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 studied.
Table 1
Main characteristics of chalcogens
| Characteristics | ABOUT | 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.
Elements of group VI of the main subgroup are called chalcogens. These include oxygen, sulfur, selenium, tellurium and polonium. The word "chalcogen" consists of two Greek words meaning "copper" or "ore" and "begotten".
Description
Chalcogens in nature are found most often in the composition of ore - sulfides, pyrites, oxides, selenides. Chalcogens include non-metals and metals. In the group from top to bottom, the properties change as follows:
- metallic properties are enhanced;
- the properties of the oxidizing agent weaken;
- electronegativity decreases;
- thermal stability weakens.
General characteristics of the chalcogen group:
- non-metals - oxygen, sulfur, selenium;
- metals - tellurium, polonium;
- valency: II - O; IV and VI - S; II, IV, VI - Se, Te, Po;
- electronic configuration - ns 2 np 4;
- hydrides - H 2 R;
- oxides - RO 2, RO 3;
- oxygen acids - H 2 RO 3, H 2 RO 4.
Rice. 1. Chalcogens.
According to their electronic structure, chalcogens are p-elements. There are six electrons in the outer energy level. Before the completion of the p-orbital, two electrons are missing, therefore, in compounds, chalcogens exhibit the properties of an oxidizing agent. With an increase in the number of energy levels in the group, the bond with external electrons weakens, so tellurium and polonium are reducing agents.
Being on the border of metals and non-metals, tellurium belongs to metalloids or semi-metals. It is an analogue of sulfur and selenium, but less active.
Rice. 2. Tellurium.
Properties
The most active element of the chalcogen group is oxygen. It is a powerful oxidizing agent that exhibits four oxidation states - -2, -1, +1, +2.
The main properties of chalcogens are presented in the table.
|
Element |
Physical Properties |
Chemical properties |
|
Oxygen (O) |
Gas. It forms two modifications - O 2 and O 3 (ozone). O 2 is odorless and tasteless, poorly soluble in water. Ozone is a bluish odorless gas that is highly soluble in water. |
Reacts with metals, non-metals |
|
A typical non-metal. Solid substance with a melting point of 115°C. Insoluble in water. There are three modifications - rhombic, monoclinic, plastic. Oxidation state - -2, -1, 0, +1, +2, +4, +6 |
Reacts with oxygen, halogens, non-metals, metals |
|
|
Brittle solid. Semiconductor. It has three modifications - gray, red, black selenium. Oxidation state - -2, +2, +4, +6 |
Reacts with alkali metals, oxygen, water |
|
|
Outwardly similar to metal. Semiconductor. Oxidation state - -2, +2, +4, +6 |
Reacts with oxygen, alkalis, acids, water, metals, non-metals, halogens |
|
|
Polonium (Po) |
Silvery radioactive metal. Oxidation level - +2, +4, +6 |
Reacts with oxygen, halogens, acids |
Artificially created livermorium (Lv) or unungexium (Uuh) are also considered chalcogens. It is the 116th element of the periodic table. Shows strong metallic properties.
Rice. 3. Livermorium.
What have we learned?
Chalcogens are elements of the sixth group of Mendeleev's periodic table. The group contains three nonmetals (oxygen, sulfur, selenium), a metal (polonium) and a semimetal (tellurium). Therefore, chalcogens are both oxidizing and reducing agents. Metallic properties are enhanced in the group from top to bottom: oxygen is a gas, polonium is a solid metal. The chalcogens also include artificially synthesized livermorium with strong metallic properties.
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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 (nickel electrodes); The main source of industrial production of oxygen is air, which is liquefied and then fraction yut. 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 already 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 ° C. +O2=СO2 When complex substances are burned in excess oxygen, oxides of the corresponding elements are formed: 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 pyrite - FeS2, chalcopyrite - CuFeS2, cinnabar - HgS. The minerals of the second group include gypsum CaSO 4 2H2O, mirabilite (Glauber's salt) - Na2SO4 10H2O, kieserite - MgSO4 H2O. Sulfur is found in animal and plant organisms, 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 3. In reactions with metals and with hydrogen, sulfur behaves like a typical oxidizing agent, and in the presence of strong oxidizing agents it exhibits reducing properties. S + 3F2 \u003d 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)
CHALCOGENES
SUB-GROUP VIA. CHALCOGENES
OXYGEN
The element oxygen O is the eighth element of the Periodic Table of Elements and the first element of the VIA subgroup (Table 7a). This element is most abundant in the earth's crust, accounting for about 50% (wt.). The air we breathe contains CHALCOGENES, 20% of oxygen is in a free (unbound) state, and 88% of oxygen is in the hydrosphere in a bound state in the form of water H2O.
