overpressure of 150 atm. There must be no grease on the cylinder valve. Oxygen is used for welding and cutting metals and in breathing apparatus, as an oxidizer for rocket fuels and as a reagent in many chemical and technological processes. Oxygen-enriched air is used in various metallurgical methods, for gasification of brown coal under pressure, etc.

liquid air. Get by the way Linde, which is as follows. The air is compressed and the heat released during this is removed; subsequent expansion is followed by cooling. By repeating this operation with intermediate cooling, liquefied air is obtained at a temperature of about -190 °C. Liquid air has a light blue color. It is stored in Dewar vessels, which are forbidden to close with a tight stopper. The color intensity of liquid air increases during storage as the more volatile colorless nitrogen evaporates. Mixtures of liquid air with activated carbon, wood flour and other particulate materials are explosive.

Trioxygen (ozone) O 3

Receipt. Ozone is formed from ordinary oxygen (in pure form or in air) under the action of a glowing electric discharge or ultraviolet radiation (3O 2 - 2O 3). The oxygen produced at the anode by the electrolysis of dilute sulfuric acid using high current density contains significant amounts of ozone.

Properties. Light blue gas with a characteristic "electric" smell. Explodes when heated. A very strong oxidizing agent, but weaker than atomic oxygen. With silver, it forms black silver peroxide (the exact formula is unknown), the latter ignite on contact with ether or alcohol.

Application. Ozone is used to disinfect drinking water, in medicine as a disinfectant, and to neutralize industrial wastewater.

Atmospheric ozone layer. In the stratosphere (25 km above the Earth's surface), ozone is formed under the action of solar radiation, and although its amount is small (compared to atmospheric oxygen), ozone is sufficient to absorb ultraviolet radiation, which is dangerous for all living organisms. Thus, the ozone layer in the stratosphere ensures the normal development of organic life on Earth.

OXYGEN COMPOUNDS

oxides

Receipt.

1. The interaction of simple substances with oxygen (oxidation of elements in free form), for example, during their combustion in an oxygen atmosphere or in air.

2. Calcination of hydroxides or hydrated oxides: Cu (OH) 2 \u003d CuO + H 2 O.

3. Heating of salts that decompose with the formation of volatile acid oxides (carbonates, sulfates, sulfites, nitrates, etc.): СuСО 3 = СuО + СO 2.

Properties. The oxides of many non-metals, with the exception of CO, NO, N 2 O, correspond to acids. They are often obtained as a result of thermal decomposition of an acid or form it upon interaction with water (acid oxides): SO 3 + H 2 O \u003d H 2 SO 4.

Metal oxides in high oxidation states (+5) - (+7) are also acidic oxides. For example, chromium trioxide reacts with water to form chromic acid:

CrO 3 + H 2 O \u003d H 2 CrO 4

Metal oxides in low oxidation states from (+1) to (+4) are basic or amphoteric oxides, they correspond to basic or amphoteric hydroxides, for example:

CaO + H 2 O = Ca(OH) 2 ; A1 2 O 3 + 3H 2 O \u003d 2A1 (OH) 3.

Most metal oxides do not react with water under normal conditions, and therefore the hydroxides corresponding to them are obtained indirectly, for example, through salts:

CuO + 2HC1 \u003d CuC1 2 + H 2 O, CuC1 2 + 2NaOH \u003d Cu (OH) 2 + 2NaC1.

Basic oxides react with typical acid oxides and acids to form the corresponding salts; so do reactions between acidic oxides and typical basic oxides or bases. Amphoteric oxides form salts with both acidic and basic oxides.

Hydroxides

Hydroxides necessarily contain the -O-H group. Depending on whether the hydroxy group is bonded to metal or non-metal atoms, hydroxides will be basic, acidic, or amphoteric.

Most metal hydroxides are slightly soluble in water and precipitate when they are obtained from an aqueous solution: CuSO 4 + 2NaOH \u003d Cu (OH) 2 (t) + Na 2 SO 4.

Hydroxides usually precipitate at room temperature as slimy, flocculent, often colored precipitates with a higher water content than indicated by the stoichiometric formula, so they are assigned the oxide polyhydrate composition. The stoichiometric composition can be achieved by heating the polyhydrated oxide, but usually partially dehydrated hydroxide oxides of the Al(OH) or Tl(OH) 2 type are formed.

Coloring of sparingly soluble hydroxides:

white: A1 (OH) 3, AlO (OH), Zn (OH) 2, Cd (OH) 2, Pb (OH) 2, Sn (OH) 2, Bi (OH) 3, VU (OH), Mg (OH) 2;

light green: Fe(OH) 2 [this hydroxide turns brown in air;

light brown: Mn(OH) 2 ;

bright green: Ni(OH) 2 ;

grey-blue: Cr(OH) 3 ;

blue: Cu(OH) 2 ;

pink: Co(OH) 2 ;

Hydroxides of silver (I) and mercury (II) are very unstable and at room temperature spontaneously decompose into oxides and water.

Peroxides

Peroxides necessarily contain an oxygen chain -O-O- (peroxo group), they can be considered as derivatives of hydrogen peroxide H-O-O-H. The most important representatives are sodium peroxide Na 2 O 2 and barium peroxide BaO 2: they contain peroxide ions O *. =Pb=O) is lead(IV) oxide Organic peroxides are widely used as polymerization catalysts.

Metal superoxides contain a chain ion O 2 - ; for example, during the combustion of potassium, cadium superoxide KO 2 is formed.

SULFUR

The element sulfur S in the form of secretions from volcanic sources has been known since the 2nd century BC. BC e.

distribution in nature. Sulfur occurs in free form (native sulfur) and forms many minerals in the form of sulfides and sulfates. It is part of natural coal, oil and protein bodies (especially a lot of sulfur is found in the keratin of hair, feathers and wool).

Minerals: sulfides (pyrite - light with a metallic luster; luster - dark with a metallic sheen; blende - dark without a metallic luster or more often light, transparent), pyrite, sulfur pyrite, iron pyrite FeS 2, molybdenite, MoS 2 molybdenum luster, chalcopyrite, copper pyrite FeCuS 2, argentite, silver shine Ag 2 S, stibnite, antimony shine, gray antimony ore Sb 2 S 3, arsenopyrite, mispikel, arsenic pyrite FeAsS, sphalerite, zinc blende ZnS, cinnabar HgS, realgar As 4 S 4 galena, lead shine PbS, chalcocite, Cu 2 S copper sheen.

physiological action. Sulfur is a vital element, in a bound form it is found in all higher organisms (a component of proteins).

For people, free sulfur is not poisonous, small amounts of it act as a laxative, fine sulfur irritates the skin (the use of medicinal sulfur-containing ointments is based on this).

Receipt.

1. Smelting of native sulfur from natural deposits, for example, using steam, and purification of raw sulfur by distillation. With a sharp cooling of sulfur vapor, sublimated sulfur is obtained in the form of a fine powder (“sulfur color”).

2. The release of sulfur during the desulfurization of coal gasification products (water, air and light gases), for example, under the action of air and a catalyst - active coal:

2H 2 S + O 2 \u003d 2H 2 O + 2S.

3. The release of sulfur during the incomplete combustion of hydrogen sulfide (see the equation above), when acidifying a solution of sodium thiosulfate: Na 2 S 2 O 3 + 2HC1 \u003d 2NaС1 + SO 2 + H 2 O + S, and during the distillation of an ammonium polysulfide solution: (NH 4) 2 S 3 .

allotropic modifications. Sulfur in free form consists of molecules of different lengths (S ¥, S 12, S 8, S 6, S 2, etc.), and these molecules can be ordered in various ways, so there are several modifications of sulfur. At room temperature, sulfur is in the form of a-sulfur (rhombic modification), which is yellow brittle crystals, colorless and odorless, insoluble in water, but easily soluble in carbon disulfide. Above 96 °C, a-sulfur slowly transforms into b-sulfur (monoclinic modification), which is almost white crystalline plates. The melting points of a- and b-sulfur are 118 and 119°C, respectively. Upon melting, yellow low-viscosity l-sulfur is formed, which, like both modifications of solid sulfur, consists of cyclic S 8 molecules. Upon further heating, the S 8 cycles are reformed into chains of different lengths. A modification of such a structure is called m-gray; it is a red-brown and very viscous liquid. As the temperature rises, the color becomes dark brown and the viscosity of liquid sulfur decreases again. Liquid sulfur boils at 444.6 °C. When molten sulfur is poured into water, supercooling of the melt occurs and the formation of yellow-brown, rubbery, knife-cut plastic sulfur (a mixture of l- and m-sulfur), which in air becomes yellow, cloudy and brittle in a few minutes.

Chemical properties. When heated in air, sulfur burns with a blue flame to sulfur dioxide SO 2 (with an admixture of sulfur trioxide SO 3). At high temperatures, it reacts with metals, giving the corresponding sulfides, and with hydrogen (and paraffin), forming hydrogen sulfide H 2 S. Sulfur dissolves in an ammonium sulfide solution to form yellow-red polysulfide ions, when sulfur is heated with a sulfite solution, the corresponding thiosulfate is obtained, and when heated with a solution of cyanide - thiocyanate.

Application. Sulfur is used to produce carbon disulfide, sulfuric acid, sodium thiosulfate, sulfur dyes, ultramarine blue, in the vulcanization of rubber, as a treatment for skin diseases, and to protect plants from powdery mildew.

Sulfur is introduced into arable land in the form of various sulfate-containing fertilizers (ammonium sulfate, superphosphate).

HYDROGEN SULFIDE. SULFIDES

Hydrogen sulfide (monosulfan) H 2 S.

distribution in nature. Hydrogen sulfide is found in sulfuric mineral springs, volcanic and natural gas, large amounts of hydrogen sulfide are formed during the natural decay of protein substances.

physiological action. hydrogen sulfide very poisonous. Inhalation of air containing 0.08% (vol.) H 2 S for 5-10 minutes leads to death. Like hydrogen cyanide, hydrogen sulfide blocks vital respiratory enzymes (cytochromes). Laboratory work with hydrogen sulfide should be carried out only in a fume hood.

Detection. By black-brown staining of "lead paper" - impregnated with a solution of lead (II) salt and dried filter paper; by black coating (formation of Ag 2 S) on silver.

Receipt.

1. Industrial method - separation from water, household, coke oven and crude synthesis gas using solutions of sodium salts of amino acids that absorb H 2 S in the cold and release when heated or using deep-cooled methanol, which also absorbs H 2 S well.

2. Treatment of iron (II) sulfide with hydrochloric acid: FeS + 2HC1 = FeC1 2 + H 2 S.

3. Heating sulfur with paraffin.

4. Direct synthesis from hydrogen and sulfur (hydrogen is passed over molten sulfur).

The last three methods are used in laboratory conditions.

Properties. Colorless gas with the smell of rotten eggs, bp. -61 °С. It burns with a blue flame and upon complete combustion forms sulfur dioxide: 2H 2 S + 3O 2 \u003d 2H 2 O + 2SO 2.

When cold objects (for example, porcelain) are introduced into the flame, they are covered with a yellow coating of sulfur due to incomplete combustion, which corresponds to black soot during incomplete combustion of hydrocarbons (methane, acetylene).

Hydrogen sulfide is slightly soluble in water; when dissolved, the so-called sulfuric water, from which sulfur precipitates in air as a result of slow oxidation. Hydrogen sulfide is one of the weakest acids in aqueous solution.

Application. Hydrogen sulfide is used to produce sulfur and as a quantitative analysis reagent in inorganic chemistry.

Sulfides

Hydrogen sulfide salts are called sulfides. In a broader sense, these are compounds of electropositive elements with sulfur, which thus has a negative oxidation state (-2).

