home / Success stories/ The elements of group 6 of the main subgroup are called. Elements of group vi b

The elements of group 6 of the main subgroup are called. Elements of group vi b

Lecture 2

SUBJECT : GROUP ELEMENTS VI B

Questions covered in the lecture:

general characteristics d elements of group VI.
Finding in nature and obtaining chromium, molybdenum and tungsten.
Physical properties of metals.
Chemical properties of chromium, molybdenum and tungsten.
The most important compounds of elements of the chromium subgroup: a) compounds

E (P); b) connections E (SH); c) connections E ( VI).

Chromium peroxide.

Side subgroup VI group is represented by the following elements: C r, Mo and W . All of them are d -elements, since they are built up with electrons d -sublevel of the pre-external level. The valence electrons of these elements are the electrons of the outer S -sublevel and pre-external d -sublevel - only 6 electrons.

Electronic configuration of outer layer and pre-external d-sublevel: C r 3 d 5 4 S 1; Mo 4 d 5 5 S 1 ; W 5 d 4 6 S 2 .

d Group 6 elements occupy 4th place in their decade d elements, therefore d The sublevel must contain 4 electrons, and the outer level must contain two s electron, as is observed for tungsten. For chromium and molybdenum, there is a “breakthrough” of one s electron from the outer level to the pre-external one d sublevel, resulting in each d the orbital will be occupied by one electron, which corresponds to the most stable state of the atom.

│││││ │ (n 1) d → ││││││ (n 1) d

nS │↓│ nS ││

Table 3

Basic parameters of element atoms VI B groups

	Atomic radius rа, nm	Ion radius r E 6+, nm	E E o → E + , eV	Ar	chemical activity
With r	0,127	0,035	6,76		│reduce- │worried
	0,137	0,065	7,10
	0,140	0,065	7,98

Analyzing these data, we can say that there is a common d -elements regularity: the radii of atoms from top to bottom in the subgroup increase, but only slightly. Since the mass of atoms in the same row increases greatly, this leads to densification of the electron shells of molybdenum and especially tungsten. It is more difficult to remove an electron from such a compacted structure, therefore the ionization energy during the transition from chromium to tungsten increases, as a result of which the chemical activity of elements from top to bottom in the subgroup decreases. Due to the fact that molybdenum and tungsten have approximately the same atomic and ionic radii, their properties are closer to each other than to chromium.

In compounds, chromium and its analogues exhibit oxidation states (S.O.) of 0, +1, +2, +3, +4, +5 and +6. Maximum S.O. corresponds to the number of valence electrons. Characteristic S.O. chromium +3 and to a lesser extent +6 and +2. Molybdenum and tungsten, like the other 4 d - and 5 d - elements, the highest S.O. is most characteristic, that is, +6. Thus, for elements of the subgroup Cr common for d elements pattern: increase in the group from top to bottom of stable S.O. Therefore, the oxidizing ability of compounds where elements exhibit a higher CO equal to +6 decreases from top to bottom in the subgroup, since the stability of compounds in this series increases. For example, in the acid series:

H 2 C rO 4 │ stability Cr +6 │ oxidizing capacity

H 2 MoO 4 │Mo increases+6 │ decreases

H 2 W O 4 │ W +6 │

↓ ↓

For Cr, Mo, W the most typical coordination numbers are 6 and 4. Derivatives are also known in which the c.n. My W reaches 8.

Examples: [Cr (OH) 4] - [Cr (H 2 O) 6] 3+

3- 2-

At the same time, they can participate in the formation of connections d -orbitals of the pre-external level, as well as S - and p-orbitals of the outer level.

The nature of the connection between elements of subgroup C r in compounds is determined largely by S.O. element. For Cr, Mo, W at low S.O. (+1, +2) ionic bonds are characteristic, and at high S.O. covalent bonds. In accordance with this C r +2 O basic oxide, C r 2 +3 O 3 amphoteric, and C r +6 O 3 acidic. Same as C r (OH) 2 base, C r (OH) 3 amphoteric hydroxide, H 2 C r O 4 acid.

Finding in nature and obtaining chromium, molybdenum and tungsten

The chromium content in the earth's crust is 0.02% (mass), molybdenum 10-3 % (mass), tungsten 7 ∙ 10-3 % (mass). The main ore of chromium is chromium iron ore Fe(CrO2)2 (chromite). Molybdenum occurs as the mineral molybdenite Mo S 2 (molybdenum luster), as well as molybdates: PvMoO 4 (wulfenite) and M gMoO 4 . The most important tungsten ores wolframite (mixture FeWO 4 and M nWO 4 ), scheelite Ca WO 4 and stolcite Pb WO 4 .