The most common isotope is 168O. The nucleus of such an isotope contains 8 protons and 8 neutrons. Significantly less common (0.2%) isotope with 10 neutrons, 188O. Even less common (0.04%) is the 9 neutron isotope, 178O. The weighted average mass of all isotopes is 16.044. Since the atomic mass of the carbon isotope with mass number 12 is exactly 12.000 and all other atomic masses are based on this standard, the atomic mass of oxygen according to this standard should be 15.9994.
Oxygen is a diatomic gas, like hydrogen, nitrogen and the halogens fluorine, chlorine (bromine and iodine also form diatomic molecules, but they are not gases). Most of the oxygen used in industry comes from the atmosphere. To do this, relatively inexpensive methods have been developed to liquefy chemically purified air using compression and refrigeration cycles. Liquefied air is slowly heated, while more volatile and easily vaporized compounds are released, and liquid oxygen accumulates. This method is called fractional distillation or distillation of liquid air. In this case, the contamination of oxygen with an admixture of nitrogen is inevitable, and in order to obtain high-purity oxygen, the rectification process is repeated until the complete removal of nitrogen.
See also AIR.
At a temperature of 182.96 ° C and a pressure of 1 atm, oxygen turns from a colorless gas into a pale blue liquid. The presence of color indicates that the substance contains molecules with unpaired electrons. At 218.7°C, oxygen solidifies. Gaseous O2 is 1.105 times heavier than air, and at 0 ° C and 1 atm 1 liter of oxygen has a mass of 1.429 g. The gas is slightly soluble in water (CHALCOGENES 0.30 cm 3 / l at 20 ° C), but this is important for the existence of life in water. Large masses of oxygen are used in the steel industry to quickly remove undesirable impurities, primarily carbon, sulfur and phosphorus, in the form of oxides during the blowing process or directly by blowing oxygen through the melt. One of the important uses of liquid oxygen is as a propellant oxidizer. Oxygen stored in cylinders is used in medicine to enrich the air with oxygen, as well as in technology for welding and cutting metals.
The formation of oxides. Metals and non-metals react with oxygen to form oxides. Reactions can occur with the release of a large amount of energy and be accompanied by a strong glow, flash, burning. Flash light is produced by the oxidation of aluminum or magnesium foil or wire. If gases are formed during oxidation, they expand as a result of the release of heat of reaction and can cause an explosion. Not all elements react with oxygen to release heat. Nitrogen oxides, for example, are formed with the absorption of heat. Oxygen reacts with elements to form oxides of the corresponding elements a) in normal or b) in high oxidation state. Wood, paper and many natural substances or organic products containing carbon and hydrogen burn according to type (a), forming, for example, CO, or according to type (b), forming CO2.
Ozone. In addition to atomic (monatomic) oxygen O and molecular (diatomic) oxygen O2, there is ozone, a substance whose molecules consist of three oxygen atoms O3. These forms are allotropic modifications. By passing a quiet electric discharge through dry oxygen, ozone is obtained:
3O2 2O3 Ozone has a strong irritating odor and is often found near electric motors or power generators. Ozone at the same temperatures is chemically more active than oxygen. It usually reacts with the formation of oxides and the release of free oxygen, for example: Hg + O3 -> HgO + O2 Ozone is effective for water purification (disinfection), for bleaching fabrics, starch, oil purification, for drying and aging wood and tea, in the production of vanillin and camphor. See OXYGEN.