Heavy metal sulfides are industrially important ores; they are converted into oxides by firing in air: 2PbS + 3O 2 \u003d 2PbO + 2SO 2.

Sulfides of alkali and alkaline earth elements, as well as ammonium sulfide, are highly soluble in water. The remaining sulfides are released in the form of characteristically colored precipitates when an ammonium sulfide solution is introduced into metal salt solutions, and practically insoluble sulfides (having extremely low solubility in water) precipitate even from acidic salt solutions when hydrogen sulfide is introduced: FeSO 4 + (NH 4) 2 S = FeS (t) + (NH 4) 2 SO 4, 2ВiС1 3 + 3Н 2 S \u003d Bi 2 S 3 (t) + 6HC1.

Sulfides precipitated from acidic solutions by hydrogen sulfide:

black - HgS, Ag 2 S, PbS, CuS orange - Sb 2 S 3, Sb 2 S 5

brown - SnS, Bi 2 S 3 yellow - As 2 S 3, As 2 S 5, SnS 2, CdS

Sulfides precipitated from ammonia solutions under the action of ammonium sulfide (NH 4) 2 S: black - FeS, NiS, CoS, pink - MnS, white - ZnS.

OXYGEN COMPOUNDS OF SULFUR

Sulfur dioxide SO 2

distribution in nature. Sulfur dioxide is found in volcanic gases and waste gases released when natural coal is burned.

Receipt.

1. Burning sulfur or hydrogen sulfide.

2. Treatment of sulfites with strong acids: Na 2 SO 3 + 2HCl \u003d 2NaCl + H 2 O + SO 2.

3. Roasting sulfide ores, such as pyrite: 4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2

4. Reductive thermal decomposition of gypsum minerals CaSO 4 2H 2 O or anhydrite CaSO 4 .

The last two methods are used in industry.

Properties. A colorless, heavy gas with a pungent odor that causes coughing. Liquefies at -10°C. Non-flammable, very easily soluble in water. In solution, SO 2 is easily oxidized, for example, by potassium permanganate (quickly) or atmospheric oxygen (slowly), to sulfuric acid H 2 SO 4.

Sulfur dioxide acts as a bleaching agent on many dyes; in contrast to the irreversible action of bleaching powder, sulfur dioxide discoloration is often reversible and the color returns after washing.

Application. SO 2 is an intermediate product in the production of sulfuric acid and other sulfur compounds. It is used for bleaching paper, straw and wool, in the processing of wine barrels, for the sulfochlorination of saturated hydrocarbons. Liquid sulfur dioxide is used for oil refining.

Sulfites

When sulfur dioxide is dissolved in water, acidic polyhydrate SO 2 *nH 2 O is formed, which was previously represented by the conditional formula H 2 SO 3 (such molecules are unknown) and called sulfuric acid. The polyhydrate SO 2 *nH 2 O in aqueous solution is an acid of medium strength; when this solution is neutralized, sulfites are formed.

The general formula of medium sulfites M I 2 SO 3, acid sulfites (hydrosulfites) M I HSO 3.

Only sulfites of alkaline elements are soluble in water; when solutions of these sulfites are boiled with sulfur, they turn into the corresponding thiosulfates. All sulfites decompose under the action of strong acids with the release of SO 2.

The most important sulfites are sodium sulfite Na 2 SO 3 and sodium hydrosulfite NaHSO 3 . A solution of calcium hydrosulfite Ca(HSO 3) 2 , called "sulfite liquor", is obtained from calcium carbonate (limestone), sulfur dioxide and water, it serves as a means for extracting lignin from wood during pulp production.

Disulfites M I 2 S 2 O 6 - derivatives of an unknown free form of disulfurous acid H 2 S 2 O 6 fd; these salts (previously called pyrosulfites or metabisulfites) can be obtained by heating hydrosulfites: 2KHSO 3 \u003d K 2 S 2 O 5 + H 2 O.

Potassium disulfite K 2 S 2 O 5 is widely used in photographic developers and fixers.

Sulfur trioxide SO 3

Receipt. Catalytic oxidation of sulfur dioxide, distillation from oleum, thermal decomposition of K 2 S 2 O 7 into K 2 SO 4 and SO 3 (laboratory method).

Properties. Three modifications of SO 3 are known. The most stable - a-SO 3 is formed in the form of silky-shiny needles, which smell strongly in the air, so pl. 40°C. They react vigorously with water to give sulfuric acid. Ice-like modification - g-SO 3 has so pl. 16.8°C and bp 44.8 °C.

7. Sulfuric acid H 2 SO 4

Receipt. Separation of sulfuric acid from sulfates with a strong acid followed by evaporation of H 2 SO 4 is impossible, since sulfuric acid itself is strong and decomposes above 300 ° C. All industrial methods for its synthesis are based on the production of sulfur dioxide SO 2, its oxidation to sulfur trioxide SO 3 and the interaction of the latter with water.

The first stage in the production of sulfuric acid - the production of sulfur dioxide - can be carried out in three ways:

The most common is the roasting of sulfide ores, such as pyrite. The process is carried out in tubular rotary or multi-hearth furnaces, as well as in fluidized bed furnaces. Technological processes of non-ferrous metallurgy are always accompanied by the production of H 2 SO 4, since sulfur dioxide is formed during the roasting of sulfide ores.

The second stage in the production of sulfuric acid is the oxidation of sulfur dioxide, this process is carried out by the contact or nitrous method.

Approximately 80% of the world's production of sulfuric acid is carried out by the contact method. The method has been known since 1900. The product is concentrated H 2 SO 4 .

physical properties. Colorless, odorless oily liquid, density 1.84 g/cm3 at 20°C. At 338 °C, it boils, forming a mist of SO 3 .

When diluted with water, strong heating occurs (formation of hydrates, for example, H 2 SO 4 *H 2 O), which is accompanied by splashing of the liquid.

The rule for diluting sulfuric acid is to pour acid into water while stirring, and not vice versa. Sulfuric acid is very hygroscopic and therefore suitable for drying many gases (but not ammonia!).

Chemical properties. A very strong dibasic acid, already at moderate dilution, it almost completely dissociates into H + ions (more precisely, H 3 O +) and SO 4 2-:

H 2 S0 4 + 2H 2 0 \u003d SO 4 2- + 2H 3 O +.

Hydrosulfate ions HSO 4 - exist only in concentrated solutions of H 2 SO 4:

H 2 SO 4 + H 2 O \u003d HSO 4 - + H 3 O +.

Sulfuric acid is low volatile and displaces many other acids from their salts, for example:

CaF 2 + H 2 SO 4 \u003d CaSO 4 + 2HF.

Diluted H 2 SO 4, when interacting with base metals (standing in the electrochemical series of voltages to the left of hydrogen), releases hydrogen.

Concentrated H 2 SO 4 never emits hydrogen (formally, even because it does not contain at all or contains few H 3 O + ions), it reacts as an oxidizing agent and most often passes into SO 2, and when interacting with strong reducing agents, into S and H 2 S. When H 2 SO 4 (conc.) is heated, it oxidizes almost all metals, including the noble metals Cu, Hg and ag: Cu + 2H 2 SO 4 (conc.) \u003d CuSO 4 + 3O 2 + 2H 2 O.

Detection.

1. Concentrated sulfuric acid is conveniently identified by the charring of a splinter immersed in it.

2. Sulfate ions SO 4 2- form with Ba 2+ ions a white finely crystalline precipitate of barium sulfate BaSO 4 .

Application. Sulfuric acid belongs to the main chemical production products. It is used in the production of chemical fibers (viscose silks, wool, polyamide fibers), fertilizers (superphosphate), explosives, detergents, wetting and emulsifying agents, dyes, medicines, as well as various sulfates, ethers and esters, some acids (hydrofluoric acid, tartaric acid, etc.), for refining mineral oils, for pickling metals, as a component of various galvanic electrolytes (for chromium plating, anodic oxidation, etc.), as an electrolyte for lead batteries, and for many other purposes.

Oleum

Fuming sulfuric acid contains an excess of sulfur trioxide, in particular in the form of disulfuric acid H 2 S 2 O 7 . Such a liquid mixture of H 2 SO 4 , H 2 S 2 O 7 and excess SO 3 is called oleum. The composition of the oleum is indicated by the percentage of SO 3 (in excess of the monohydrate SO 3 * H 2 O, i.e. 100% H 2 SO 4).

sulfates- salts of sulfuric acid.

Sulfates of lead (II), calcium, strontium and barium are very slightly soluble in water, most other sulfates are easily soluble in water. The method for detecting them is similar to the method for detecting SO 4 2-sulfuric acid ions. Many sulfates are found in the earth's crust in the form of minerals.

The most important natural sulfates: mirabilite (Glauber's salt) - Na 2 SO 4 * 10H 2 O, epsomite (bitter, or English salt) MgSO 4 * 7H 2 O.

Vitriols are crystalline hydrates of sulfates of some divalent metals:

ferrous sulfate (light green) FeSO 4 * 7H 2 O; copper sulfate (blue) CuSO 4 * 5H 2 O; nickel vitriol (green) NiSO 4 * 7H 2 O; cobalt vitriol (dark red) CoSO 4 * 7H 2 O zinc sulfate (white) ZnSO 4 -7H 2 O.

Alum is a crystalline hydrate of double sulfates:

potassium alum K 2 SO 4 * A1 2 (SO 4) 3 * 24H 2 O;

potassium chromium alum K 2 SO 4 *Cr 2 (SO 4) 3 * 24H 2 O;

iron-potassium alum K 2 SO 4 * Fe 2 (SO 4) 3 * 24H2O.

Mohr's salt is not alum, its composition is (NH 4) 2 SO 4 * FeSO 4 * 6H 2 O.

OTHER SULFUR COMPOUNDS

Disulfates are salts of disulfuric acid H 2 S 2 O 7.

Thiosulfuric acid H 2 S 2 O 3 is stable only at low temperatures (below -72°C). Its thiosulfate salts are formed by boiling solutions of metal sulfites with excess sulfur:

Na 2 S 2 O 3 + S \u003d H 2 S 2 O 3.

It is not possible to obtain acid H 2 S 2 O 3 by displacing it from thiosulfates with a strong acid, since it decomposes: Na 2 S 2 O 3 + 2HC1 \u003d 2NaC1 + SO 2 + S + H 2 O.

peroxodisulfuric acid H 2 S 2 O 8, or more accurately H 2 S 2 O 6 (O 2), contains a peroxo group - O - O -, in a free form it is very unstable. Its salts - peroxodisulfates - are very strong oxidizing agents, for example, potassium peroxodisulfate K 2 S 2 O 8. Peroxomonosulfuric acid is also known (caro acid) H 2 SO 3 (O 2).

Dithionic acid H 2 S 2 O 4 is not known in free form, but its sodium dithionite Na 2 S 2 O 4 salt has been obtained, which is used as a reducing agent, for example, in the synthesis of vat dyes, in etch printing and in bleaching processes. Sodium dithionite is produced by passing sulfur dioxide into an aqueous suspension of zinc: Zn + 2SO 2 \u003d Zn 2+ + S 2 O 4, followed by removal of Zn 2+ ions from the solution by adding sodium carbonate and crystallization of Na 2 S 2 O 4. The S 2 O 4 2- ion contains a direct sulfur-sulfur bond.

Dithionic acid H 2 S 2 O 6, its dithionate salts, and tetrathionic acid H 2 S 4 O 6 , its tetrathionate salts, exist only in dilute aqueous solution. They contain two and four sulfur atoms linked directly in a chain. Manganese dithionate (P) is formed by treating manganese dioxide (pyrolusite) with sulfur dioxide: MnO 2 + 2SO 2 = MnS 2 O 6.

Sodium tetrathionate is obtained by reacting sodium thiosulfate with iodine:

2Na 2 S 2 O 3 + I 2 = Na 2 S 4 O 6 + 2NaI.

