Chemistry of the Elements Non-metals of VIA-subgroup

Elements of the VIA subgroup are non-metals, except for Po.

Oxygen is very different from other subgroup elements and plays a special role in chemistry. Therefore, the chemistry of oxygen is highlighted in a separate lecture.

Sulfur is the most important among the other elements. The chemistry of sulfur is very extensive, since sulfur forms a huge variety of compounds. Its compounds are widely used in chemical practice and in various industries. When discussing nonmetals of the VIA subgroup, the greatest attention will be paid to the chemistry of sulfur.

Key Issues Addressed in the Lecture

General characteristics of non-metals of the VIA-subgroup. Natural compounds Sulfur

Simple substance Sulfur compounds

Hydrogen sulfide, sulfides, polysulfides

Sulphur dioxide. Sulfites

Sulfur trioxide

Sulphuric acid. oxidative properties. sulfates

Other sulfur compounds

selenium, tellurium

Simple substances Compounds of selenium and tellurium

Selenides and tellurides

Se and Te compounds in oxidation state (+4)

Selenic and telluric acids. oxidative properties.

				Elements of the VIA subgroup
					general characteristics
						The p-elements belong to the VIA subgroup: acid-
					genus O, sulfur S, selenium Se, tellurium Te, polonium Po.
									The general formula for valence electrons
									thrones - ns 2 np 4 .
			oxygen
						Oxygen, sulfur, selenium and tellurium are non-metals.
					They are often grouped under the common name "chalcogens",
					which means "forming ores". Indeed many
					metals are found in nature in the form of oxides and sulfides;
					in sulfide ores				in small quantities with
					there are selenides and tellurides.
						Polonium is a very rare radioactive element that
					which is a metal.
			molybdenum
						To create a stable eight-electron
					chalcogen atoms lack only two electro-
					new The minimum oxidation state (–2) is
			tungsten		resistant to all elements. It is this degree of oxidation
					elements show in natural compounds - ok-
					sides, sulfides, selenides and tellurides.
						All elements of the VIA-subgroup, except for O, exhibit
			seaborgium		positive oxidation states +6 and +4. Most-
					the highest oxidation state of oxygen is +2, it exhibits
	only in conjunction with F.

The most characteristic oxidation states for S, Se, Te are

xia: (–2), 0, +4, +6, for oxygen: (–2), (–1), 0.

In the transition from S to Te, the stability of the highest oxidation state is +6

decreases, and the stability of the +4 oxidation state increases.

For Se, Te, Po, - the most stable oxidation state is +4.

Some characteristics of atoms of elements ViB - subgroups

		Relative	First energy
		elektrootri-	ionization,
		value	kJ/mol
		(according to Polling)
				an increase in the number of
				throne layers;
				an increase in the size of an atom;
				decrease in energy io-
				decrease in electrical
				values

As can be seen from the above data , oxygen is very different from other elements of the subgroup high value of ionization energy, ma-

large orbital radius of the atom and high electronegativity, only F has a higher electronegativity.

Oxygen, which plays a very special role in chemistry, was considered from

sensibly. Among the other elements of the VIA group, sulfur is the most important.

			Sulfur forms a very large number of various
			different connections. Its compounds are known from almost all
			mi elements, except for Au, Pt, I and noble gases. Cro-
			me of widespread compounds S in powers
	3s2 3p4
			oxidation (–2), +4, +6, are known, as a rule,
			stable compounds in oxidation states: +1 (S2 O), +2

(SF2 , SCl2 ), +3 (S2 O3 , H2 S2 O4 ). The variety of sulfur compounds is also confirmed by the fact that only about 20 oxygen-containing acids S are known.

The strength of the bond between S atoms turns out to be commensurate with the

bonds S with other non-metals: O, H, Cl, therefore, S is characterized by

including the very common mineral pyrite, FeS2, and polythionic acids (eg H2 S4 O6 ). Thus, the chemistry of sulfur is quite extensive.

The most important sulfur compounds used in industry

The most widely used sulfur compound in industry and the laboratory is sulfuric acid. The world volume of production of ser-

acid is 136 million tons. (no other acid is produced in such large quantities). Common compounds include

whether sulfuric acid - sulfates, as well as salts of sulfurous acid - sulfites.

natural sulfides are used to obtain the most important non-ferrous metals

thalls: Cu, Zn, Pb, Ni, Co, etc. Other common sulfur compounds include: hydrosulfide acid H2 S, di- and trioxides of sulfur: SO2

and SO3, thiosulfate Na2 S2 O3 ; acids: disulfuric (pyrosulfuric) H2 S2 O7, perox-

codisulfate H2 S2 O8 and peroxodisulfates (persulphates): Na2 S2 O8 and

(NH4 )2 S2 O8 .

