The structure of selenium and tellurium oxide molecules. Abstract: Oxygen chemical characteristics

Slide 12

Chemical properties

In air, sulfur burns, forming sulfur dioxide - a colorless gas with a pungent odor: S + O2 → SO2 The reducing properties of sulfur are manifested in the reactions of sulfur with other non-metals, but at room temperature sulfur reacts only with fluorine: S + 3F2 → SF6

Slide 13

Molten sulfur reacts with chlorine, with the possible formation of two lower chlorides 2S + Cl2 → S2Cl2 S + Cl2 → SCl2 When heated, sulfur also reacts with phosphorus, forming a mixture of phosphorus sulfides, among which is the higher sulfide P2S5: 5S + 2P → P2S2 In addition , when heated, sulfur reacts with hydrogen, carbon, silicon: S + H2 → H2S (hydrogen sulfide) C + 2S → CS2 (carbon disulfide)

Slide 14

Of the complex substances, we should first of all note the reaction of sulfur with molten alkali, in which sulfur is disproportioned similarly to chlorine: 3S + 6KOH → K2SO3 + 2K2S + 3H2O With concentrated oxidizing acids, sulfur reacts only with prolonged heating: S+ 6HNO3 (conc) → H2SO4 + 6NO2 + 2H2O S+ 2 H2SO4 (conc) → 3SO2 + 2H2O

Slide 15

At 100–160°C it is oxidized by water: Te+2H2O= TeO2+2H2 When boiled in alkaline solutions, tellurium disproportionates to form telluride and tellurite: 8Te+6KOH=2K2Te+ K2TeO3+3H2O.

Slide 16

Dilute HNO3 oxidizes Te to telluric acid H2TeO3: 3Te+4HNO3+H2O=3H2TeO3+4NO. Strong oxidizing agents (HClO3, KMnO4) oxidize Te to weak telluric acid H6TeO6: Te+HClO3+3H2O=HCl+H6TeO6. Tellurium compounds (+2) are unstable and prone to disproportionation: 2TeCl2=TeCl4+Te.

Slide 17

When heated in air, it burns to form colorless crystalline SeO2: Se + O2 = SeO2. Reacts with water when heated: 3Se + 3H2O = 2H2Se + H2SeO3. Selenium reacts when heated with nitric acid to form selenous acid H2SeO3: 3Se + 4HNO3 + H2O = 3H2SeO3 + 4NO.

Slide 18

When boiled in alkaline solutions, selenium disproportionates: 3Se + 6KOH = K2SeO3 + 2K2Se + 3H2O. If selenium is boiled in an alkaline solution through which air or oxygen is passed, then red-brown solutions containing polyselenides are formed: K2Se + 3Se = K2Se4

In the VIA group of the periodic system of elements D.I. Mendeleev's elements include oxygen, sulfur, selenium, tellurium, and polonium. The first four of them are non-metallic in nature. General name of the elements of this group chalcogens, which is translated from Greek. means “forming ores,” indicating their occurrence in nature.

Electronic formula of the valence shell of atoms of group VI elements.

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 in 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 outer energy level allows the oxygen atom to form additional bonds through 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, atoms of sulfur and other chalcogens are capable of forming 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, 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 F2O, 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 nonmetallic properties of chalcogens during the transition from oxygen to tellurium.

As the atomic number of chalcogens increases, the characteristic coordination numbers increase. This is due to the fact that during the transition from p-elements of the fourth period to p-elements of the fifth and sixth periods in the formation of σ- and π-bonds d begin to play an increasingly important role - and even f-orbitals. So, if for sulfur and selenium the most characteristic coordination numbers are 3 and 4, then for tellurium - 6 and even 8.

Under normal conditions, hydrogen compounds H 2 E of group VIA elements, 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 explained by an increase in the radii of the E 2- ions and a corresponding weakening of the E-H 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 in these compounds have an intermediate oxidation state of + 4.

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

Chemical properties of oxygen compounds. Oxygen is the most abundant 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 occurs due to 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, excess oxygen (4/5) is necessary to create a high oxygen concentration in the blood. This, in accordance with Fick's law, ensures a sufficient rate of oxygen diffusion through the walls of the capillaries. 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 residues. 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 ozone content is very small. The maximum volume fraction of ozone at an altitude of 22 km is only 10 -6%.

The oxygen molecule O2 is very stable in the absence of other substances. The presence of two unpaired electrons in the molecule determines its high reactivity. Oxygen is one of the most active non-metals. It reacts directly with most simple substances, forming oxides E x O y. The oxidation state of oxygen in them is -2. In accordance with the change in the structure of the electronic 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 naturally. Thus, in the 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 V gradually decreases. In accordance with this, the basic properties are weakened and the acidic 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 acidic oxides. Acid-base properties change similarly in other periods.