Other oxygenated sulfur acids are sulfoxylic acid H 2 SO 2, thiosulfur acid H 2 S 2 O 2, tri-, penta- and hexathionic acids H 2 S 3 O 6, H 2 S 5 O 6 and H 2 S 6 O 6, their salts tri-, penta- and hexathionates.

Disulfur dichloride S 2 C1 2 - orange-yellow, sometimes colorless, liquid fuming in humid air with a characteristic suffocating odor. It is formed when sulfur is heated with a lack of chlorine. Used in the vulcanization of rubber.

Sulfur hexafluoride SF 6 is a colorless, odorless gas. Chemically inert. In engineering, it is used as a gas electrical insulator.

Sulfuryl chloride SCl 2 O 2 and thionyl chloride SC1 2 O - colorless liquids that form a mist in the air and cause a strong cough. They are completely hydrolyzed by water:

SC1 2 O 2 + 2H 2 O \u003d H 2 SO 4 + 2HC1; SC1 2 O + H 2 O \u003d SO 2 + 2HC1.

Chlorosulfonic acid HSO 3 C1 is also known.

SELENIUM, TELLURIUM, POLONIUM AND THEIR COMPOUNDS

Opening. Selenium Se was discovered in 1817 in the sludge of lead chambers (production towers) of a sulfuric acid plant (Bertzelius, Sweden).

Spreading. Selenium is a rare element; does not have its own minerals. It is found in small amounts (together with tellurium) in native sulfur and sulfide ores.

Receipt. Isolation from anode sludge of copper electrolytic installations. For this purpose, the sludge is treated with a solution of sodium hydroxide and sulfur dioxide:

2SeO 2 + 2SO 2 + 2OH - \u003d 2SO 4 2- + Se + H 2 O.

Solid selenium is separated and purified by distillation.

Properties. Selenium has two allotropic modifications.

Gray (metallic) selenium is a gray substance with a faint luster. Does not dissolve in carbon disulfide. The electrical resistance of this modification sharply (by a factor of ≈1000) decreases in the light (compared to the electrical resistance in the dark). sustainable modification.

Red selenium is a red non-metallic substance. Soluble in CS 2 to form a yellow solution. Thermodynamically unstable modification.

Selenium burns in air with a blue flame, spreading the characteristic smell of rotten radish. As a result, white solid selenium dioxide SeO 2 is formed. Gray selenium turns into red selenium when dissolved in hot concentrated sulfuric acid and pouring the resulting green solution into a large volume of water.

Application. Selenium is used in the production of photovoltaic cells and electric current rectifiers.

Selenium compounds. The properties of selenium compounds are similar to those of sulfur compounds. The best known hydrogen selenide H 2 Se (derivatives - selenides); selenium dioxide SeO 2 is a white solid, with water forms selenious acid H 2 SeO 3 (salts - selenites); selenic acid H 2 SeO 4 , equal in strength to sulfuric acid; its salts are selenates, of which barium selenate ВаSeO 4 is very slightly soluble in water.

Tellurium Those, as a rule, accompanies selenium and sulfur in natural sulfides, a rather rare element. Discovered in 1782 in gold-bearing rocks (Müller von Reichenstein, Hungary). It is a silvery-white soft but brittle metal. Used in semiconductor technology. Hydrogen telluride H 2 Te (derivatives - tellurides) exhibits stronger acidic properties than hydrogen selenide, but is much more stable with respect to atmospheric oxygen.

Polonium Rho was discovered in 1898 in uranium resin ore (M. Sklodowska-Curie and P. Curie, France). A very rare radioactive element. Obtained artificially by irradiation of bismuth in nuclear reactors; the longest-lived isotope is polonium-209 (half-life 102 years). It is a silvery-white lustrous metal that glows with a constant blue luminescence. In all compounds, polonium behaves like a typical metal.

§8 Elements VI And the groups.

Oxygen, sulfur, selenium, tellurium, polonium.

General information of elements VI A group:

Group VI A elements (except polonium) are called chalcogenides. There are six valence electrons (ns2 np4) on the outer electronic level of these elements, therefore they show valency 2 in the normal state, and -4 or 6 in the excited state (except for oxygen). The oxygen atom differs from the atoms of other elements of the subgroup by the absence of a d-sublevel in the outer electron layer, which causes high energy costs for the “pairing” of its electrons, which are not compensated by the energy of the formation of new covalent bonds. Therefore, the covalence of oxygen is two. However, in some cases, the oxygen atom, which has unshared electron pairs, can act as an electron donor and form additional covalent bonds according to the donor-acceptor mechanism.

The electronegativity of these elements gradually decreases in the order O-S-Se-Te-Rho. The degree of oxidation is from -2, +2, +4, +6. The radius of the atom increases, which weakens the non-metallic properties of the elements.

The elements of this subgroup form compounds of the form H2 R with hydrogen (H2 O, H2 S, H2 Se, H2 Te, H2 Po). These compounds, dissolving in water, form acids. Acid properties increase in the direction H2 O→H2 S→H2 Se→H2 Te→H2 Po. S, Se and Te form compounds of the RO2 and RO3 type with oxygen. From these oxides, acids of the type H2 RO3 and H2 RO4 are formed. As the atomic number increases, the strength of the acids decreases. All of them have oxidizing properties. Acids like H2 RO3 also exhibit reducing properties.

Oxygen

Natural compounds and preparations: Oxygen is the most abundant element in the earth's crust. In a free state, it is found in atmospheric air (21%); in a bound form, it is part of water (88.9%), minerals, rocks and all substances from which plant and animal organisms are built. Atmospheric air is a mixture of many gases, the main part of which is nitrogen and oxygen, and a small amount of noble gases, carbon dioxide and water vapor. Carbon dioxide is formed in nature during the combustion of wood, coal and other fuels, the respiration of animals, and decay. In some parts of the world, CO2 is released into the air through volcanic activity and also from underground sources.

Natural oxygen consists of three stable isotopes: 816 O (99.75%), 817 O (0.04), 818 O (0.20). The isotopes 814 O, 815 O, 819 O were also obtained artificially.

Oxygen was first obtained in pure form by K.W. Scheele in 1772, and then in 1774 by D.Yu. Priestley, who isolated it from HgO. However, Priestley did not know that the gas he received was part of the air. Only a few years later, Lavoisier, who studied the properties of this gas in detail, established that it is the main part of the air.

In the laboratory, oxygen is obtained by the following methods:

E water electrolysis. To increase the electrical conductivity of water, an alkali solution (usually 30% KOH) or alkali metal sulfates is added to it:

In general: 2H2 O → 2H2 + O2

At the cathode: 4H2 O+4e¯ → 2H2 +4OH¯

At the anode: 4OH−4е→2H2 О+О2

- Decomposition of oxygen-containing compounds:

Thermal decomposition of Bertolet's salt under the action of MnO2 catalyst.

KClO3 →2KCl+3O2

Thermal decomposition of potassium permanganate

KMnO4 → K2 MnO4 + MnO2 + O2.

Thermal decomposition of alkali metal nitrates:

2KNO3 →2KNO2 +О2.

Decomposition of peroxides:

2H2 O2 → 2H2 O + O2.

2ВаО2 → 2ВаО+О2.

Thermal decomposition of mercury oxide (II):

2HgO→2HgO+О2.

The interaction of peroxides of alkali metals with carbon monoxide (IV):

2Na2 O2 + 2CO2 → 2Na2 CO3 + O2.

Thermal decomposition of bleach in the presence of a catalyst - cobalt salts:

2Ca(OCl)Cl → 2CaCl2 + O2.

Oxidation of hydrogen peroxide with potassium permanganate in an acidic medium:

2KMnO4 + H2 SO4 + 5H2 O2 → K2 SO4 + 2Mn SO4 + 8H2 O + 5O2.

In industry: At present, oxygen is produced in industry by fractional distillation of liquid air. With weak heating of liquid air, nitrogen is first separated from it (tboil (N2) = -196ºC), then oxygen is released (tboil (O2) = -183ºС).

The oxygen obtained by this method contains nitrogen impurities. Therefore, to obtain pure oxygen, the resulting mixture is re-distilled and ultimately 99.5% oxygen is obtained. In addition, some oxygen is obtained by electrolysis of water. The electrolyte is a 30% KOH solution.

Oxygen is usually stored in blue cylinders at a pressure of 15 MPa.

Physiochemical properties: Oxygen is a colorless, odorless, tasteless gas, slightly heavier than air, slightly soluble in water. Oxygen at a pressure of 0.1 MPa and a temperature of -183ºС passes into a liquid state, at -219ºС it freezes. In the liquid and solid state, it is attracted by a magnet.

According to the method of valence bonds, the structure of the oxygen molecule, represented by the scheme -:Ö::Ö: , does not explain the great strength of a molecule that has paramagnetic properties, that is, unpaired electrons in the normal state.

As a result of the bonding of the electrons of two atoms, one common electron pair is formed, after which the unpaired electron in each atom forms a mutual bond with an unshared pair of another atom, and a three-electron bond is formed between them. In an excited state, the oxygen molecule exhibits diamagnetic properties, which correspond to the structure according to the scheme: Ö=Ö: ,

Two electrons are missing to fill the electron level in the oxygen atom. Therefore, oxygen in chemical reactions can easily add two electrons and exhibit an oxidation state of -2. Oxygen only in compounds with a more electronegative element fluorine exhibits the oxidation state +1 and +2: O2 F2, OF2.

Oxygen is a strong oxidizing agent. It does not interact only with heavy inert gases (Kr, Xe, He, Rn), with gold and platinum. The oxides of these elements are formed in other ways. Oxygen is included in the reactions of combustion, oxidation, both with simple substances and with complex ones. When non-metals interact with oxygen, acid or salt-forming oxides are formed, and when metals interact, amphoteric or mixed oxides are formed. Thus, oxygen reacts with phosphorus at a temperature of ~ 60 ° C,

4P+5O2 → 2P2 O5

With metals - oxides of the corresponding metals

4Al + 3O2 → 2Al2O3

3Fe + 2O2 → Fe3O4

when alkali metals are heated in dry air, only lithium forms oxide Li2 O, and the rest are peroxides and superoxides:

2Na+O2 →Na2 O2 K+O2 →KO2

Oxygen interacts with hydrogen at 300 °C:

2H2 + O2 = 2H2 O.

When interacting with fluorine, it exhibits reducing properties:

O2 + F2 = F2 O2 (in electrical discharge),

with sulfur - at a temperature of about 250 ° C:

Oxygen reacts with graphite at 700 °C

C + O2 = CO2.

The interaction of oxygen with nitrogen begins only at 1200°C or in an electric discharge:

N2 + O22NO - Q.

Oxygen also reacts with many complex compounds, for example, with nitric oxide (II), it reacts even at room temperature:

2NO + O2 = 2NO2.

During the oxidation of hydrogen sulfide, when heated, sulfur is formed, or sulfur oxide (IV), depending on the ratio between oxygen and hydrogen sulfide:

2H2 S + O2 = 2S + 2H2 O

2Н2 S + ЗО2 = 2SO2 + 2Н2 О

In most oxidation reactions involving oxygen, heat and light are released - such processes are called combustion.

Ozone

Ozone-O3 is the second allotropic modification of the element oxygen. The O3 molecule has an angular structure (the angle between the bonds is 116º, the length of the O=O bond, l=0.1278 nm) it is a blue gas. Liquid ozone is dark blue. It is poisonous and explosive especially in liquid and solid state). Ozone is formed in the atmosphere during lightning discharges, and has a specific smell of freshness.