To obtain pure chromium, the oxide is first prepared Cr2O3 , which is then reduced aluminothermically:

Cr 2 O 3 + 2Al → Al 2 O 3 + 2Cr.

For metallurgical purposes, chromium is obtained in the form of an alloy with iron (ferrochrome). To do this, chromium iron ore is reduced with coal in an electric furnace.

Fe(CrO 2 ) 2 + 4C → Fe + 2Cr + 4CO.

Molybdenum and tungsten are obtained by converting the above minerals into oxides, from which the metal is reduced with hydrogen at high temperatures:

2 Mo S 2 + 7O 2 → 2MoO 3 + 4SO 2

MoO 3 + 3H 2 → Mo + 3H 2 O.

Physical properties of metals

In the form of simple substances, chromium, molybdenum and tungsten are grayish-white shiny metals. They are all refractory, and tungsten is the most refractory of the metals (T pl. = 3380 o C).

The electrical conductivity of metals in the transition from chromium to tungsten generally increases and for molybdenum and tungsten is approximately 30% of the electrical conductivity of silver. The properties of metals are greatly influenced by impurities. Thus, technical chromium is one of the hardest metals, while pure chromium is ductile.

Chemical properties of chromium, molybdenum and tungsten

Chemical activity in a series CrMoW decreases noticeably. Under normal conditions, all three metals noticeably interact only with fluorine:

Me + 3F 2 → MeF 6 (CrF 3).

Under normal conditions, these metals are resistant to air oxygen and water.

In the series of standard electrode potentials of metals, chromium is before hydrogen between zinc and iron, molybdenum is also before hydrogen, but not far from it, and tungsten is after hydrogen. Therefore, chromium displaces hydrogen from dilute HC l and H 2 SO 4.

C r + 2 HCl → CrCl 2 + H 2

Cr + H 2 SO 4 → CrSO 4 + H 2

In concentrated H 2 SO 4 and H NO 3 In the cold, chrome becomes passivated. When heated, chromium slowly dissolves in these acids

2Cr + 6H 2 SO 4 conc. → Cr 2 (SO 4 ) 3 + 3SO 2 + 6H 2 O.

Hydrochloric acid and dilute H 2 SO 4 on Mo and W don't work. Molybdenum dissolves only in hot conc. N 2 SO 4 . Tungsten dissolves only in a hot mixture of hydrofluoric and nitric acids

E o + 2H N +5 O 3 + 8 HF → H 2 [E +6 F 8 ] + 2 N +2 O + 4 H 2 O, where E = Mo, W.

At high temperatures, especially in a finely crushed state, C r, Mo, W are quite easily oxidized by many non-metals:

│ O 2 → Cr 2 O 3

│ t o

│ S → CrS

│ t o

Cr + │ Cl 2 → CrCl 3

│ t o

│ N 2 → c floating

│ t o

│ C → c melts

In this case, in the case of chromium, compounds with the most stable CO are most often formed. chromium (+3). When Mo and W with nonmetals, as a rule, compounds are formed in which S.O. element is +6.

Common to elements of the chromium subgroup is the lack of interaction with hydrogen.

The most important compounds of elements of the chromium subgroup

I .Connections E (P), that is, S.O. = +2.

1.Black chromium oxide (P) C rO very difficult to get. It is formed by the oxidation of chromium amalgam (that is, there is no oxide film) with air under normal conditions: C r + ½ O 2 → CrO.

When heated, oxidation continues to C r 2 O 3 .

CrO unstable connection of a basic nature:

With rO + 2HCl → CrCl 2 + H 2 O.

2.Chromium hydroxide (P) C r(OH)2 a yellow substance insoluble in water, which is obtained by alkalizing solutions of chromium salts (P):

With rCl 2 + 2NaOH → Cr(OH) 2 ↓ + 2NaCl.

Chromium hydroxide (P) is basic in nature, that is, it interacts only with acids and does not dissolve in alkali solutions:

C r (OH) 2 + 2 HCl ↔ CrCl 2 + 2 H 2 O.

Сr(OH)2 is a weak base.