Sulfur in nature

tea in the form of a simple substance, forming large underground deposits,

and in the form of sulfide and sulfate minerals , as well as in the form of compounds,

which are impurities in coal and oil. Coal and oil are obtained as a result of

those decompositions of organic substances, and sulfur is a part of animals and plants

body proteins. Therefore, when coal and oil are burned, sulfur oxides are formed,

polluting the environment.

Natural sulfur compounds

Rice. Pyrite FeS2 is the main mineral used to produce sulfuric acid.

native sulfur;

sulfide minerals:

FeS2 - pyrite or iron pyrite

FeCuS2 - chalcopyrite (copper quanti-

FeAsS - arsenopyrite

PbS - galena or lead luster

ZnS - sphalerite or zinc blende

HgS - cinnabar

Cu2 S- chalcocite or copper luster

Ag2 S - argentite or silver sheen

MoS2 - molybdenite

Sb2 S3 - stibnite or antimony shine

As4 S4 - realgar;

sulfates:

Na2 SO4 . 10 H2 O - mirabilite

CaSO4 . 2H2 O - gypsum

CaSO4 - anhydrite

BaSObarite or heavy spar

SrSO4 is celestine.

Rice. Gypsum CaSO4. 2H2O

simple substance

In a simple substance, sulfur atoms are bonded with two neighboring ones.

The most stable is the structure consisting of eight sulfur atoms,

united in a corrugated ring resembling a crown. There are several modifications of sulfur: rhombic sulfur, monoclinic and plastic sulfur. At ordinary temperature, sulfur is in the form of yellow brittle crystals.

rhombic shape (-S), formed by

ionic molecules S8 . Another modification - monoclinic sulfur (-S) also consists of eight-membered rings, but differs in location

arrangement of S8 molecules in the crystal. When dis-

melting sulfur rings are torn. At the same time, mo-

tangled threads can form, which

Rice. Sulfur

make the melt viscous, with further

As the temperature rises, the polymer chains can break down and the viscosity will decrease. Plastic sulfur is formed during the sharp cooling of the molten

sulfur and consists of entangled chains. Over time (within a few days), it will be converted into rhombic sulfur.

Sulfur boils at 445o C. Equilibria take place in sulfur vapor:

450 o C	650 o C	900 o C	1500 o C
S 8  S 6	 S 4	 S 2	 S

S2 molecules have a structure similar to O2.

Sulfur can be oxidized (usually to SO2) and can be reduced

upgraded to S(-2). At ordinary temperatures, almost all reactions involving solid sulfur are inhibited; only reactions with fluorine, chlorine, and mercury proceed.

This reaction is used to bind the smallest droplets of spilled mercury.

Liquid and vaporous sulfur are highly reactive . Sulfur vapor burns Zn, Fe, Cu. When passing H 2 over molten sulfur is formed

H 2 S. In reactions with hydrogen and metals, sulfur acts as an oxidizing

Sulfur can be easily oxidized under the action of halogens.

and oxygen. When heated in air, sulfur burns with a blue flame, oxidizing

up to SO2.

S + O2 = SO2

Sulfur is oxidized with concentrated sulfuric and nitric acids:

S + 2H2 SO4 (conc.) = 3SO2 + 2H2 O,

S + 6HNO3 (conc.) = H2 SO4 + 6 NO2 + 2H2 O

In hot alkali solutions, sulfur disproportionates.

3S + 6 NaOH = 2 Na2 S + Na2 SO3 + 3 H2 O.

When sulfur reacts with a solution of ammonium sulfide, yellow-red polysulfide ions(–S–S–)n or Sn 2– .

When sulfur is heated with a solution of sulfite, thiosulfate is obtained, and

when heated with a solution of cyanide - thiocyanate:

S + Na 2 SO3 = Na2 S2 O3, S + KCN = KSCN

Potassium thiocyanate or thiocyanate is used for the analytical detection of Fe3+ ions:

3+ + SCN – = 2+ + H2O

The resulting complex compound has a blood-red color,

even at a low concentration of hydrated Fe3+ ions in the

About 33 million tons of native sulfur are mined annually in the world. The main amount of extracted sulfur is processed into sulfuric acid and used

used in the rubber industry for the vulcanization of rubber. Add sulfur

binds to double bonds of rubber macromolecules, forming disulfide bridges

ki -S- S-, thereby, as if "stitching" them, which gives the rubber strength and elasticity. When a large amount of sulfur is introduced into rubber, ebo-

nit, which is a good insulating material used in electrical engineering. Sulfur is also used in pharmaceuticals to make skin ointments and in agriculture to control plant pests.