SULFUR, SELENIUM, TELLURIUM, POLONIUM
In the transition from oxygen to polonium in the VIA subgroup, the change in properties from non-metallic to metallic is less pronounced than in the elements of the VA subgroup. The electronic structure of ns2np4 chalcogens suggests the acceptance of electrons rather than their return. Partial withdrawal of electrons from the active metal to the chalcogen is possible with the formation of a compound with a partially ionic bond, but not to the same degree of ionicity as a similar compound with oxygen. Heavy metals form chalcogenides with a covalent bond, the compounds are colored and completely insoluble.
molecular forms. The formation of an octet of electrons around each atom is carried out in the elemental state due to the electrons of neighboring atoms. As a result, for example, in the case of sulfur, a cyclic S8 molecule is obtained, constructed according to the corona type. There is no strong bond between the molecules, so sulfur melts, boils and evaporates at low temperatures. Selenium, which forms the Se8 molecule, has a similar structure and set of properties; tellurium probably forms Te8 chains, but this structure has not been definitely established. The molecular structure of polonium is also not clear. The complexity of the structure of molecules determines the various forms of their existence in the solid, liquid and gaseous state (allotropy); this property, obviously, is a distinctive feature of chalcogens among other groups of elements. The most stable form of sulfur is the a-form, or rhombic sulfur; the second metastable form b, or monoclinic sulfur, which can be converted to a-sulfur on storage. Other modifications of sulfur are shown in the diagram:
A-Sulfur and b-Sulfur are soluble in CS2. Other forms of sulfur are also known. m-Form The viscous liquid is likely formed from the "crown" structure, which explains its rubbery state. With a sharp cooling or condensation of sulfur vapor, powdered sulfur is formed, which is called "sulfur color". Vapors, as well as purple powder, obtained by rapid cooling of vapors, according to the results of studies in a magnetic field, contain unpaired electrons. For Se and Te, allotropy is less characteristic, but has a general similarity with sulfur, with selenium modifications similar to sulfur modifications.
reactivity. All elements of the VIA subgroup react with one-electron donors (alkali metals, hydrogen, methyl radical HCH3), forming compounds of the RMR composition, i.e. showing a coordination number of 2, such as HSH, CH3SCH3, NaSNa and ClSCl. Six valence electrons coordinate around the chalcogen atom, two on the valence s-shell and four on the valence p-shell. These electrons can participate in the formation of a bond with a stronger electron acceptor (for example, oxygen), which pulls them away to form molecules and ions. Thus, these chalcogens exhibit oxidation states II, IV, VI, forming predominantly covalent bonds. In the chalcogen family, the manifestation of the VI oxidation state weakens with increasing atomic number, since the ns2 electron pair is less and less involved in the formation of bonds in heavier elements (the effect of an inert pair). Compounds with such oxidation states include SO and H2SO2 for sulfur(II); SO2 and H2SO3 for sulfur(IV); SO3 and H2SO4 for sulfur(IV). Compounds of other chalcogens have similar compositions, although there are some differences. There are relatively few odd oxidation states. Methods for extracting free elements from natural raw materials are different for different chalcogens. Large deposits of free sulfur are known in rocks, in contrast to minor amounts of other chalcogens in the free state. Sedimentary sulfur can be extracted by the geotechnological method (flash process): superheated water or steam is pumped through the inner pipe to melt the sulfur, then the molten sulfur is squeezed out to the surface through the outer concentric pipe with compressed air. In this way, clean, cheap sulfur is obtained from deposits in Louisiana and under the Gulf of Mexico off the coast of Texas. Selenium and tellurium are extracted from gas emissions from copper, zinc and lead metallurgy, as well as from silver and lead electrometallurgy sludge. Some plants, where selenium is concentrated, become sources of poisoning of the animal world. Free sulfur finds great use in agriculture as a powdered fungicide. Only in the USA about 5.1 million tons of sulfur is used annually for various processes and chemical technologies. A lot of sulfur is consumed in the production of sulfuric acid.
Separate classes of chalcogen compounds, especially halides, differ greatly in properties.
Hydrogen compounds. Hydrogen reacts slowly with chalcogens to form H2M hydrides. There is a big difference between water (oxygen hydride) and hydrides of other chalcogens, which have a disgusting smell and are poisonous, and their aqueous solutions are weak acids (the strongest of them is H2Te). Metals react directly with chalcogens to form chalcogenides (eg sodium sulfide Na2S, potassium sulfide K2S). Sulfur in aqueous solutions of these sulfides forms polysulfides (for example, Na2Sx). Chalcogen hydrides can be displaced from acidified solutions of metal sulfides. Thus, H2Sx sulfanes are isolated from acidified Na2Sx solutions (where x can be greater than 50; however, only sulfanes with x ∼ 6 have been studied).