Usually, ozone is produced in ozonizers by passing a quiet electrical discharge through oxygen (the reaction is endothermic and highly reversible; the ozone yield is 5%):

3О22О3 ΔН=-285 kJ. Under laboratory conditions, ozone is obtained by acidifying persulfate with nitric acid.

(NH4)2 S2 O8 →H2 S2 O8 +2NH4+

H2 S2 O8 →2SO2 +O3 +H2O

O3 is formed in low yield as a result of the reaction:

3F2 +H2 O(g)→6HF+O3

O3 is the strongest oxidizing agent, oxidizes all metals (except gold and platinum metals) and most non-metals. It converts lower oxides into higher ones, and metal sulfides into their sulfates. In reactions involving O3, O2 is usually formed, for example:

2Ag+O3 →Ag2 O+O2

PbS+4O3 →PbSO4 +4O2

NH2 +3O3 →HNO2 +H2O

Pb(OH)2 +O3 →PbO2 +H2O+O2

When exposed to O3 on alkali metals, ozonides can be obtained - unstable compounds that decompose:

2KO3 →2KO2 +O2

As a strong oxidizing agent, ozone kills bacteria and is therefore used to disinfect the air. A stable ozone layer is located in the atmosphere at a height of ~22 km. This ozone layer protects the Earth from life-damaging pure ultraviolet radiation.

When ozone interacts with a solution of potassium iodide, iodine is released, while this reaction does not occur with oxygen:

2KI + O3 + H2 O \u003d I2 + 2KOH + O2.

The reaction is often used as a qualitative one for the detection of I- or ozone ions. To do this, starch is added to the solution, which gives a characteristic blue complex with released iodine, and it is also of high quality because ozone does not oxidize Cl - and Br- ions.

Water

Physical and chemical properties of water: Pure water is a colorless, tasteless, odorless, transparent liquid. Density of water at the transition her from a solid to a liquid state does not decrease, as with almost all other substances, but increases.

Water is a familiar and unusual substance. There is no substance on earth that is more important to us than ordinary water, and at the same time there is no other substance whose properties would have as many contradictions and anomalies as its properties.

Almost ¾ of the surface of our planet is occupied by oceans and seas. Solid water - snow and ice - covers 20% of the land. The planet's climate depends on water. Geophysicists say that the Earth would have cooled down long ago and turned into a lifeless piece of stone, if not for water. She has a very high heat capacity. As it heats up, it absorbs heat, and as it cools, it releases it. Terrestrial water both absorbs and returns a lot of heat, thereby equalizing the climate. The Earth is protected from cosmic cold by those molecules that are scattered in the atmosphere - in clouds and in the form of vapors.

Water in physical properties differs significantly from other solvents: At 4ºС, water has a maximum density, and only with further heating does its density decrease. If, with a decrease in temperature and during the transition from a liquid to a solid state, water changed similarly to other substances, then when winter approached, the surface layers of natural waters would cool to 0 ° C and sink to the bottom until the entire mass of the reservoir would acquire a temperature of 0 ° C. The water would freeze, the ice floes would sink to the bottom, and the pond would freeze to its full depth. Many forms of life in water would be impossible. In reality, the cooled layer, which has a lower density, remains on the surface, freezes, and thus protects the underlying layers from cooling.

Water has an abnormally high heat capacity (4.18 J/g∙K), so at night, as well as during the transition from summer to winter, the water cools down slowly. And during the day, or during the transition from winter to summer, it also heats up slowly, thus being the temperature regulator on the globe.

Water in its normal state is a liquid, while H2 S, H2 Se, H2 Te are gases. The temperatures of crystallization and evaporation of water are significantly higher than the corresponding temperatures of these compounds.

Water has a very high dielectric constant (78.5 at 298K).

Water is a good solvent for polar liquids and compounds with ionic bonds; it forms crystalline hydrates with many chemical compounds.

For a long time, the unusual properties of water were a mystery to scientists. They are mainly due to the following reasons:

The polar nature of the molecules;

The presence of unshared electron pairs at the oxygen atom;

Hydrogen bonds.

The bond between hydrogen and oxygen atoms is polar, which leads to asymmetry in the distribution of electronic charges and, consequently, to the polarity of the molecule. The bond length is 96 nm, and the angle between bonds is ~ 105º.

The presence of lone pairs of electrons in oxygen and the shift of shared electron pairs from hydrogen atoms to oxygen cause the formation of hydrogen bonds. The binding energy is 25 kJ/mol. The oxygen atom in the water molecule is in a state of sp3 hybridization. Therefore, the HOH bond angle is close to the tetrahedral angle (109.5º).

The molecular weight of vaporous water is 18 and corresponds to its simplest formula. However, the molecular weight of the liquid is higher. This indicates that the association of molecules occurs in the liquid phase; their combination into more complex aggregates, due to the formation of hydrogen bonds between molecules.

In solid water (ice), the oxygen atom of each molecule is involved in the formation of two hydrogen bonds with neighboring water molecules.

The structure of ice belongs to the least dense structures; there are voids in it, the dimensions of which are somewhat larger than the dimensions of a water molecule. When ice melts, its structure is destroyed, but hydrogen bonds remain in the liquid phase, associates are formed, but they exist for a short time: the destruction of some and the formation of other aggregates is constantly occurring. In the voids of such "ice" aggregates, single water molecules can be placed, while the packing of water molecules becomes dense. That is why when ice melts, the volume occupied by water decreases, and its density increases. When water is heated, part of the heat is spent on breaking hydrogen bonds. This explains the high heat capacity of water. Hydrogen bonds between water molecules are completely broken only when water passes into steam.

On Earth, there is one deuterium atom for every 6800 protium atoms, and in interstellar space one deuterium atom is already for 200 protium atoms.

Water is a highly reactive substance.

Water reacts with many metals with hydrogen evolution:

2Na + 2H2 O = H2 + 2NaOH (violently)

2K + 2H2O = H2 + 2KOH (violently)

3Fe + 4H2 O = 4H2 + Fe3 O4 (only when heated)

Not all, but only sufficiently active metals can participate in redox reactions of this type. Alkali and alkaline earth metals react most easily.

From non-metals for example, carbon and its hydrogen compound (methane) react with water. These substances are much less active than metals, but still able to react with water at high temperatures:

C + H2 O ® H2 + CO

CH4 + 2H2 O ® 4H2 + CO2

Water decomposes into hydrogen and oxygen under the action of an electric current. It is also a redox reaction, where water is both an oxidizing agent and a reducing agent:

2H2O 2H2+O2

Water reacts with many oxides non-metals. Unlike the previous ones, these reactions are not redox, but compound reactions:

P2 O5 +3H2 O→2H3 PO4 ; N2 O5 +H2 O→2HNO3

Alkali and alkaline earth metal oxides react with water to form the corresponding alkalis:

CaO+H2O→Ca(OH)2

Not all metal oxides are capable of reacting with water. Some of them are practically insoluble in water and therefore do not react with water. These are ZnO, TiO2, Cr2 O3, from which, for example, water-resistant paints are prepared. Iron oxides are also insoluble in water and do not react with it. Many compounds of metals with non-metals easily interact with water to form the corresponding metal hydroxides and hydrogen compounds of non-metals:

PCl3 +3H2O → H3PO3 + 3HCl

Al2 S3 +6H2 O→2Al(OH)3 +3H2 S

Ca3 P2+6H2 O→3Ca(OH)2 +2PH3

Na3 N+3H2 O→3NaOH+NH3

KH+H2O→KOH+H2

Water forms numerous compounds in which its molecule is completely preserved. These are the so-called hydrates. If the hydrate is crystalline, then it is called crystalline hydrate, for example:

CuSO4 +5 H2O→CuSO4 . 5H2O

H2 SO4 + H2 O = H2 SO4 . H2O (sulfuric acid hydrate)

NaOH + H2O = NaOH . H2 O (caustic soda hydrate)

Compounds that bind water into hydrates and crystalline hydrates are used as desiccants. With their help, for example, remove water vapor from moist atmospheric air.

A special reaction of water - photosynthesis - the synthesis of starch (C6 H10 O5) n and other similar compounds (carbohydrates) by plants, occurring with the release of oxygen:

6n CO2 + 5n H2 O = (C6 H10 O5)n + 6n O2 (under the action of light)

Water has catalytic activity. In the absence of traces of moisture, ordinary reactions practically do not occur, for example, sodium, white phosphorus do not oxidize, chlorine does not interact with metals, hydrogen fluoride does not cut glass.

Hydrogen peroxide

Hydrogen peroxide H2 O2 is a hydrogen-oxygen compound containing a record amount of oxygen - 94% by mass. H2O2 molecules contain peroxide groups –О–О–, which largely determine the properties of this compound.

Due to the asymmetric distribution of H-O bonds, the H2O2 molecule is highly polar. A fairly strong hydrogen bond arises between H2O2 molecules, leading to their association. Therefore, under normal conditions, hydrogen peroxide is a pale blue syrupy liquid (density 1.44) with a rather high boiling point (150ºС). When storing H2 O2 decomposes.

Selenium is obtained from waste products of sulfuric acid, pulp and paper production and anode sludge from the electrolytic refining of copper. Selenium is present in sludge along with sulfur, tellurium, heavy and noble metals. To extract selenium, the sludge is filtered and subjected to either oxidative roasting (about 700 °C) or heating with concentrated sulfuric acid. The resulting volatile SeO2 is captured in scrubbers and electrostatic precipitators. From solutions, commercial selenium is precipitated with sulfur dioxide. Sintering of the sludge with soda is also used, followed by leaching of sodium selenate with water and isolation of selenium from the solution. To obtain high-purity selenium used as a semiconductor material, crude selenium is refined by vacuum distillation, recrystallization, and others.

Physical and chemical properties of selenium. The configuration of the outer electron shell of the Se 4s2 4p4 atom; the spins of two p-electrons are paired, while the other two are not paired, so selenium atoms are able to form Se2 molecules or chains of Sen atoms. Chains of selenium atoms can be closed into ring Se8 molecules. The diversity of the molecular structure determines the existence of selenium in various allotropic modifications: amorphous (powdered, colloidal, glassy) and crystalline (monoclinic α- and β-forms and hexagonal γ-forms). Amorphous (red) powdered and colloidal selenium (density 4.25 g / cm3 at 25 ° C) is obtained by reduction from a solution of selenious acid H2 SeO3, rapid cooling of selenium vapor and other methods. Vitreous (black) selenium (density 4.28 g/cm3 at 25°C) is obtained by heating any modification of selenium above 220°C followed by rapid cooling. Vitreous selenium has a vitreous luster and is brittle. Thermodynamically the most stable is hexagonal (gray) selenium. It is obtained from other forms of selenium by heating to melting with slow cooling to 180-210 ° C and holding at this temperature. Its lattice is built from parallel helical chains of atoms. The atoms within the chains are covalently bonded. All modifications of selenium have photoelectric properties. Hexagonal selenium up to the melting temperature is an impurity semiconductor with hole conductivity. Selenium is a diamagnet (its pairs are paramagnetic).

Selenium is stable in air; oxygen, water, hydrochloric and dilute sulfuric acids do not affect it, it is highly soluble in concentrated nitric acid and aqua regia, it dissolves disproportionately in alkalis:

Se + 4HNO3 → H2 SeO3 + 4NO2 + H2O

3Se + 6KOH → K2SeO3 + 2K2Se + 3H2O

Selenium in compounds has oxidation states -2, +2, +4, +6. With oxygen, selenium forms a number of oxides: SeO, Se2 O3, SeO2, SeO3. The last two are anhydrides of selenous H2 SeO3 and selenic H2 SeO4 acids (salts - selenites and selenates). SeO2 is the most stable. SeO2 and H2 SeO3 with strong oxidizing agents exhibit reducing properties:

3H2 SeO3 + HClO3 → 3H2 SeO4 + HCl

With halogens, selenium gives compounds SeF6, SeF4, SeCl4, SeBr4, Se2 Cl2 and others. Sulfur and tellurium form a continuous series of solid solutions with selenium. With nitrogen, selenium gives Se4 N4, with carbon - CSe2. Compounds with phosphorus P2 Se3, P4 Se3, P2 Se5 are known. Hydrogen interacts with selenium at t>=200 °C, forming H2 Se; H2Se solution in water is called hydroselenic acid. When interacting with metals, selenium forms selenides. Numerous complex compounds of selenium have been obtained. All selenium compounds are poisonous.