With r (P) forms a number of complexes. For chromium in S.O. +2 is characterized by a coordination number of 6. For example, in aqueous solutions the C ion r 2+ hydrates, forming blue aqua complexes [ Cr(H2O)6]2+ . Chromium halides (C) absorb ammonia gas, forming ammonia:

Cr Cl 2 + 6NH 3 → Cl 2 .

P. Connections E (Sh), that is, S.O. = +3

At chrome S.O. +3 in connections is the most stable.

Chromium oxide (III) C r 2 O 3 is obtained:

a) when metallic chromium powder is heated in air:

4Cr + 3O 2 → 2Cr 2 O 3

b) calcination of chromium oxide ( VI ) or ammonium dichromate:

t o t o

4CrO 3 → 2Cr 2 O 3 + 3O 2 (NH 4 ) 2 Cr 2 O 7 → Cr 2 O 3 + N 2 + 4H 2 O;

c) when heating chromium hydroxide (III):

2 C r(OH) 3 → Cr 2 O 3 + 3H 2 O.

Amorphous oxide Cr 2 O 3 dark green powder. Crystal modification Cr2O3 black powder. It is highly refractory and chemically inert.In water, acids and alkali solutions do not dissolves . However, when fusing the oxide Cr (III) with alkalis and basic oxides salts of metachromic acid are formed:

t o t o

Cr 2 O 3 + 2KOH → 2K Cr O 2 + H 2 O; Cr 2 O 3 + CaO → Ca (Cr O 2 ) 2.

When fusing Cr 2 O 3 with potassium disulfate, chromium sulfate (III) is formed.

3K 2 S 2 O 7 = 3K 2 SO 4 + 3SO 3;

Cr 2 O 3 + 3SO 3 = Cr 2 (SO 4 ) 3

─────────────────────────

Cr 2 O 3 + 3K 2 S 2 O 7 = Cr 2 (SO 4 ) 3 + 3K 2 SO 4 .

These reactions show amphoteric character Cr 2 O 3 .

Chromium hydroxide (III) Cr(OH)3 precipitated from solutions of chromium (III) salts with alkalis in the form of a voluminous gelatinous grayish-greenish precipitate, insoluble in water.

C r +3 + 3 OH - → Cr (OH) 3 ↓;

CrCl 3 + 3NaOH → Cr(OH) 3 ↓ + 3NaCl.

Chromium hydroxide Cr(OH)3 It is amphoteric in nature and freshly obtained chromium hydroxide (III) is easily dissolved in acids and alkali solutions.

Cr (OH) 3 + 3HC l ↔ CrCl 3 + 3 H 2 O

Cr(OH) 3 + NaOH ↔ Na.

The basic and especially acidic properties of chromium hydroxide (III) are weakly expressed. Therefore salt Cr +3 undergo significant hydrolysis in solutions, and soluble chromites in the absence of excess alkali are hydrolyzed almost completely.

Cr 3+ + HOH ↔ Cr(OH) 2+ + H + ;

3- + HOH ↔ 2- + OH - .

Alum. Cr (Ш), like A l (Ш), forms with active metals and NH4+ double salts alum. Example: KS r (SO 4 ) 2 ∙12 H 2 O and (NH 4 )Cr (SO 4 ) 2 ∙12 H 2 O. They are formed during the interaction of solutions of M 2 +1 SO 4 and Cr 2 (SO 4 ) 3 . In solution these salts dissociate:

K Cr(SO 4 ) 2 ↔ K + + Cr 3+ + 2 SO 4 2- .

Cr 3+ + H 2 O ↔ Cr (OH) 2+ + H + - acidic environment.

Cr (W) as well as Cr (P) - active complexing agent. Coordination number Cr(Ш) is equal to 6 and 4.

Examples: aqua complex [ Cr(H2O)6]3+ - blue-violet color;

hydroxo complex [ Cr(OH)6]3- - emerald green color;

amino complex [Cr (NH 3 ) 6 ] 3+ - purple color.

Sh. Compounds E (VI), that is, S.O. = +6

Compounds in which S.O. element is +6, most characteristic of Mo, W and to a lesser extent for Cr.

oxides EO 3 (Cr O 3 , MoO3 AndWABOUT3 ).

MoO3 AndWABOUT3 are formed when metals are heated in air:

2E + 3O2 → 2EO3.

CrABOUT3 can be obtained only indirectly, since when heatedCris formed in the airCr2 ABOUT3 .