Halides. Chalcogens react directly with halogens to form halides of various compositions. The range of reacting halogens and the stability of the resulting compounds depend on the ratio of the chalcogen and halogen radii. The possibility of forming a halide with a high oxidation state of chalcogen decreases with increasing atomic mass of the halogen, since the halide ion will be oxidized to halogen, and the chalcogen will be reduced to free chalcogen or chalcogen halide in a low oxidation state, for example: TeI6 -> TeI4 + I2 The oxidation state I for sulfur may be realized in the compound (SCl)2 or S2Cl2 (this composition is not established reliably enough). The most unusual of the sulfur halides is SF6, which is highly inert. Sulfur in this compound is so strongly shielded by fluorine atoms that even the most aggressive substances have practically no effect on SF6. From Table. 7b that sulfur and selenium do not form iodides.
Complex chalcogen halides are known, which are formed by the interaction of a chalcogen halide with halide ions, for example,
TeCl4 + 2Cl= TeCl62.
Oxides and oxoacids. Chalcogen oxides are formed by direct interaction with oxygen. Sulfur burns in air or oxygen to form SO2 and SO3 impurities. Other methods are used to obtain SO3. When SO2 interacts with sulfur, the formation of SO is possible. Selenium and tellurium form similar oxides, but they are much less important in practice. The electrical properties of oxides of selenium and, especially, pure selenium determine the growth of their practical application in electronics and the electrical industry. Alloys of iron and selenium are semiconductors and are used to make rectifiers. Since the conductivity of selenium depends on light and temperature, this property is used in the manufacture of photocells and temperature sensors. Trioxides are known for all elements of this subgroup, except for polonium. The catalytic oxidation of SO2 to SO3 underlies the industrial production of sulfuric acid. Solid SO3 has allotropic modifications: feather-shaped crystals, asbestos-like structure, ice-like structure and polymeric cyclic (SO3)3. Selenium and tellurium dissolve in liquid SO3, forming interchalcogenic compounds such as SeSO3 and TeSO3. Obtaining SeO3 and TeO3 is associated with certain difficulties. SeO3 is obtained from a gas mixture of Se and O2 in a discharge tube, and TeO3 is formed by intense dehydration of H6TeO6. Said oxides hydrolyze or react vigorously with water to form acids. Sulfuric acid is of the greatest practical importance. To obtain it, two processes are used - the constantly developing contact method and the outdated nitrous tower method (see also SULFUR).
Sulfuric acid is a strong acid; it actively interacts with water to release heat by the reaction H2SO4 + H2O H3O+ + HSO4 Therefore, care must be taken when diluting concentrated sulfuric acid, as overheating can cause vapors to escape from the acid container (burns from sulfuric acid are often associated with the addition of a small amount of water to it). Due to its high affinity for water, H2SO4 (conc.) interacts intensively with cotton clothing, sugar and human living tissues, taking away water. Enormous amounts of acid are used for the surface treatment of metals, in agriculture for the production of superphosphate (see also PHOSPHORUS), in the processing of crude oil to the rectification stage, in the technology of polymers, dyes, in the pharmaceutical industry and many other industries. Sulfuric acid is the most important inorganic compound from an industrial point of view. Oxoacids of chalcogens are given in table. 7th century It should be noted that some acids exist only in solution, others only in the form of salts.
Among the other sulfur oxo acids, an important place in industry is occupied by sulfurous acid H2SO3, which is formed when SO2 is dissolved in water, a weak acid that exists only in aqueous solutions. Its salts are quite stable. Acid and its salts are reducing agents and are used as "anti-chlorinators" to remove excess chlorine from bleach. Thiosulfuric acid and its salts are used in photography to remove excess unreacted AgBr from photographic film: AgBr + S2O32 [] + Br
The name "sodium hyposulfite" for the sodium salt of thiosulfuric acid is unfortunate, the correct name "thiosulfate" reflects the structural bond of this acid with sulfuric acid, in which one atom of unhydrated oxygen is replaced by a sulfur atom ("thio"). Polythionic acids represent an interesting class of compounds in which a chain of sulfur atoms is formed between two SO3 groups. There are many data on H2S2O6 derivatives, but polythionic acids can also contain a large number of sulfur atoms. Peroxoacids are important not only as oxidizers, but also as intermediates for the production of hydrogen peroxide. Peroxodisulfuric acid is obtained by electrolytic oxidation of the HSO4 ion in the cold. Peroxosulfuric acid is formed by the hydrolysis of peroxodisulfuric acid: 2HSO4 -> H2S2O8 + 2e
H2S2O8 + H2O -> H2SO5 + H2SO4 The range of selenium and tellurium acids is much smaller. Selenous acid H2SeO3 is obtained by evaporating water from a solution of SeO2. It is an oxidizing agent, unlike sulfurous acid H2SO3 (reducing agent) and easily oxidizes halides to halogens. The 4s2 electron pair of selenium is not actively involved in the formation of a bond (the effect of an inert pair; see above in the section on the reactivity of sulfur), and therefore selenium easily passes into the elemental state. Selenic acid, for the same reason, easily decomposes to form H2SeO3 and Se. The Te atom has a larger radius and is therefore inefficient in the formation of double bonds. Therefore, telluric acid does not exist in its usual form.