Application of selenium . Due to its cheapness and reliability, selenium is used in converter technology in rectifier semiconductor diodes, as well as for photoelectric devices (hexagonal), electrophotographic copiers (amorphous selenium), synthesis of various selenides, as phosphors in television, optical and signal devices, thermistors, etc. n. selenium is widely used to bleach green glass and obtain ruby ​​glasses; in metallurgy - to give cast steel a fine-grained structure, improve the mechanical properties of stainless steels; in the chemical industry - as a catalyst; selenium is also used in the pharmaceutical industry and other industries.

8.4 Tellurium

Natural compounds and obtaining. Basic. sources of tellurium are sludge from electrolytic refining of copper and sludge from sulfuric acid production, as well as alkaline dross from lead refining. During the processing of sulfuric acid sludge by the roasting method (see Selenium), tellurium remains in the cinder, which is leached with hydrochloric acid. Se is precipitated from the hydrochloric acid solution by passing SO2, after which the solution is diluted to an acid content of 10-12% and when heated by the action of SO2, tellurium is precipitated.

During sintering of sludge with soda and subsequent leaching, tellurium passes into a solution and, upon neutralization, precipitates in the form of TeO2. Tellurium is obtained either by direct reduction of TeO2 with coal, or by precipitation by the action of SO2 on hydrochloric acid solutions of TeO2. During the processing of sludge by the sulfide method (leaching with a Na2 S solution), tellurium is isolated from the solution (after Se precipitation by aeration) by the action of dry Na2 S2 O3:

Na2 TeS3 + 2Na2 SO3 → Te + 2Na2 S2 O3 + Na2 S

During the processing of copper electrolyte sludge, tellurium is mainly converted into soda slag, resulting from the remelting of residues into a gold-silver alloy (“Dore metal”). When sulfatization is used, part of the tellurium passes into sulfate solutions together with Cu. Of these, tellurium is precipitated by the action of metallic copper:

H2 TeO3 + 4H2 SO4 + 6Cu → Te + Cu2 Te + 4CuSO4 + 6H2 O

Tellurium is extracted from soda slags after dissolution in water or by neutralization with precipitation of TeO2 (it is purified by reprecipitation from sulfide or acid solutions, dissolved in alkali and tellurium is isolated by electrolysis), or rough tellurium is precipitated directly from a soda solution by electrolysis. It is reduced by A1 in an alkaline solution:

6Te + 2A1 + SNaOH → 3Na2 Te2 + 2NaAlO2 + 4H2 O. Then tellurium is precipitated by aeration:

2Na2 Te2 + 2H2 O + O2 → 4Te + 4NaOH

To obtain tellurium of high purity, its volatile compounds are used, in particular TeCl4, which is purified by distillation or rectification and extraction from hydrochloric acid solution. After hydrolysis of TeO2 chloride, H2 is reduced. Sometimes H2Te is also used for purification. At the final stages of purification, vacuum sublimation, distillation or rectification of tellurium, as well as zone melting or directional crystallization are used.

Physical and chemical properties. Tellurium is a silvery-gray substance with a metallic luster, in thin layers reddish-brown in light, golden yellow in pairs. Tellurium melt above ~ 700 °C has metallic conductivity. Tellurium is diamagnetic, magnetic. susceptibility - 0.31 10-9. Mohs hardness 2.3, Brinell 180-270 MPa; tear resistance 10.8 MPa. Tellurium is brittle and becomes ductile when heated.

For tellurium, the normal electrode potential is 0.56 V. Tellurium, even dispersed, is stable in air, but when heated, it burns (a blue flame with a green halo) to form TeO2. Crystalline tellurium reacts with water above 100°C, amorphous - above 50°C. Concentrated alkali solutions dissolve tellurium to form tellurides and tellurites. Hydrochloric acid and dilute H2 SO4 do not affect tellurium, conc. H2 SO4 dissolves it, the resulting red solutions contain the cation. HNO3 oxidizes tellurium to tellurous acid H2 TeO3 (tellurite salts):

Te + HNO3 → H2 TeO3 + 4NO2 + H2O

Strong oxidizing agents (HClO3, KMnO4, etc.) oxidize to telluric acid H2 TeO4 (tellurate salts):

4Te + 3HClO4 + 4H2O → 4H2 TeO4 + 3HCl

Te + 3H2 O2 → H2 TeO4 + 2H2 O

Tellurium dissolves in solutions of sulfides and polysulfides of alkali metals (with the formation of thiotellurides and thiotellurites). Reacts with Ag salt solutions. Does not dissolve in CS2. It reacts with Cl2, F2 and Br2 at room temperature, with I2 - when heated, alloys with S, P (it does not form compounds), As (giving As2 Te3), with Si (with the formation of Si2 Te3 and SiTe), with Se (forming solid solutions during crystallization). It does not directly interact with boron and carbon; when heated, it forms gaseous unstable TeCO carbonyl. When fused with metals, tellurides are obtained.

Hydrogen telluride H2 Te is a colorless gas with an unpleasant odor; in the liquid state greenish-yellow, crystalline-lemon-yellow; t. kip. - 2°C, so pl. - 51 °С; dense 5.81 g/l; for gas; and in dry air at room temperature it slowly decomposes, in moist air it oxidizes to tellurium; when heated in air, it burns, giving TeO2; solubility in water 0.1 M, aqueous solution-weak acid, K1 2 10-3; strong reducing agent; obtained by the interaction of Al2 Te3 with hydrochloric acid, as well as by electrolysis of a solution of H2 SO4 with a tellurium cathode at 0°C; used to produce high purity tellurium.

TeF6 hexafluoride is a colorless gas; m.p. - 37.8°С, temp. -38.6°C; dense 10.7 g/l; stable in dry air, does not affect glass; dissolves in water, gradually hydrolyzing with the formation of fluorotelluric acids TeFn (OH) 6-n, where n is from 1 to 4, and ultimately telluric acid; forms compounds with metal fluorides, for example. Ag and Ba; obtained by fluorination of tellurium when heated. Tetrafluoride TeF4 - orthorhombic crystals; m.p. 129.6°С, b.p. 194°C (with decomposition); density 4.22 g/cm3; very hygroscopic, easily hydrolyzed; with alkali metal fluorides forms pentafluorotellurates M; obtained by the action of SeF4 on TeO2. Fluorides tellurafluorinating agents.

TeCl4 tetrachloride - yellow crystals; m.p. 224°С, b.p. 381.8°C; dense 3.01 g/cm3; ur-tion of the temperature dependence of vapor pressure \ gp (mm Hg) \u003d 8.791 - - 3941 / T (497 - 653); very hygroscopic, hydrolyzes with water; in concentrated HC1 solution, forming chlorotelluric acid H2 TeC16; from hydrochloric acid solutions it is extracted with tributyl phosphate and other organic solvents; with alkali metal chlorides it forms hexa-M2 [TeCl6] and pentachlorotellurates M[TeC15], with chlorides of Al, Fe(III), Zr and other complexes with cations, for example, TeC13; obtained by chlorination of tellurium; TeCl4 is the starting material for the production of high purity tellurium. Brown TeCl2 dichloride is stable in vapors and can be condensed into a liquid. Two crystalline lower chlorides were also obtained - silver-gray Te2 Cl3 and metastable black Te2 Cl with a metallic sheen.

TeS2 and TeS3 sulfides, which decompose when heated, can be obtained by precipitation from aqueous solutions; TeS7 and Te7 S10 are known. Thiotellurates (eg, Na2 TeS3) can be obtained by dissolving tellurium in a solution of alkali metal polysulfides or sulfur in solutions of polytellurides, as well as by fusion. Thiotellurates are intermediates in some tellurium recovery processes.

Application. The most important field of application of tellurium is the synthesis of the decomposition of tellurides with semiconductor properties. Tellurium is also used in metallurgy for alloying cast iron and steel, Pb, Cu (to increase their mechanical and chemical resistance). Tellurium and its compounds are used in the production of catalysts, spec. glasses, insecticides, herbicides, etc.

Polonium

Natural compounds and obtaining polonium. A radioactive chemical element of group VI of the periodic system, an analogue of tellurium. Atomic number 84. Has no stable isotopes. There are 27 known radioactive isotopes of polonium with mass numbers from 192 to 218, of which seven (with mass numbers from 210 to 218) are found in nature in very small quantities as members of the radioactive series of uranium, thorium and actinium, the remaining isotopes were obtained artificially. The longest-lived isotopes of polonium are artificially produced 209 Rho ( t 1/2 = 102 years) and 208 Rho ( t 1/2 \u003d 2.9 years), as well as 210 Rho contained in radium-uranium ores ( t 1/2 = 138.4 days). The content of 210 Rho in the earth's crust is only 2 10–14%; 1 ton of natural uranium contains 0.34 g of radium and fractions of a milligram of polonium-210. The shortest-lived known isotope of polonium is 213 Po ( t 1/2 = 3 10–7 s). The lightest isotopes of polonium are pure alpha emitters, while the heavier isotopes simultaneously emit alpha and gamma rays. Some isotopes decay by electron capture, and the heaviest ones also exhibit very weak beta activity. Different isotopes of polonium have historical names adopted as early as the beginning of the 20th century, when they were obtained as a result of a chain of decays from the "parent element": RaF (210 Po), AcC "(211 Po), ThC" (212 Po), RaC " (214 Po), AcA (215 Po), ThA (216 Po), RaA (218 Po).

Polonium-210 is synthesized by neutron irradiation of natural bismuth (it contains only 208 Bi) in nuclear reactors (the beta-active isotope of bismuth-210 is formed intermediately): 208 Bi + n → 210 Bi → 210 Po + e. When bismuth is irradiated with accelerated protons, polonium-208 is formed, it is separated from bismuth by sublimation in a vacuum - as M. Curie did. In the USSR, Zinaida Vasilievna Ershova (1905–1995) developed the method for isolating polonium. In 1937 she was sent to Paris to the Institute of Radium in the laboratory of M.Curie (headed at that time by Irene Joliot-Curie). As a result of this business trip, colleagues began to call her "Russian Madame Curie." Under the scientific guidance of Z.V. Ershova, a permanent, environmentally friendly production of polonium was created in the country, which made it possible to implement the national program for launching lunar rovers, in which polonium was used as a heat source.

Long-lived isotopes of polonium have not yet received significant practical application due to the complexity of their synthesis. Nuclear reactions can be used to obtain them.

207Pb + 4He® 208Po + 3n,

208 Bi + 1 H® 208 Po + 2n,

208 Bi + 2D® 208 Po + 3n,

208 Bi + 2D® 208 Po + 2n,

where 4 He are alpha particles, 1 H are accelerated protons, 2 D are accelerated deuterons (deuterium nuclei).

properties of polonium. Tellurium already partially exhibits metallic properties, while polonium is a soft silvery-white metal. Due to the strong radioactivity, it glows in the dark and gets very hot, so continuous heat removal is needed. The melting point of polonium is 254 ° C (slightly higher than that of tin), the boiling point is 962 ° C, therefore, even with a slight heating, polonium sublimates. The density of polonium is almost the same as that of copper - 9.4 g/cm3. In chemical research, only polonium-210 is used; longer-lived isotopes are practically not used due to the difficulty of obtaining them with the same chemical properties.