CrABOUT3 precipitates when adding excess concentratedH2 SO4 to a saturated chromate solution:

TO2 CrABOUT4 + H2 SO4 conc. = CrABOUT3 ↓ + K2 SO4 + N2 ABOUT

MoO3 colorless crystals;

WABOUT3 light yellow crystals;

CrABOUT3 dark red crystals.

MoO3 AndWABOUT3 are stable and, when heated, pass into the gas phase without decomposition. When heatedCrABOUT3 easily decomposes, releasing O2 .

4 CrABOUT3 → 2 Cr2 ABOUT3 + 3О2 .

CrABOUT3 easily dissolves in water, forming chromic acid

CrABOUT3 + N2 O → N2 CrABOUT4 .

MoO3 AndWABOUT3 do not dissolve in water. The acidic nature of these oxides manifests itself when dissolved in alkali solutions:

EO3 + 2KON → K2 EO4 + N2 ABOUT.

In this case, salts of chromic, molybdic and tungstic acids are formed, respectively.

Hydroxides E (VI) - H2 EO4

N2 CrABOUT4 , N2 MoO4 , N2 WABOUT4 .

Chromic acid is prepared by dissolvingCrABOUT3 in N2 A. Molybdic and tungstic acids are obtained indirectly - by acidifying solutions of their salts:

(NH4 ) 2 MOO4 + 2HNO3 →H2 MOO4 ↓ + 2NH4 NO3 .

Strength of acids in the H series2 WITHrO4 N2 MoO4 - N2 WO4 decreases.

Chromic acid H2 WITHrO4 medium strength acid (K1 = 2∙10 -1 , TO2 =3∙10 -7 ), not highlighted in the free state.

H2 MoO4 highlighted in free form. It is a white powder, almost insoluble in water. Constants of the first stage of acid and base dissociationH2 MoO4 are of order 10 respectively-2 and 10-13 .

3.Salt.

Most salts of acids H2 EO4 sparingly soluble in H2 A. Only salts dissolve wellNa+ and K+ . Chromates are colored yellow by the C ionrO4 2- , molybdates and tungstates are colorless. All salts of chromic acids are poisonous.

When a chromate solution is acidified, hydrochromate is formed, which is very unstable and, releasing water, turns into dichromate:

2 CrO4 2- + 2H+ ↔ 2Н СrO4 - ↔ Cr2 ABOUT7 2- + N2 ABOUT.

In this case, the yellow color of the solution changes to orange, characteristic of the C ionr2 ABOUT7 2- . This balance is very fluid. It can be shifted by changing the pH of the medium: adding acids (H ions) to the solution+ ) shifts the equilibrium towards the formation of dichromate, and the addition of alkali - to the left (due to the binding of H ions+ ). Thus, in the presence of an excess of OH ions- practically only C ions exist in solutionrO4 2- , and with an excess of hydrogen ions - C ionsr2 ABOUT7 2- .

2K2 WITHrO4 + H2 SO4 → K2 WITHr2 ABOUT7 + K2 SO4 + N2 ABOUT;

TO2 WITHr2 ABOUT7 + 2KON → 2K2 WITHrO4 + N2 ABOUT.

Dichromic acid H2 WITHr2 ABOUT7 much stronger than chromium, K2 = 2∙10 -2 . It is also not highlighted in free form.

ConnectionsCr (VI) - strong oxidizing agents, pass in redox reactions into derivatives Cr(III). Most strongly oxidizing propertiesCr(VI) are expressed in an acidic environment.

TO2 WITHr2 ABOUT7 + 6KJ + 7H2 SO4 →Cr2 (SO4 ) 3 + 3J2 +4K2 SO4 + 7H2 O.

In this case, the orange color of the potassium dichromate solution changes to a green or greenish-violet color of the solutionsCr3+ .

In contrast to chromium, the oxidizing properties of Mo(VI) AndW (VI) even in an acidic environment, appear only when interacting with the strongest reducing agents, for example, hydrogen at the time of release.

Chromium peroxide

When an acidic solution of chromate or dichromate is treated with hydrogen peroxide, chromium peroxide C is formedrO(O2 ) 2 or withrO5 .

WITHr2 ABOUT7 2- + 4H2 ABOUT2 + 2H+ = 2 CrO (O2 ) 2 + 5 H2 O.

CrO (O2 ) 2 blue in color, unstable in aqueous solution and decomposes into oxygen and aqua complexes [Cr(H2 O) 6 ] 3+ .