and 6 hydroxo groups are coordinated by tellurium to form H6TeO6, or Te(OH)6.
Oxohalides. Oxoacids and chalcogen oxides react with halogens and PX5 to form oxohalides of composition MOX2 and MO2X2. For example, SO2 reacts with PCl5 to form SOCl2 (thionyl chloride):
PCl5 + SO2 -> POCl3 + SOCl2
The corresponding fluoride SOF2 is formed by the interaction of SOCl2 and SbF3, and thionyl bromide SOBr2 from SOCl2 and HBr. Sulfuryl chloride SO2Cl2 is obtained by chlorination with chlorine SO2 (in the presence of camphor), sulfuryl fluoride SO2F2 is similarly obtained. Chlorofluoride SO2ClF is formed from SO2Cl2, SbF3 and SbCl3. Chlorosulfonic acid HOSO2Cl is obtained by passing chlorine through fuming sulfuric acid. Fluorosulfonic acid is formed similarly. Selenium oxohalides SeOCl2, SeOF2, SeOBr2 are also known.
Nitrogen- and sulfur-containing compounds. Sulfur forms various compounds with nitrogen, many of which are poorly understood. When S2Cl2 is treated with ammonia, N4S4 (tetrasulfur tetranitride), S7HN (heptasulfur imide), and other compounds are formed. S7HN molecules are constructed as a cyclic S8 molecule in which one sulfur atom is replaced by nitrogen. N4S4 is also formed from sulfur and ammonia. It is converted to tetrasulfur tetraimide S4N4H4 by the action of tin and hydrochloric acid. Another nitrogen derivative of sulfamic acid NH2SO3H is of industrial importance, a white, non-hygroscopic crystalline substance. It is obtained by the interaction of urea or ammonia with fuming sulfuric acid. This acid is close in strength to sulfuric acid. Its ammonium salt NH4SO3NH2 is used as a flame retardant, and alkali metal salts as herbicides.
Polonium. Despite the limited availability of polonium, the chemistry of this last VIA subgroup element has been relatively well understood through exploitation of its radioactivity property (usually mixed with tellurium as a carrier or co-reagent in chemical reactions). The half-life of the most stable isotope 210Po is only 138.7 days, so the difficulties of studying it are understandable. To obtain 1 g of Po, it is necessary to process more than 11.3 tons of uranium pitch. 210Po can be obtained by neutron bombardment of 209Bi, which first transforms into 210Bi and then ejects a b-particle, forming 210Po. Apparently, polonium exhibits the same oxidation states as other chalcogens. Polonium hydride H2Po, oxide PoO2 have been synthesized, salts with oxidation states II and IV are known. Apparently PoO3 doesn't exist.
Collier Encyclopedia. - Open Society. 2000 .