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

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

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

The polarity of the E-OH bond in E(OH) n hydroxides, and therefore 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 increasing ion radii, the polarity of the E-OH bond increases. As a result, ionization E-OH → E + + OH - occurs more easily in water. Accordingly, the basic properties of hydroxides are enhanced. Thus, 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, ionization of EON ⇄ EO - + H + occurs more easily in water. Accordingly, acidic properties are enhanced in this direction. Thus, 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 creature. The human body contains 70-80% water. Over the course of 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 and 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 is vital oxygenyl complexes with various substances. Previously, examples of O 2 oxygenyl complexes with metal ions - oxygen carriers in living organisms - oxyhemoglobin and oxyhemocyanin were considered:

НbFe 2 + + О 2 → НbFe 2+ ∙О 2

НсСu 2+ + О 2 → НсСu 2+ ∙О 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. To a certain extent, these compounds can be considered as substitutes for hemoglobin.

One of the allotropic modifications of oxygen is ozone O 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:

3О 2 ⇄2О 3 ,

Ozone is produced by the action of an electrical discharge in oxygen. Ozone is formed from O 2 and under the influence 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, and at low temperatures oxidizes compounds with which oxygen does not react:

O 3 + 2Ag = Ag 2 O + O 2

РbS + 4О 3 = РbSO 4 + 4O 2

A well-known qualitative reaction is:

2KI + O 3 + H 2 O = I 2 + 2KON + O 2

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

RН + О 3 → RО 2 ∙ + HE ∙

Radicals initiate radical chain reactions with bioorganic molecules - lipids, proteins, DNA. Such reactions lead to cell damage and death. In particular, ozone kills microorganisms contained 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. In its properties, sulfur is close to oxygen. But unlike it, it exhibits in compounds not only the oxidation state -2, but also positive oxidation states +2, +4 and +6. Sulfur, like oxygen, is characterized by allotropy - the existence of several elemental substances - orthorhombic, monoclinic, plastic sulfur. Due to its 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 a significant biological role in life processes. Thus, 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. Thanks to the S-S disulfide bond, the polypeptide chains are held together (disulfide bridge).

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

The amino acid methionine is very important biologically.

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 biosynthesis processes is transferred to various acceptors of methyl groups RH:

Ad-S-CH 3 + RN → Ad-SN + R-CH 3

Sulfur is quite widespread on Earth (0.03%). It is present in nature 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 the native state. Precipitated sulfur powder is used externally in the form of ointments (5-10-20%) and powders in the treatment of skin diseases (seborrhea, psoriasis). The body produces sulfur oxidation products - 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 reactive non-metal. Even with slight heating, it oxidizes many simple substances, but it itself is easily oxidized by oxygen and halogens (redox duality).

Sulfur exhibits oxidation state -2 in hydrogen sulfide and its derivatives - sulfides.

Hydrogen sulfide (dihydrogen sulfide) often found in nature. Contained in so-called sulfur 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 influence of microorganisms. Some photosynthetic bacteria, such as green sulfur bacteria, use dihydrogen sulfide as a hydrogen donor. These bacteria, instead of oxygen O2, produce elemental sulfur - a product of the oxidation of H2S.

Dihydrogen sulfide is a very toxic substance, as it is an inhibitor of the enzyme cytochrome oxidase, an electron transporter in the respiratory chain. It blocks the transfer of electrons from cytochrome oxidase to oxygen O2.

Aqueous solutions of H 2 S give a slightly acidic reaction to litmus. Ionization occurs in two stages:

Н 2 S ⇄ Н + + НS - (I stage)

NS - ⇄ N + + S 2- (II stage)

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

Salts of hydrosulfide 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 = 2SO 2 + 2H 2 O (in air)

2H 2 S + O 2 = 2H 2 O + 2S (in air)

3H 2 S + 4HClO 3 = 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-SН + Ag + → E-S-Аg + H +

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

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

S + O 2 = SO 2

4FеS 2 + 11О 2 = 2Fe 2 О 3 + 8SO 2

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

Oxide SO2, acid H2SO3 and its salts are characterized by redox duality, since sulfur in these compounds has an intermediate oxidation state of +4:

2Na 2 SO 3 + O 2 = 2Na 2 SO 4

SO 2 + 2H 2 S = 3S° + 2H 2 O

However, the reducing properties of sulfur (IV) compounds predominate. Thus, sulfites in solutions are oxidized even by dioxygen in the air 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 small amounts 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 = H 2 SO 3