Chromium peroxide is stable in ether and forms a peroxo complex

CrO (O2 ) 2 ∙ L, WhereLether, pyridine, etc. These complexes have the shape of a pentagonal pyramid with an oxygen atom at the top:

Chromium peroxide contains two peroxide groups (-O-O-), due to which it exhibits oxidizing properties.

13541 0

Group 16 includes O, S, Se, Te, Po (Tables 1 and 2). The valence shell of elements of this group is formed by two electrons in the s orbital and four in the p orbital (s 2 p 4). The word "chalcogen" comes from two Greek words meaning "copper" and "born". Most copper ores consist of copper compounds with oxygen and sulfur, and some of them also contain Se and Te. The most important ores contain compounds with sulfur, for example, “chalcocite” - copper(I) sulfide Cu 2 S, “chalcopyrite” - CuFeS 2. Elements that have an affinity for sulfur are called chalcophiles. These include Cu, Pb, Zn, Hg, As, Sb. Ores with these metals are known - “galena” (lead luster PbS), “sphalerite” (zinc blende ZnS), “cinnabar” (HgS), “realgar” (As 4 S 4), “stibnite” (Sb 2 S 3) .

Table 1. Some physical and chemical properties of group 16 metals

Name	Relates, at. weight	Electronic formula	Radius, pm	Main isotopes (%)
Oxygen Oxygen [from Greek. ohu genes - acid-forming]			covalent (single bond) 66
Sulfur Sulfur [Sanskrit, sulvere - sulfur, lat. Sulphurium]			atomic 104(S 8) covalent 104
Selenium Selenium [from Greek. Selene - Moon]		3d 10 4s 2 4р 4	atomic 215.2 (gray) covalent 117
Tellurium Tellurium [from lat. tellus - earth]		4d 10 5s 2 5 R 4	atomic 143.2, covalent 137	128 Te (31.73)
Polonium Polonium [in honor of Poland]		4f 14 5d 10 6s 2 6р 4	atomic 167, covalent 153	210,211*,216,218 Po (traces)

As a rule, elements of group 16 form compounds in which they have an oxidation state of -2, especially in compounds with N and reactive metals. In oxides, the most common valences are +4 and +6. Like p-elements of other groups, when moving to the bottom of the group they show a gradual change from non-metallic properties to metallic ones: ABOUT And S- typical non-metals, Se And Those- semimetals, Ro- metal (highly radioactive).

Table 2. Content in the body, toxic (TD) and lethal doses (LD) of group 16 metals

In the earth's crust (%)	In the ocean (%)	In the human body
In the earth's crust (%)	In the ocean (%)	Average (with body weight 70 kg)		Blood (mg/l)
	Included in water				non-toxic in the form of O 2, toxic in the form of O,
					non-toxic
	(0.15-1.8)x10 -11		(0.42-1.9) x10 -4		TD 5 mg, LD nd
	(0.7-1.9)x10 -11				TD 0.25 mg, LD 2 g
Traces in uranium ores

Oxygen (O) - a colorless, odorless gas. Extremely reactive, forming oxides with all elements except noble gases. In industry, it is used in steel smelting, metal cutting and chemical production. Oxygen-containing compounds with H, Si, Ca, Al, F e constitute 49% of the mass of the earth's crust, 89% of the mass of the world's oceans and, in the form of diatomic molecules ABOUT 2.21% of the earth's atmosphere. It is part of many hundreds of thousands of compounds and is essential for life, as it participates in the respiration processes of living organisms. It is the most important factor in chemical and biological evolution on Earth. Violation of the processes of neutralization of active forms ABOUT 2, formed during metabolism, is believed to accelerate the aging process of the body.

Oxygen has a high electronegativity (3.5 on the corresponding scale), which provides strong oxidizing properties. The reactions of oxide formation are highly exothermic, and can be accompanied by combustion of the combined ABOUT 2 elements or compound formed. Due to its small atomic size coupled with high electronegativity, oxygen is able to stabilize atoms of other elements in states with high oxidation states, e.g. Cl 2 O 7 2- or in Cr 2 O 7 2- . Oxides of metallic elements are usually basic, while oxides of non-metallic elements are acidic. Therefore, they can combine with each other to form salts.