Sulfur compounds

Hydrogen sulfide, sulfides, polysulfides

Hydrogen sulfide H 2 S occurs naturally in sulfuric mineral waters,

present in volcanic and natural gas, formed during the decay of white

kov bodies.

Hydrogen sulfide is a colorless gas with a rotten egg odor and is highly toxic.

It is slightly soluble in water; at room temperature, three volumes of gaseous H2 S dissolve in one volume of water. The concentration of H 2 S in saturated

nom solution is ~ 0.1 mol/l . When dissolved in water, it forms

hydrosulfide acid, which is one of the weakest acids:

		H2 S  H+ + HS – , K1 = 6. 10 –8 ,
		HS -  H+ + S 2–,								K2 = 1.10 –14
Executor:											Many natural sulfides are known (see the list of sulfide minerals). Sulfides of many heavy non-ferrous metals (Cu, Zn, Pb, Ni, Co, Cd, Mo) are are industrially important ores. They are converted into oxides by firing in air, for example, 2 ZnS + 3 O2 = 2 ZnO + 2 SO2 then the oxides are most often reduced with coal: ZnO + C = Zn + CO Sometimes oxides are brought into solution by the action of an acid, and then the solution is subjected to electrolysis in order to reduce the metal. Sulfides of alkali and alkaline earth metals are practically chemically ionic compounds. Sulfides of other metals - the advantage vein-covalent compounds, as a rule, of non-stoichiometric composition. Many nonmetals also form covalent sulfides: B, C, Si, Ge, P, As, Sb. Natural sulfides As and Sb are known. Sulfides of alkali and alkaline earth metals, as well as sulfides ammonium feed is highly soluble in water, the rest of the sulfides are insoluble rhymes. They are isolated from solutions in the form of characteristically colored precipitates, for example, Pb(NO3 )2 + Na2 S = PbS (t.) + 2 NaNO3 This reaction is used to detect H2 S and S2– in solution. Some of the water-insoluble sulfides can be brought into solution by acids, due to the formation of a very weak and volatile hydrosulphuric acid. native acid, for example, NiS + H2SO4 = H2S + NiSO4 Sulfides can be dissolved in acids: FeS, NiS, CoS, MnS, ZnS. Metal sulfides and PR values Sulfides Sediment color PR value 5 . 10–18 1 . 10–24 2 . 10–25 2 . 10–27 6 . 10–36 4 . 10–53 brown 2 . 10–27 2 . 10–28 2 . 10–10 2 . 10–24 Sulfides, characterized by a very low value of the solubility product, cannot dissolve in acids with the formation of H2 S. In ki- sulfides do not dissolve in slots: CuS, PbS, Ag2 S, HgS, SnS, Bi2 S3, Sb2 S3, Sb2 S5, CdS, As2 S3, As2 S5, SnS2. If the reaction of dissolution of sulfide due to the formation of H2 S is impossible, then it can be transferred into a solution by the action of concentrated nitric acid slots or aqua regia. CuS + 8HNO3 = CuSO4 + 8NO2 + 4H2O The sulfide anion S 2– is a strong proton acceptor (os- innovation according to Brønsted). That's why highly soluble sulfides

The oxygen subgroup includes five elements: oxygen, sulfur, selenium, tellurium and polonium (a radioactive metal). These are the p-elements of the VI group of the periodic system of D.I. Mendeleev. They have a group name - chalcogens, which means "forming ores."

Properties of elements of the oxygen subgroup

Properties				Those	Ro
1. Order number
2. Valence electrons	2 s 2 2p 4	Z s 2 3p 4	4 s 2 4r 4	5s 2 5p 4	6s 2 6p 4
3. Energy Ionization of atom, eV	13,62	10,36	9,75	9,01	8,43
4. Relative electronegativity	3,50		2,48	2,01	1,76
5. The oxidation state in connections	1, -2,	2, +2, +4, +6	4, +6	4, +6	2, +2
6. Atomic radius, nm	0,066	0,104	0,117 0,137		0,164

Chalcogen atoms have the same structure of the external energy level - ns 2 nr 4 . This explains the similarity of their chemical properties. All chalcogens in compounds with hydrogen and metals exhibit an oxidation state of -2, and in compounds with oxygen and other active non-metals, usually +4 and +6. For oxygen, as well as for fluorine, an oxidation state equal to the group number is not typical. It exhibits an oxidation state of usually -2 and in combination with fluorine +2. Such values ​​of oxidation states follow from the electronic structure of chalcogens