The chemical properties of metallic polonium are close to those of its closest analogue, tellurium; it exhibits oxidation states of –2, +2, +4, +6. In air, polonium slowly oxidizes (quickly when heated to 250 ° C) with the formation of red dioxide PoO2 (when cooled, it becomes yellow as a result of rearrangement of the crystal lattice). Hydrogen sulfide from solutions of polonium salts precipitates black sulfide PoS.

The strong radioactivity of polonium is reflected in the properties of its compounds. So, in dilute hydrochloric acid, polonium slowly dissolves with the formation of pink solutions (the color of Po2+ ions):

Po + 2HCl ® PoCl2 + H2 ,

however, under the action of its own radiation, the dichloride is converted to yellow PoCl4. Dilute nitric acid passivates polonium, while concentrated nitric acid quickly dissolves it:

Po + 8HNO3 → Po(NO3)4 + 4NO2 + 4H2O

With non-metals of group VI, polonium is related by the reaction with hydrogen to form the volatile hydride PoH2 (mp. -35 ° C, b.p. + 35 ° C, easily decomposes), the reaction with metals (when heated) with the formation of solid black polonides (Na2Po, MgPo, CaPo, ZnPo, HgPo, PtPo, etc.) and reaction with molten alkalis to form polonides:

3Po + 6NaOH ® 2Na2Po + Na2PoO3 + H2O.

Polonium reacts with chlorine when heated to form bright yellow PoCl4 crystals, red PoBr4 crystals are obtained with bromine, and polonium reacts with iodine already at 40 ° C to form black volatile iodide PoI4. White polonium tetrafluoride PoF4 is also known. When heated, the tetrahalides decompose to form more stable dihalides:

PoCl4 ® PoCl2 + Cl2 .

In solutions, polonium exists in the form of Po2+, Po4+ cations, PoO32–, PoO42– anions, as well as various complex ions, for example, PoCl62–.

The use of polonium Polonium-210 emits alpha rays with an energy of 5.3 MeV, which are decelerated in solid matter, passing only thousandths of a millimeter and giving up their energy in the process. Its lifetime makes it possible to use polonium as an energy source in atomic batteries of spacecraft: only 7.5 g of polonium is enough to obtain a power of 1 kW. In this respect, it is superior to other compact "atomic" energy sources. Such an energy source worked, for example, on Lunokhod-2, heating the equipment during a long moonlit night. Of course, the power of polonium energy sources decreases over time - by half every 4.5 months, but longer-lived polonium isotopes are too expensive. Polonium is also conveniently used to study the effects of alpha radiation on various substances. As an alpha emitter, polonium mixed with beryllium is used to make compact neutron sources:

9 Be + 4 He ® 12 C + n.

Boron can be used instead of beryllium in such sources. In 2004, inspectors from the International Atomic Energy Agency (IAEA) were reported to have discovered a polonium production program in Iran. This led to the suspicion that it could be used in a beryllium source to "start" with the help of neutrons a nuclear chain reaction in uranium, leading to a nuclear explosion.

Polonium, when it enters the body, can be considered one of the most toxic substances: for 210 Rho, the maximum permissible content in the air is only 40 billionths of a microgram per 1 m3 of air, i.e. Polonium is 4 trillion times more toxic than hydrocyanic acid. The alpha particles emitted by the polonium (and to a lesser extent also the gamma rays) cause damage, which destroy tissues and cause malignant tumors. Polonium atoms can be formed in human lungs as a result of the decay of radon gas in them. In addition, metallic polonium is able to easily form the smallest aerosol particles. Therefore, all work with polonium is carried out remotely in sealed boxes.

The discovery of polonium. The existence of an element with the atomic number 84 was predicted by D.I. Mendeleev in 1889 - he called it ditellurium (in Sanskrit - the “second” tellurium) and suggested that its atomic mass would be close to 212. Of course, Mendeleev could not foresee that this element is unstable. Polonium is the first radioactive element, discovered in 1898 by the Curies in search of a source of strong radioactivity in certain minerals. When it turned out that uranium resin ore radiates more strongly than pure uranium, Marie Curie decided to chemically isolate a new radioactive chemical element from this compound. Before that, only two weakly radioactive chemical elements were known - uranium and thorium. Curie began with the traditional qualitative chemical analysis of the mineral according to the standard scheme, which was proposed by the German analytical chemist K.R. Fresenius (1818–1897) as early as 1841 and by which many generations of students for almost a century and a half determined cations by the so-called "hydrogen sulfide method ". At the beginning she had about 100 g of the mineral; then American geologists gave Pierre Curie another 500 g. Carrying out a systematic analysis, M. Curie each time checked individual fractions (precipitates and solutions) for radioactivity using a sensitive electrometer invented by her husband. Inactive fractions were discarded, active ones were analyzed further. She was assisted by one of the leaders of the chemical workshop at the School of Physics and Industrial Chemistry, Gustav Bemon.

First of all, Curie dissolved the mineral in nitric acid, evaporated the solution to dryness, dissolved the residue in water, and passed a stream of hydrogen sulfide through the solution. At the same time, a precipitate of metal sulfides precipitated; according to the Fresenius method, this precipitate could contain insoluble sulfides of lead, bismuth, copper, arsenic, antimony, and a number of other metals. The precipitate was radioactive, despite the fact that the uranium and thorium remained in solution. She treated the black precipitate with ammonium sulfide to separate the arsenic and antimony, which under these conditions form soluble thiosalts such as (NH4)3 AsS4 and (NH4)3 SbS3. The solution did not detect radioactivity and was discarded. Lead, bismuth and copper sulfides remained in the sediment.

The part of the precipitate that did not dissolve in ammonium sulfide was again dissolved by Curie in nitric acid, sulfuric acid was added to the solution, and it was evaporated on a burner flame until thick white SO3 fumes appeared. Under these conditions, volatile nitric acid is completely removed, and metal nitrates are converted to sulfates. After cooling the mixture and adding cold water, insoluble lead sulfate PbSO4 turned out to be in the precipitate - there was no radioactivity in it. She discarded the precipitate, and added a strong solution of ammonia to the filtered solution. At the same time, a precipitate fell out again, this time - white; it contained a mixture of basic bismuth sulfate (BiO)2 SO4 and bismuth hydroxide Bi(OH)3. In the solution, complex copper ammonia SO4 of bright blue color remained. The white precipitate, unlike the solution, turned out to be highly radioactive. Since the lead and copper had already been separated, the white precipitate contained bismuth and an admixture of the new element.

Curie again converted the white precipitate into dark brown sulfide Bi2 S3, dried it, and heated it in an evacuated ampoule. Bismuth sulfide did not change at the same time (it is resistant to heat and melts only at 685 ° C), however, some vapors were released from the precipitate, which settled in the form of a black film on the cold part of the ampoule. The film was radioactive and apparently contained a new chemical element - an analogue of bismuth in the periodic table. It was polonium - the first discovered radioactive element after uranium and thorium, inscribed in the periodic table (in the same year 1898, radium was discovered, as well as a group of noble gases - neon, krypton and xenon). As it turned out later, polonium easily sublimates when heated - its volatility is about the same as that of zinc.

The Curies were in no hurry to call the black coating on the glass a new element. One radioactivity was not enough. A colleague and friend of Curie, the French chemist Eugene Anatole Demarce (1852–1903), a specialist in the field of spectral analysis (he discovered europium in 1901), examined the emission spectrum of black plaque and found no new lines in it that could indicate the presence of a new element. Spectral analysis is one of the most sensitive methods, allowing the detection of many substances in microscopic quantities invisible to the eye. Nevertheless, in an article published on July 18, 1898, the Curies wrote: “We think that the substance we isolated from uranium resin contains a metal that is not yet known, which is analogous to bismuth in analytical properties. If the existence of a new metal is confirmed, we propose to call it polonium, after the birthplace of one of us” (Polonia in Latin - Poland). This is the only case when a new chemical element, not yet identified, has already received a name. However, it was not possible to obtain weight amounts of polonium - there was too little of it in uranium ore (later polonium was obtained artificially). And it was not this element that glorified the Curie spouses, but radium.

In the VIA-group of the periodic system of elements D.I. Mendeleev includes oxygen, sulfur, selenium, tellurium, polonium. The first four of them are non-metallic in nature. The common name of the elements of this group chalcogens, which is translated from Greek. means "forming ores", indicating their presence in nature.

The electronic formula of the valence shell of atoms of the elements of the VIA group.

The atoms of these elements have 6 valence electrons in the s- and p-orbitals of the outer energy level. Of these, two p-orbitals are half filled.

The oxygen atom differs from the atoms of other chalcogens by the absence of a low-lying d-sublevel. Therefore, oxygen, as a rule, is able to form only two bonds with atoms of other elements. However, in some cases, the presence of lone pairs of electrons at the external energy level allows the oxygen atom to form additional bonds by the donor-acceptor mechanism.

For atoms of other chalcogens, when energy is supplied from outside, the number of unpaired electrons can increase as a result of the transition of s- and p-electrons to the d-sublevel. Therefore, the atoms of sulfur and other chalcogens are able to form not only 2, but also 4, and 6 bonds with atoms of other elements. For example, in an excited state of a sulfur atom, the electrons of the outer energy level can acquire the electronic configuration 3s 2 3p 3 3d 1 and 3s 1 3p 3 3d 2:

Depending on the state of the electron shell, different oxidation states (CO) appear. In compounds with metals and hydrogen, the elements of this group exhibit CO = -2. In compounds with oxygen and non-metals, sulfur, selenium and tellurium can have CO = +4 and CO = +6. In some compounds they exhibit CO = +2.

Oxygen is second only to fluorine in electronegativity. In fluoroxide F 2 O, the oxidation state of oxygen is positive and equal to +2. With other elements, oxygen usually exhibits an oxidation state of -2 in compounds, with the exception of hydrogen peroxide H 2 O 2 and its derivatives, in which oxygen has an oxidation state of -1. In living organisms, oxygen, sulfur and selenium are part of biomolecules in the -2 oxidation state.

In the series O - S - Se-Te - Po, the radii of atoms and ions increase. Accordingly, the ionization energy and relative electronegativity naturally decrease in the same direction.

With an increase in the serial number of elements of the VIA group, the oxidative activity of neutral atoms decreases and the reducing activity of negative ions increases. All this leads to a weakening of the non-metallic properties of chalcogens during the transition from oxygen to tellurium.

With an increase in the atomic number of chalcogens, the characteristic coordination numbers increase. This is due to the fact that during the transition from the p-elements of the fourth period to the p-elements of the fifth and sixth periods, d - and even f-orbitals. So, if for sulfur and selenium the most typical coordination numbers are 3 and 4, then for tellurium - 6 and even 8.

Under normal conditions, hydrogen compounds H 2 E of elements of group VIA, with the exception of water, are gases with a very unpleasant odor. The thermodynamic stability of these compounds decreases from water to hydrogen telluride H 2 Te. In aqueous solutions, they exhibit slightly acidic properties. In the series H 2 O-H 2 S-H 2 Se-H 2 Te, the strength of acids increases.

This is due to the increase in the radii of the E 2- ions and the corresponding weakening of the E-N bonds. In the same direction, the reducing ability of H 2 E increases.

Sulfur, selenium, tellurium form two series of acidic oxides: EO 2 and EO 3. They correspond to acid hydroxides of the composition H 2 EO 3 and H 2 EO 4 . Acids H 2 EO 3 in the free state are unstable. The salts of these acids and the acids themselves exhibit redox duality, since the elements S, Se and Te have an intermediate oxidation state of + 4 in these compounds.

Acids of the composition H 2 EO 4 are more stable and behave like oxidizing agents in reactions (the highest oxidation state of the element is +6).