2H 2 SO 3 + O 2 = 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 heating a solution of Na 2 SO 3 with sulfur powder, sodium thiosulfate:

Na 2 SO 3 + S = Na 2 S 2 O 3

Crystalline hydrate Na 2 S 2 O 3 ∙5H 2 O is released 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 compounds, 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 entering the body with food or air form poorly soluble non-toxic sulfites:

Рb 2+ + SO 3 2- = РbSO 3

Cyanide ions react with elemental sulfur, forming less toxic thiocyanates:

СN - + S° = NСS -

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

Na 2 S 2 O 3 + 2HCl = 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 = 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 occurs to a lesser extent:

НSO 4 - ⇄ Н + + SO 4 2- ,

Concentrated sulfuric acid is a strong oxidizing agent. It oxidizes metals and non-metals. Typically, 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 - hydrosulfates(NaHSO 4, KHSO 4, etc.). Most sulfates are highly soluble in water. Many sulfates are isolated from solutions in the form of crystalline hydrates: FeSO 4 ∙7H 2 O, CuSO 4 ∙5H 2 O. The practically insoluble sulfates include BaSO 4, SrSO 4 and PbSO 4. 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, emetic, barium sulfate BaSO 4 - as a contrast agent for x-ray examination of the esophagus and stomach

Compounds of selenium and tellurium. The chemical properties of tellurium and especially selenium are similar to sulfur. However, strengthening the metallic properties of Se and Te increases their tendency to form stronger ionic bonds. The similarity of physicochemical characteristics: radii of E 2- ions, coordination numbers (3, 4) - determines the interchangeability of selenium and sulfur in compounds. Thus, 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 similar to H 2 S, very poisonous gases 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 are isolated as medium salts - selenides(Na 2 Se, etc.), and acid salts - hydroselenides(NaHSe, etc.). For H 2 Te, only medium salts are known - tellurides.

Compounds of Se (IV) and Te (IV) with oxygen, unlike SO 2, are solid crystalline substances SeO 2 and TeO 2.

Selenous acid H 2 SeO 3 and its salts Selenites, for example, Na 2 SeO 3 are oxidizing agents of medium strength. Thus, 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 = Se + 2H 2 SO 4

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

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

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

Compounds with oxidation state –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 as dibasic acids stronger than hydrosulfide acid. They form salts - selenides and tellurides. Tellurium and hydrogen selenide, as well as their salts, are extremely toxic. Selenide and tellurides have properties similar to sulfides. Among them there 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 = H 2 CO 3 + 2H 2 Se

A large group of selenides and tellurides are semiconductors. The most widely used are selenides and tellurides of elements of the zinc subgroup.

Compounds with oxidation state +4. Selenium(IV) and tellurium(IV) oxides are formed by the oxidation of simple substances with oxygen and are solid polymer 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 = H 2 SeO 3

Tellurium(IV) oxide is insoluble in water, but reacts with aqueous solutions of alkalis, forming tellurites.

TeO 2 + 2NaOH = Na 2 TeO 3

H 2 TeO 3 is prone to polymerization, therefore, when acids act on tellurites, a precipitate of variable composition TeO 2 nH 2 O is released.

SeO 2 and TeO 2 are stronger oxidizing agents compared to SO 2:

2SO 2 + SeO 2 = Se + 2SO 3

Compounds with oxidation state +6. Selenium(VI) oxide is a white solid (mp 118.5 ºС, decomposes > 185 ºС), known in glassy and asbestos-like modifications. Obtained by the action of SO 3 on selenates:

K 2 SeO 4 + SO 3 = SeO 3 + K 2 SO 4

Tellurium(VI) oxide also has two modifications: orange and yellow. Prepared by dehydration of orthotelluric acid:

H 6 TeO 6 = TeO 3 + 3H 2 O

Oxides of selenium(VI) and tellurium(VI) 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 = 1·10 3, K 2 = 1.2·10 -2), chars 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 telluric acid salts - tellurates.

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

When solutions of tellurates are exposed to hydrochloric acid, orthotelluric acid is released - H 6 TeO 6 - a white crystalline substance that is highly soluble in hot water. By dehydrating H 6 TeO 6, telluric acid can be obtained. Telluric acid is very weak, K1 = 2·10 -8, K2 = 5·10 -11.

Na 2 TeO 4 + 2HCl + 2H 2 O = 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; tellurium compounds are much less toxic. Poisoning with selenium and tellurium compounds is accompanied by the appearance of a persistent disgusting odor in the victim.

Literature: p. 359 - 383, p. 425 - 435, p. 297 - 328