The oxygen atom has two unpaired electrons in the 2p sublevel. Its electrons cannot be separated, since there is no d-sublevel at the outer (second) level, that is, there are no free orbitals. Therefore, the valency of oxygen is always equal to two, and the oxidation state is -2 and +2 (for example, in H 2 O and OF 2). These are the same valencies and oxidation states of the atom of sulfur in the unexcited state. Upon transition to an excited state (which takes place during the supply of energy, for example, during heating), at the sulfur atom, the 3 R— and then 3s electrons (shown by arrows). The number of unpaired electrons, and, consequently, the valency in the first case is four (for example, in SO 2), and in the second - six (for example, in SO 3). Obviously, even valencies 2, 4, 6 are characteristic of sulfur analogues - selenium, tellurium and polonium, and their oxidation states can be equal to -2, +2, +4 and +6.

Hydrogen compounds of elements of the oxygen subgroup are responsible formula H 2 R (R - element symbol): H 2 O, H 2 S , H 2 S e, H 2 Te. They callare hydrogen chalcides. When dissolved in water, they formacids. The strength of these acids increases with increasing atomic number of the element, which is explained by a decrease in energy bonds in the series of compounds H 2 R . Water dissociating into H + and O ions H - , is amphoteric electrolyte.

Sulfur, selenium and tellurium form the same forms of compounds with oxygen of the type R O 2 and R About 3- . They correspond to acids of the type H 2 R O 3 and H 2 R About 4- . With an increase in the ordinal number of the element, the strength of these acids decreases.vaet. All of them exhibit oxidizing properties, and acids of the type H 2 R About 3 are also restorative.

The properties of simple substances naturally change: with an increase incharge of the nucleus, non-metallic ones weaken and metallic ones increase. properties. So, oxygen and tellurium are non-metals, but the latter hasmetallic luster and conducts electricity.

Selenium is not widely distributed in nature. The content of selenium in the earth's crust is . Its compounds are found as impurities in natural sulfur compounds with metals and. Therefore, selenium is obtained from waste products generated in the production of sulfuric acid, in the electrolytic refining of copper, and in some other processes.

Tellurium is one of the rare elements: its content in the earth's crust is only .

In the free state, selenium, like sulfur, forms several allotropic modifications, of which the most famous are amorphous selenium, which is a red-brown powder, and gray selenium, which forms brittle crystals with a metallic luster.

Tellurium is also known in the form of an amorphous modification and in the form of light gray crystals with a metallic luster.

Selenium is a typical semiconductor (see § 190). An important property of it as a semiconductor is a sharp increase in electrical conductivity when illuminated. At the boundary of selenium with a metal conductor, a barrier layer is formed - a section of the circuit that can pass electric current in only one direction. In connection with these properties, selenium is used in semiconductor technology for the manufacture of rectifiers and photocells with a barrier layer. Tellurium is also a semiconductor, but its use is more limited. Selenides and tellurides of some metals also have semiconductor properties and are used in electronics. In small amounts, tellurium serves as an alloying addition to lead, improving its mechanical properties.

Hydrogen selenide and hydrogen telluride are colorless gases with a disgusting odor. Their aqueous solutions are acids, the dissociation constants of which are somewhat larger than the dissociation constant of hydrogen sulfide.

Chemically, hydrogen selenide and hydrogen telluride are extremely similar to hydrogen sulfide. Like hydrogen sulfide, they are highly reducing properties. When heated, they both decompose. At the same time, it is less stable than: just as it happens in the series of hydrogen halides, the strength of the molecules decreases during the transition. Salts of hydrogen selenide and hydrogen telluride - selenides and tellurides - are similar to sulfides in terms of solubility in water and acids. By acting on selenides and tellurides with strong acids, hydrogen selenide and hydrogen telluride can be obtained.

When selenium and tellurium are burned in air or in oxygen, dioxides and are obtained, which under normal conditions are in a solid state and are anhydrides of selenous and tellurous acids.

Unlike sulfur dioxide, and exhibit predominantly oxidizing properties, easily recovering to free selenium and tellurium, for example:

By the action of strong oxidizing agents, selenium and tellurium dioxides can be converted into selenic and telluric acids, respectively.

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 l 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 purifying (disinfecting) water, bleaching fabrics, starch, refining oils, 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 Oxidation state I for sulfur, it may be realized in the compound (SCl)2 or S2Cl2 (this composition has not been 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 with the release of heat by the reaction H2SO4 + H2O H3O+ + HSO4 Therefore, care should be taken when diluting concentrated sulfuric acid, as overheating can cause the release of vapors from the acid tank (burns from sulfuric acid are often associated with the addition of a small amount of sulfuric acid to it water). 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