There is a classification of oxides according to composition: 1. Normal oxides contain bonds only between the element and oxygen, e.g. MgO, SO 3 , SiO 2 . 2. Peroxides contain bonds not only between an element and oxygen, but also between two oxygen atoms, for example, Na 2 O 2 and N 2 O 2. Peroxides are strong oxidizing agents. 3. Mixed oxides are a mixture of two oxides, for example, trilead tetroxide (red lead) Pb 3 O 4 - a mixture of two parts РbО and one part РbO 2 .

One of the allotropes of oxygen is triatomic ozone(O 3), which in nature is formed in the upper layers of the atmosphere under the influence of ultraviolet radiation from the Sun or in electrical discharges during thunderstorms. In laboratory conditions, ozone is obtained in ozonizers by passing O 2 through a weak electrical discharge. Currently, ozone is used to disinfect drinking water at waterworks, since it is a stronger oxidizing agent than regular ozone. O 2. When ozone enters the body, it affects the lungs, forming peroxide metabolites.

O 2 has a pronounced electron affinity (142 kJ/mol). This provides a high ability to form anionic superoxide ion * O 2 - which is a highly reactive radical. These properties of superoxide ions determine their high toxicity. Hyperoxia and excess ozone initiate homolytic (when shared electrons upon breaking a bond are distributed equally between two atoms) splitting of chemical bonds in biomolecules. In this case, radicals with an unpaired electron are formed. For example, when reacting ROOH With ABOUT 2- carbon peroxide is formed * ROO- and hydrogen peroxide * NOO- radicals. Superoxide ion reacts actively with organic substances of the RH type, especially with unsaturated bonds. The resulting organic radicals initiate a chain process of oxidation of organic substances. Accumulated organic peroxides are normally destroyed peroxidases, as well as antioxidants - tocopherol(vitamin E) and thiol compounds ( glutathione, cysteine).

In a healthy body, there are several levels of defense mechanisms against oxygen radicals: cytochrome oxidase(almost undamaged by excess oxygen), various amines, γ-aminobutyric acid, etc.

Sulfur (S) - found in nature in native form, as well as in sulfide ores of metals (for example, in pyrite - “iron pyrite” - FeS 2, zinc blende ZnS, galena PbS), in natural gas H 2 S. Sulfur is a key element for the chemical industry. Has several allotropic modifications, the most stable are enantiotropes S 8 . They consist of rhombic lemon yellowα -sulfur And monoclinic honey yellowβ -sulfur. Among other allotropes are known cuttings, amorphous, colloidal and plastic sulfur. Sea water contains sulfate ions.

Atoms S have 6 electrons in the outer shell and can add two more electrons to their half-filled 3p orbitals to form a sulfide ion S 2- . Atoms can exist in states with valence -2, +2, +4, +6. Several oxides are known, of which two are the most significant: dioxide SO 2 and trioxide SO 3 .

Dioxide sulfur is a dense colorless gas with a sharp suffocating odor, easily dissolving in water to form weak sulfurous acid. It is used in the pulp industry, for bleaching fabrics, as an antiseptic for long-term storage of vegetables and fruits. In the atmosphere, oxidizing to trioxide, it causes the formation of acid rain. Its oxidation is catalyzed by trace amounts of iron and manganese contained in the atmosphere.

Trioxide is a powerful oxidizing agent and has pronounced acidic properties. Reacts exothermically with water, forming strong sulfuric acid. Saturated solution MgSO 4 *7H 2 O(“Epsom salt”) is used in medicine as an anti-inflammatory agent.

Sulfur is one of the 6 organogens ( S, N, N, O, S, P), making up the bulk of organic molecules. It is part of the biological tissues of all living beings in the form of the amino acids cysteine, cystine and methionine. Like phosphorus, it functions as a carrier of functional groups and energy. The presence of paired cysteine residues causes the formation of disulfide bonds in proteins (- S-S-), determining their spatial structure. Sulfhydryl (“thiol”) groups (- SH) cysteine molecules are part of the active centers of many enzymes.

S easily donates electrons to metal atoms, forming coordination compounds with high stability constants, for example, in structures with a high content of keratin (hair, nails, feathers, claws, hooves).

The last three elements of group 16 ( Se, Te, Po) form hex-valent fluorides, although the oxidation process is difficult, especially for elements lower in the Periodic Table. Have an effect inert pair- the behavior of an element as if two of its valence electrons were missing. Selenides, tellurides and metal polonides are almost always isomorphic with the corresponding sulfides. This explains their presence together with sulfur in nature.