Chemical properties of oxygen compounds. Oxygen is the most common element in the earth's crust (49.4%). The high content and high chemical activity of oxygen determine the predominant form of existence of most elements of the Earth in the form of oxygen-containing compounds. Oxygen is a part of all vital organic substances - proteins, fats, carbohydrates.

Numerous extremely important life processes, such as respiration, oxidation of amino acids, fats, and carbohydrates, are impossible without oxygen. Only a few plants, called anaerobic, can survive without oxygen.

In higher animals (Fig. 8.7), oxygen enters the blood, combines with hemoglobin, forming an easily dissociating compound oxyhemoglobin. With the blood flow, this compound enters the capillaries of various organs. Here, oxygen is split off from hemoglobin and diffuses through the walls of the capillaries into the tissues. The connection between hemoglobin and oxygen is fragile and is carried out due to the donor-acceptor interaction with the Fe 2+ ion.

At rest, a person inhales about 0.5 m 3 of air per hour. But only 1/5 of the oxygen inhaled with air is retained in the body. However, an excess of oxygen (4 / 5) is necessary to create a high concentration in the blood. This, in accordance with Fick's law, provides a sufficient rate of oxygen diffusion through the capillary walls. Thus, a person actually uses about 0.1 m 3 of oxygen per day.

Oxygen is consumed in tissues. for the oxidation of various substances. These reactions ultimately lead to the formation of carbon dioxide, water and energy storage.

Oxygen is consumed not only in the process of respiration, but also in the process of decay of plant and animal remains. As a result of the process of decay of complex organic substances, their oxidation products are formed: CO 2, H 2 O, etc. Oxygen regeneration occurs in plants.

Thus, as a result of the oxygen cycle in nature, its constant content in the atmosphere is maintained. Naturally, the oxygen cycle in nature is closely related to the carbon cycle (Fig. 8.8).

The element oxygen exists in the form of two simple substances (allotropic modifications): dioxygen(oxygen) O 2 and trioxygen(ozone) O 3 . In the atmosphere, almost all oxygen is contained in the form of oxygen O 2, while the content of ozone is very small. The maximum volume fraction of ozone at a height of 22 km is only 10 -6%.

The oxygen molecule O 2 in the absence of other substances is very stable. The presence of two unpaired electrons in a molecule determines its high reactivity. Oxygen is one of the most active non-metals. With most simple substances, it reacts directly, forming oxides E x O y The degree of oxidation of oxygen in them is -2. In accordance with the change in the structure of the electron shells of atoms, the nature of the chemical bond, and consequently, the structure and properties of oxides in the periods and groups of the system of elements, change regularly. So, in a series of oxides of elements of the second period Li 2 O-BeO-B 2 O 3 -CO 2 -N 2 O 5, the polarity of the chemical bond E-O from group I to group V gradually decreases. In accordance with this, the basic properties are weakened and acid properties are enhanced: Li 2 O is a typical basic oxide, BeO is amphoteric, and B 2 O 3, CO 2 and N 2 O 5 are acid oxides. The acid-base properties change similarly in other periods.

In the main subgroups (A-groups), with an increase in the ordinal number of the element, the ionicity of the E-O bond in oxides usually increases.

Accordingly, the main properties of oxides in the Li-Na-K-Rb-Cs group and other A-groups increase.

The properties of oxides, due to a change in the nature of the chemical bond, are a periodic function of the charge of the nucleus of an atom of an element. This is evidenced, for example, by the change in periods and groups of melting temperatures, enthalpies of oxide formation depending on the charge of the nucleus.

The polarity of the E-OH bond in E(OH) n hydroxides, and, consequently, the properties of the hydroxides naturally change according to the groups and periods of the system of elements.

For example, in IA-, IIA- and IIIA-groups from top to bottom with an increase in the radii of the ions, the polarity of the E-OH bond increases. As a result, ionization E-OH → E + + OH - is easier in water. Accordingly, the basic properties of hydroxides are enhanced. So, in group IA, the main properties of alkali metal hydroxides are enhanced in the series Li-Na-K-Rb-Cs.

In periods from left to right, with decreasing ionic radii and increasing ion charge, the polarity of the E-OH bond decreases. As a result, the ionization of EON ⇄ EO - + H + is easier in water. Accordingly, acidic properties are enhanced in this direction. So, in the fifth period, the hydroxides RbOH and Sr(OH) 2 are bases, In(OH) 3 and Sn(OH) 4 are amphoteric compounds, and H and H 6 TeO 6 are acids.

The most common oxide on earth is hydrogen oxide or water. Suffice it to say that it makes up 50-99% of the mass of any living being. The human body contains 70-80% water. For 70 years of life, a person drinks about 25,000 kg of water.

Due to its structure, water has unique properties. In a living organism, it is a solvent of organic and inorganic compounds, participates in the processes of ionization of molecules of dissolved substances. Water is not only the medium in which biochemical reactions take place, but also actively participates in hydrolytic processes.

The ability of oxygen to form oxygenic complexes with various substances. Previously, examples of O 2 oxygenyl complexes with metal ions - oxygen carriers in living organisms - oxyhemoglobin and oxyhemocyanin were considered:

HbFe 2 + + O 2 → HbFe 2+ ∙O 2

HcCu 2+ + O 2 → HcCu 2+ ∙O 2

where Hb is hemoglobin, Hc is hemocyanin.

Having two lone pairs of electrons, oxygen acts as a donor in these coordination compounds with metal ions. In other compounds, oxygen forms various hydrogen bonds.

At present, much attention is paid to the preparation of oxygenyl complexes of transition metals, which could perform functions similar to those of the corresponding bioinorganic complex compounds. The composition of the internal coordination sphere of these complexes is similar to natural active centers. In particular, complexes of cobalt with amino acids and some other ligands are promising in terms of their ability to reversibly add and donate elemental oxygen. These compounds, to a certain extent, can be considered as substitutes for hemoglobin.

One of the allotropic modifications of oxygen is ozone About 3 . In its properties, ozone is very different from oxygen O 2 - it has higher melting and boiling points, and has a pungent odor (hence its name).

The formation of ozone from oxygen is accompanied by the absorption of energy:

3O 2 ⇄2O 3,

Ozone is produced by the action of an electrical discharge in oxygen. Ozone is formed from O 2 and under the action of ultraviolet radiation. Therefore, during the operation of bactericidal and physiotherapeutic ultraviolet lamps, the smell of ozone is felt.

Ozone is the strongest oxidizing agent. Oxidizes metals, reacts violently with organic substances, at low temperatures oxidizes compounds with which oxygen does not react:

O 3 + 2Ag \u003d Ag 2 O + O 2

PbS + 4O 3 \u003d PbSO 4 + 4O 2

A well-known qualitative reaction:

2KI + O 3 + H 2 O \u003d I 2 + 2KOH + O 2

The oxidative effect of ozone on organic substances is associated with the formation of radicals:

RN + O 3 → RO 2 ∙ + OH ∙

Radicals initiate radical chain reactions with bioorganic molecules - lipids, proteins, DNA. These reactions lead to cell damage and death. In particular, ozone kills microorganisms found in air and water. This is the basis for the use of ozone for the sterilization of drinking water and swimming pool water.

Chemical properties of sulfur compounds. Sulfur is similar in properties to oxygen. But unlike it, it exhibits in compounds not only the oxidation state -2, but also the positive oxidation states +2, +4 and +6. For sulfur, as well as for oxygen, allotropy is characteristic - the existence of several elemental substances - rhombic, monoclinic, plastic sulfur. Due to the lower electronegativity compared to oxygen, the ability to form hydrogen bonds in sulfur is less pronounced. Sulfur is characterized by the formation of stable polymer homochains having a zigzag shape.

The formation of homochains from sulfur atoms is also characteristic of its compounds, which play an essential biological role in life processes. So, in the molecules of the amino acid - cystine there is a disulfide bridge -S-S-:

This amino acid plays an important role in the formation of proteins and peptides. Due to the S-S disulfide bond, the polypeptide chains are bonded to each other (disulfide bridge).

Sulfur is also characterized by the formation of a hydrogen sulfide (sulfhydryl) thiol group -SH, which is present in the amino acid cysteine, proteins, and enzymes.

Biologically important is the amino acid methionine.

The donor of methyl groups in living organisms is S-adenosylmethionine Ad-S-CH 3 - an activated form of methionine, in which the methyl group is connected through S to adenine Ad. The methyl group of methionine in the processes of biosynthesis is transferred to various acceptors of methyl groups RN:

Ad-S-CH 3 + RN → Ad-SH + R-CH 3

Sulfur is quite widespread on Earth (0.03%). In nature, it is present in the form of sulfide (ZnS, HgS, PbS, etc.) and sulfate (Na 2 SO 4 ∙10H 2 O, CaSO 4 ∙2H 2 O, etc.) minerals, as well as in a native state. Powder "precipitated sulfur" is used externally in the form of ointments (5-10-20%) and powders in the treatment of skin diseases (seborrhea, psoriasis). In the body, sulfur oxidation products are formed - polythionic acids with the general formula H 2 S x O 6 ( x = 3-6)

S + O 2 → H 2 S x O 6

Sulfur is a fairly active non-metal. Even with a slight heating, it oxidizes many simple substances, however, it itself is easily oxidized by oxygen and halogens (redox duality).

The oxidation state -2 sulfur shows in hydrogen sulfide and its derivatives - sulfides.

Hydrogen sulfide (dihydrogen sulfide) often found in nature. Contained in the so-called sulfuric mineral waters. It is a colorless gas with an unpleasant odor. It is formed during the decay of plant and, in particular, animal residues under the action of microorganisms. Some photosynthetic bacteria, such as green sulfur bacteria, use dihydrogen sulfide as a hydrogen donor. These bacteria, instead of oxygen O 2, emit elemental sulfur - the product of the oxidation of H 2 S.

Dihydrogen sulfide is a very toxic substance, as it is an inhibitor of the enzyme cytochrome oxidase, an electron carrier in the respiratory chain. It blocks the transfer of electrons from cytochrome oxidase to oxygen O 2 .

Aqueous solutions of H 2 S give a weakly acid reaction according to litmus. Ionization occurs in two stages:

H 2 S ⇄ H + + HS - (I stage)

HS - ⇄ H + + S 2- (stage II)

Sulfuric acid is very weak. Therefore, ionization in the second stage proceeds only in very dilute solutions.

Salts of hydrosulphuric acid are called sulfides. Only alkali, alkaline earth and ammonium sulfides are soluble in water. Acid salts - hydrosulfides E + NS and E 2+ (HS) 2 - are known only for alkali and alkaline earth metals

Being salts of a weak acid, sulfides undergo hydrolysis. The hydrolysis of sulfides of multiply charged metal cations (Al 3+ , Cr 3 + , etc.) often comes to an end, it is practically irreversible.

Sulfides, especially hydrogen sulfide, are strong reducing agents. Depending on the conditions, they can be oxidized to S, SO 2 or H 2 SO 4:

2H 2 S + 3O 2 \u003d 2SO 2 + 2H 2 O (in air)

2H 2 S + O 2 \u003d 2H 2 O + 2S (in air)

3H 2 S + 4HClO 3 \u003d 3H 2 SO 4 + 4HCl (in solution)

Some proteins containing cysteine ​​HSCH 2 CH (NH 2) COOH and an important metabolite coenzyme A, having hydrogen sulfide (thiol) groups -SH, behave in a number of reactions as bioinorganic dihydrogen sulfide derivatives. Proteins containing cysteine, like dihydrogen sulfide, can be oxidized with iodine. With the help of a disulfide bridge formed during the oxidation of thiol groups, cysteine ​​residues of polypeptide chains connect these chains with a cross-link (a crosslink is formed).