See what "CHALCOGENES" are in other dictionaries:
CHALCOGENES, chemical elements of group VI of the periodic system: oxygen, sulfur, selenium, tellurium. Compounds of chalcogens with more electropositive chemical elements chalcogenides (oxides, sulfides, selenides, tellurides) ... Modern Encyclopedia
Chemical elements of group VI of the Periodic system oxygen, sulfur, selenium, tellurium ... Big Encyclopedic Dictionary
Group → 16 ↓ Period 2 8 Oxygen ... Wikipedia
Chemical elements of group VI of the periodic system oxygen, sulfur, selenium, tellurium. * * * CHALCOGENES CHALCOGENES, chemical elements of Group VI of the Periodic Table oxygen, sulfur, selenium, tellurium ... encyclopedic Dictionary
chalcogens- chalkogenai statusas T sritis chemija apibrėžtis S, Se, Te, (Po). atitikmenys: engl. chalcogens rus. chalcogens ... Chemijos terminų aiskinamasis žodynas
Chem. elements VIa gr. periodic systems: oxygen O, sulfur S, selenium Se, tellurium Te, polonium Po. Ext. the electron shell of X atoms has the s2p4 configuration. With an increase in at. n. covalent and ionic radii X increase, energy decreases ... ... Chemical Encyclopedia
Compounds with an oxidation state of –2. H 2 Se and H 2 Te are colorless gases with a disgusting odor, soluble in water. In the series H 2 O - H 2 S - H 2 Se - H 2 Te, the stability of the molecules decreases, therefore, in aqueous solutions, H 2 Se and H 2 Te behave like dibasic acids stronger than hydrosulfide acid. They form salts - selenides and tellurides. Telluro- and hydrogen selenide, as well as their salts, are extremely toxic. Selenides and tellurides are similar in properties to sulfides. Among them are basic (K 2 Se, K 2 Te), amphoteric (Al 2 Se 3 , Al 2 Te 3) and acidic compounds (CSe 2 , CTe 2).
Na 2 Se + H 2 O NaHSe + NaOH; CSe 2 + 3H 2 O \u003d H 2 CO 3 + 2H 2 Se
A large group of selenides and tellurides are semiconductors. Selenides and tellurides of elements of the zinc subgroup are most widely used.
Compounds with an oxidation state of +4. Selenium(IV) and tellurium(IV) oxides are formed during the oxidation of simple substances with oxygen and are solid polymeric compounds. Typical acid oxides. Selenium(IV) oxide dissolves in water, forming selenous acid, which, unlike H 2 SO 3 , is isolated in a free state and is a solid.
SeO 2 + H 2 O \u003d H 2 SeO 3
Tellurium(IV) oxide is insoluble in water, but interacts with aqueous solutions of alkalis, forming tellurites.
TeO 2 + 2NaOH \u003d Na 2 TeO 3
H 2 TeO 3 is prone to polymerization, therefore, under the action of acids on tellurites, a precipitate of variable composition TeO 2 nH 2 O is formed.
SeO 2 and TeO 2 are stronger oxidizing agents compared to SO 2:
2SO 2 + SeO 2 \u003d Se + 2SO 3
Compounds with an oxidation state of +6. Selenium(VI) oxide is a white solid (mp 118.5 ºС, decomposes > 185 ºС), known in vitreous and asbestos modifications. Obtained by the action of SO 3 on selenates:
K 2 SeO 4 + SO 3 \u003d SeO 3 + K 2 SO 4
Tellurium(VI) oxide also has two modifications, orange and yellow. Obtained by dehydration of orthotelluric acid:
H 6 TeO 6 \u003d TeO 3 + 3H 2 O
Selenium(VI) and tellurium(VI) oxides are typical acidic oxides. SeO 3 dissolves in water forming selenic acid - H 2 SeO 4 . Selenic acid is a white crystalline substance, in aqueous solutions it is a strong acid (K 1 \u003d 1 10 3, K 2 \u003d 1.2 10 -2), carbonizes organic compounds, a strong oxidizing agent.
H 2 Se +6 O 4 + 2HCl -1 = H 2 Se +4 O 3 + Cl 2 0 + H 2 O
Salts - barium and lead selenates are insoluble in water.
TeO 3 is practically insoluble in water, but interacts with aqueous solutions of alkalis, forming salts of telluric acid - tellurates.
TeO 3 + 2NaOH \u003d Na 2 TeO 4 + H 2 O
Under the action of hydrochloric acid solutions of tellurates, orthotelluric acid is released - H 6 TeO 6 - a white crystalline substance that is highly soluble in hot water. Dehydration of H 6 TeO 6 can produce telluric acid. Telluric acid is very weak, K 1 \u003d 2 10 -8, K 2 \u003d 5 10 -11.
Na 2 TeO 4 + 2HCl + 2H 2 O \u003d H 6 TeO 6 + 2NaCl; H 6 TeO 6 ¾® H 2 TeO 4 + 2H 2 O.
Selenium compounds are toxic to plants and animals, while tellurium compounds are much less toxic. Poisoning with compounds of selenium and tellurium is accompanied by the appearance of a persistent disgusting smell in the victim.
Literature: p. 359 - 383, p. 425 - 435, p. 297 - 328