Selenium (Se) - found in some sulfide ores. It is obtained during the electrolytic purification of copper (as a by-product) in the form of a silver allotropic modification, the crystal structure of which consists of helical (twisted in one direction) chains Se∞ or as a less stable red amorphous powder consisting of cycles Se 8 in the shape of a crown. Selenium burns in air. Below the melting point (490°K) it is a semiconductor. Important property Se is the ability to generate electric current in light. Therefore, it is used in photovoltaic cells, photocopiers, solar cells and semiconductors.

In oxides it most often exhibits oxidation states of +4 and +6. Oxides correspond to selenium ( H2SeO3) and selenium ( H2SeO4) acids. Like sulfur trioxide SeO 3 is a strong oxidizing agent, but due to thermodynamic instability, selenates in living organisms are reduced to selenites, which can easily react with sulfhydryl groups of bioorganic compounds. Acids are dibasic and form two sets of salts with metal ions.

Many connections Se very toxic, especially H 2 Se. The maximum permissible concentration of hydrogen selenide is an order of magnitude lower than that of such a well-known poison as hydrocyanic acid HCN. Even in very low concentrations it causes headache and nausea, and in high concentrations it causes acute irritation of the mucous membranes. All selenides, many organic compounds Se, soluble selenites and selenates upon contact with the skin cause eczema and local inflammation. Selenide intoxication is manifested by impaired sense of smell and increased sweating; their release from the body occurs slowly. Of the selenium compounds, only sulfides from Se 2 S before SeS 3 (Bagnall, 1971). Sulfide SeS 2 are used in cosmetics. Excess Se in soil causes the disease “alcololysis” in livestock.

With organic acids, Se forms salts with a valence of +2. Only simple salts are known: methyl thiosulfonates Se(S 2 O 2 CH 3) 2, dialkyl dithiocarbamates Se(S 2 CNR 2) 2 and alkyl xanthogenates Se(S 2 COR) 2. They are easily destroyed when heated. A variety of carbon compounds are also known, from simple carbon selenides CSe 2 and CSSe to saturated and unsaturated heterocyclic molecules such as selenanthrene, cycloselenopropane and selenonaphthene (Fig. 1). Biological reactions of carbon compounds Se poorly studied.

Rice. 1.

For chronic exposure Se accumulates in the liver and kidneys, as well as in other organs: in noticeable quantities in bones, hair and nails, in minimal quantities in the brain. Se is part of selenoproteins, in particular, the prosthetic group glutathione peroxidase, which together with tocopherol (vitamin E) protects cell membranes from damage by free radicals. Highly active free radical compounds can be formed in a number of important processes, for example, during activation of phagocytes or exposure to ionizing radiation.

Selepoproteins are such important enzymes as deiodiasis, ensuring homeostasis of thyroxine and through calcitonin - homeostasis Sa, selenoprotein N regulating myocyte regeneration. Apparently, selenoproteins play a significant role in the body’s antiviral defense. Selenium deficiency has been identified in some areas of China and is manifested by endemic cardiomyopathy (“Keshan disease”). Antioxidant properties Se used for cancer prevention.

Tellurium (Te) - accompanies other metals (for example, gold in the mineral calaverite); it is obtained from anode sludge during copper purification. Occurs as a rare mineral tellurite. Pure metal Those It looks silvery-white, burns in air, and is toxic in any form. The vapor has a garlicky odor. In industry, it is used in alloys to improve their mechanical properties, for the production of chemical reagents, catalysts, and in electronics - as a semiconductor.

Polonium (Po) is a very rare and volatile radioactive silver-gray metal. Formed when bismuth atoms are bombarded by neutrons. It is used as a heat source in space equipment and a source of alpha particles for scientific research. Extremely toxic due to high decay energy.

Medical bioinorganics. G.K. Barashkov

Group 6 elements include: oxygen (8 O), sulfur (16 S), selenium (34 Se), tellurium (52 Te) and polonium (84 Po). The name of the group “Chalcogens” is literally translated as “giving birth to salts” from the Greek. "chalcos" - copper and "genos" - genus, origin. In nature, chalcogens actually occur most often in the form of copper compounds (except for oxygen) - these are copper sulfides and selenides. Copper(II) sulfide Copper(I) selenide

Ø When moving from oxygen to polonium, the size of the atom increases, non-metallic properties weaken, and metallic properties increase: oxygen and sulfur are typical non-metals, selenium and tellurium are metalloids with non-metallic properties, polonium is a metal. Ø The ionization energy of atoms in the same series decreases E 0 → E+, which means an increase in reducing properties (the ability to give up an electron). Ø Due to the high electron density and strong interelectron repulsion, the electron affinity energy for oxygen is less than for all other elements of the 6-A group. As a result, the S 2 anion is much more stable than similar selenium and tellurium anions, and the O 2 anion practically does not exist in free form. Ø A decrease in electronegativity in the series from oxygen to polonium means a decrease in the degree of polarity of covalent bonds in the series of chalcogens.