Many sulfur-containing E-SH enzymes are irreversibly poisoned by heavy metal ions such as Cu 2+ or Ag+. These ions block thiol groups to form mercaptans, bioinorganic analogues of sulfides:

E-SH + Ag + → E-S-Ag + H +

As a result, the enzyme loses its activity. The affinity of Ag + ions for thiol groups is so high that AgNO 3 can be used to quantify -SH groups by titration.

Sulfur(IV) oxide SO 2 is an acidic oxide. It is obtained by burning elemental sulfur in oxygen or by burning pyrite FeS 2:

S + O 2 \u003d SO 2

4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2

SO 2 - gas with a suffocating odor; very poisonous. When SO 2 is dissolved in water, sulfurous acid H 2 SO 3. This is a medium strength acid. Sulfurous acid, being dibasic, forms salts of two types: medium - sulfites(Na 2 SO 3, K 2 SO 3, etc.) and acidic - hydrosulfites(NaHSO 3 , KHSO 3 and others). Only alkali metal salts and hydrosulfites of the E 2+ (HSO 3) 2 type are soluble in water, where E are elements of various groups.

Oxide SO 2, acid H 2 SO3 and its salts are characterized by redox duality, since sulfur has an intermediate oxidation state of +4 in these compounds:

2Na 2 SO 3 + O 2 \u003d 2Na 2 SO 4

SO 2 + 2H 2 S \u003d 3S ° + 2H 2 O

However, the reducing properties of sulfur compounds (IV) prevail. Thus, sulfites in solutions are oxidized even by air dioxygen at room temperature.

In higher animals, SO 2 oxide acts primarily as an irritant to the mucous membrane of the respiratory tract. This gas is also toxic to plants. In industrial areas, where a lot of coal containing a small amount of sulfur compounds is burned, sulfur dioxide is released into the atmosphere. Dissolving in the moisture on the leaves, SO 2 forms a solution of sulfurous acid, which, in turn, is oxidized to sulfuric acid H 2 SO 4:

SO 2 + H 2 O \u003d H 2 SO 3

2H 2 SO 3 + O 2 \u003d 2H 2 SO 4

Atmospheric moisture with dissolved SO 2 and H 2 SO 4 often falls in the form of acid rain, leading to the death of vegetation.

When a solution of Na 2 SO 3 is heated with sulfur powder, sodium thiosulfate:

Na 2 SO 3 + S \u003d Na 2 S 2 O 3

Crystal hydrate Na 2 S 2 O 3 ∙ 5H 2 O stands out from the solution. Sodium thiosulfate - salt thiosulfuric acid H 2 S 2 O 3.

Thiosulfuric acid is very unstable and decomposes into H 2 O, SO 2 and S. Sodium thiosulfate Na 2 S 2 O 3 ∙5H 2 O is used in medical practice as an antitoxic, anti-inflammatory and desensitizing agent. As an antitoxic agent, sodium thiosulfate is used for poisoning with mercury, lead, hydrocyanic acid and its salts. The mechanism of action of the drug is obviously associated with the oxidation of thiosulfate ion to sulfite ion and elemental sulfur:

S 2 O 3 2- → SO 3 2- + S °

Lead and mercury ions that enter the body with food or air form poorly soluble non-toxic sulfites:

Pb 2+ + SO 3 2- = PbSO 3

Cyanide ions interact with elemental sulfur to form less toxic thiocyanates:

СN - + S° = NСS -

Sodium thiosulfate is also used to treat scabies. After rubbing the solution into the skin, repeated rubbing of a 6% HCl solution is done. As a result of the reaction with HCl, sodium thiosulfate decomposes into sulfur and sulfur dioxide:

Na 2 S 2 O 3 + 2HCl \u003d 2NaCl + SO 2 + S + H 2 O

which have a detrimental effect on scabies mites.

Oxide sulfur (VI) SO 3 is a volatile liquid. When interacting with water, SO 3 forms sulfuric acid:

SO 3 + H 2 O \u003d H 2 SO 4

The structure of sulfuric acid molecules corresponds to sulfur in sp 3 - hybrid state.

Sulfuric acid is a strong dibasic acid. In the first stage, it is almost completely ionized:

H 2 SO 4 ⇄ H + + HSO 4 -,

Ionization in the second stage proceeds to a lesser extent:

HSO 4 - ⇄ H + + SO 4 2-,

Concentrated sulfuric acid is a strong oxidizing agent. It oxidizes metals and non-metals. Usually, the product of its reduction is SO 2, although depending on the reaction conditions (metal activity, temperature, acid concentration), other products (S, H 2 S) can be obtained.

Being a dibasic acid, H 2 SO 4 forms two types of salts: medium - sulfates(Na 2 SO 4, etc.) and acidic - hydrosulphates(NaHSO 4 , KHSO 4 and others). Most sulfates are highly soluble in water. Many sulfates are released from solutions in the form of crystalline hydrates: FeSO 4 ∙7H 2 O, CuSO 4 ∙5H 2 O. Sulfates BaSO 4, SrSO 4 and PbSO 4 are practically insoluble. Slightly soluble calcium sulfate CaSO 4 . Barium sulfate is insoluble not only in water, but also in dilute acids.

In medical practice, sulfates of many metals are used as medicines Na 2 SO 4 ∙ 10H 2 O - as a laxative, MgSO 4 ∙ 7H 2 O - for hypertension, as a laxative and as a choleretic agent, copper sulfate CuSO 4 ∙ 5H 2 O and ZnSO 4 ∙7H 2 O - as antiseptic, astringent, emetics, barium sulfate BaSO 4 - as a contrast agent in x-ray examination of the esophagus and stomach

Selenium and tellurium compounds. Tellurium and especially selenium are chemically similar to sulfur. However, strengthening the metallic properties of Se and Te increases their tendency to form stronger ionic bonds. The similarity of physical and chemical characteristics: the radii of E 2- ions, coordination numbers (3, 4) - determines the interchangeability of selenium and sulfur in compounds. So, selenium can replace sulfur in the active centers of enzymes. Replacing the hydrogen sulfide group -SH with the hydrogen selenide group -SeH changes the course of biochemical processes in the body. Selenium can act as both a synergist and an antagonist of sulfur.

With hydrogen, Se and Te form very poisonous gases, similar to H 2 S, H 2 Se and H 2 Te. Dihydrogen selenide and dihydrogen telluride are strong reducing agents. In the series H 2 S-H 2 Se-H 2 Te, the reducing activity increases.

For H 2 Se isolated as medium salts - selenides(Na 2 Se, etc.), and acid salts - hydroselenides(NaHSe and others). For H 2 Te, only medium salts are known - tellurides.

Compounds Se (IV) and Te (IV) with oxygen, in contrast to SO 2, are solid crystalline substances SeO 2 and TeO 2.

Selenic acid H 2 SeO 3 and its salts selenites, for example, Na 2 SeO 3, are oxidizing agents of medium strength. So, in aqueous solutions, they are reduced to selenium by such reducing agents as SO 2, H 2 S, HI, etc.:

H 2 SeO 3 + 2SO 2 + H 2 O \u003d Se + 2H 2 SO 4

Obviously, the ease of reduction of selenites to the elemental state determines the formation in the body of biologically active selenium-containing compounds, such as selenocysteine.

SeO 3 and TeO 3 are acidic oxides. Oxygen acids Se (VI) and Te (VI) - selenic H 2 SeO 4 and tellurium H 6 TeO 6 - crystalline substances with strong oxidizing properties. The salts of these acids are named accordingly. selenates and tellurates.

In living organisms, selenates and sulfates are antagonists. Thus, the introduction of sulfates leads to the excretion of excess selenium-containing compounds from the body.

ELEMENTS VI A subgroups

(O, S, Se, Te, Po)

general characteristics

Oxygen

Sulfur

Selenium and tellurium

General characteristics of the elements

The VI A subgroup of PS includes the elements: oxygen, sulfur, selenium, tellurium and polonium. For sulfur, selenium, tellurium and polonium, a common name is used - chalcogens. Oxygen, sulfur, selenium and tellurium are non-metals, while polonium is a metal. Polonium is a radioactive element, in nature it is formed in small quantities during the radioactive decay of radium, therefore its chemical properties are poorly understood.

Table 1

Main characteristics of chalcogens

Characteristics	O	S	Se	Those
Atomic radius, nm	0,066	0,104	0,117	0,136
Ionic radius E 2-, nm	0,140	0,184	0,198	0,221
Ionization potential, eV	13,62	10,36	9,75	9,01
Electron affinity, eV	1,47	2,08	2,02	1,96
Electronegativity (according to Pauling)	3,44	2,58	2,55	2,10
Bond enthalpy, kJ/mol E –E E = E	- 146 - 494	- 265 - 421	- 192 - 272	- 218 - 126
Melting point, °С
Boiling point, °C	- 183
Density, g / cm 3	1.43 (liquid)	2,07	4,80	6,33
Content in the earth's crust, % (wt.)	49,13	0,003	1.4 10 -5	1 10 -7
Mass numbers of natural isotopes	16, 17, 18	32, 33, 34, 35	74, 76, 77, 78, 80, 82	120, 122, 123, 124, 125, 126 128, 130
The state of aggregation at Art. conditions of the most stable allotropic form. color	colorless gas	Crystal. yellow substance	Crystal. gray matter	Crystal. silvery white substance
Crystal cell	Molecular in TV. form	molecular	molecular	molecular
Composition of molecules	About 2	S8	Se ∞	Te ∞

According to the structure of the outer electronic layer, the considered elements belong to the p-elements. Of the six electrons in the outer layer, two are unpaired, which determines their valence of two. For atoms of sulfur, selenium, tellurium and polonium in an excited state, the number of unpaired electrons can be 4 and 6. That is, these elements can be four - and hexavalent. All elements have high electronegativity values, and the EO of oxygen is second only to fluorine. Therefore, in compounds they exhibit art. oxidation -2, -1, 0. The ionization potentials of sulfur, selenium and tellurium atoms are small, and these elements in compounds with halogens have oxidation states of +4 and +6. Oxygen has a positive oxidation state in fluorine compounds and in ozone.

Atoms can form molecules with a double bond O 2, ... and join in chains E - E - ... - E -, which can exist both in simple and in complex substances. In terms of chemical activity and oxidizing ability, chalcogens are inferior to halogens. This is indicated by the fact that in nature oxygen and sulfur exist not only in a bound, but also in a free state. The lower activity of chalcogens is largely due to a stronger bond in the molecules. In general, chalcogens are among the highly reactive substances, the activity of which sharply increases with increasing temperature. Allotropic modifications are known for all substances of this subgroup. Sulfur and oxygen practically do not conduct electric current (dielectrics), selenium and tellurium are semiconductors.

When moving from oxygen to tellurium, the tendency of elements to form double bonds with small atoms (C, N, O) decreases. The inability of large atoms to form π-bonds with oxygen is especially evident in the case of tellurium. So, in tellurium there are no acid molecules H 2 TeO 3 and H 2 TeO 4 (meta-forms), as well as TeO 2 molecules. Tellurium dioxide exists only in the form of a polymer, where all oxygen atoms are bridging: Te - O - Te. Telluric acid, in contrast to sulfuric and selenic acid, occurs only in the ortho form - H 6 TeO 6, where, as in TeO 2, the Te atoms are connected to the O atoms only by σ-bonds.

The chemical properties of oxygen differ from those of sulfur, selenium and tellurium. On the contrary, there is much in common in the properties of sulfur, selenium and tellurium. When moving through the group from top to bottom, one should note an increase in acidic and reducing properties in a series of compounds with hydrogen H 2 E; an increase in oxidizing properties in a series of similar compounds (H 2 EO 4, EO 2); decrease in thermal stability of hydrogen chalcogens and salts of oxygen acids.

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