The higher the binding energy, the stronger the bond and the more stable the connection. Due to the high electron density and interelectron repulsion forces, the single bond between oxygen atoms (O – O) is the least stable. On the contrary, the formation of a double bond is much more favorable for oxygen, since the energy of a double bond significantly exceeds the energy of two single bonds.

Forms of existence of compounds 6 -A group E Simple. substance N 2 E EO 3 N 2 EO 4 CO O O 2 H 2 O - - -2, -1, 0, +1, +2 S S Н 2 S SO 3 Н 2 SO 4 -2, 0, + 2, +4, +6 Se Se Н 2 Se Se. O 3 H 2 Se. O 4 -2, 0, +2, +4, +6 Te Te H 2 Te Te. O 3 H 2 Te. O 4 -2, 0, +2, +4, +6 Po Po H 2 Po Po. O 3 -2, 0, +2, +4, +6 -

The electronic configuration of an unexcited atom is 1 s 22 p 4. In most compounds it exhibits an oxidation state of -2, but compounds with an oxidation state of -1 are known; 0; +1; +2; +4. The most common element in the earth's crust (58%). 3 stable isotopes were detected: 168 O (99.759%), 178 O (0.037%), 18 O (0.204%). 8 More than 1,400 minerals are known to contain oxygen.

Features of the structure of the oxygen atom 1) There is no low-energy d-sublevel; 2) High electronegativity (place after fluorine), unable to donate more than 2 electrons (all other elements of the subgroup exhibit the highest oxidation state of +6) 3) Small atomic radius

Structure of the oxygen molecule 1) 12 electrons of the outer energy level are located in 8 molecular orbitals (MO); 2) The orbitals of oxygen atoms are close in energy, so non-bonding orbitals are not formed; 3) 8 electrons are located in bonding orbitals and 4 electrons are in antibonding orbitals, therefore the bond order is (8 – 4)/2 = 2; 4) The presence of 2 unpaired electrons in antibonding π* orbitals gives the oxygen molecule paramagnetic properties.

Physical properties of oxygen 1) The diatomic molecule is the most stable, since the dissociation energy O = O is 494 kJ/mol, while the O – O bond is only 210 kJ/mol; 2) Oxygen molecules are weakly polarized, so the intermolecular bonds between them are weak. Tpl. = -218 °C Bp. = -183 °C. 3) Poorly soluble in water (5 volumes in 100 volumes of water at 0 °C). 4) Liquid and solid oxygen are attracted by a magnet, because its molecules have paramagnetic properties. 5) Solid oxygen is blue, and liquid oxygen is blue.

There are 3 known allotropic modifications of oxygen O 2, O 3 (ozone) and O 4 - an unstable tetraoxygen.

Allotropic modifications of oxygen A bluish gas with a characteristic pungent odor. The ozone molecule is polar. Slightly soluble in water Tm. = -193 °C Bp. = -112 °C

Chemical properties of oxygen 1) Strong oxidizing agent. 2) Does not directly interact only with inert gases, halogens, silver, gold and platinum group metals (except osmium). Gold (III) oxide Au – yellow O – red

Oxygen production 1) Biological origin of oxygen: 6 H 2 O + 6 CO 2 hν C 6 H 12 O 6 + 6 O 2 400 -500 °C 2) Decomposition of potassium chlorate (Berthollet salt): KCLO 3 3 O 2 + 2 KCl 3) Decomposition of potassium permanganate: KMn. O 4 210 -240 °C K 2 Mn. O 4 + Mn. O2 + O2

Oxygen compounds (-2) The water molecule has a tetrahedral (not planar) structure. The oxygen atom is in the state. sp 3 hybridization The BC method does not explain why one of the two lone electron pairs (LEP) is more active than the other

Structure of the water molecule (MO method) v Two 1 s atomic orbitals (one from each hydrogen atom), one 2 s orbital and three 2 p orbitals of the oxygen atom take part in the formation of a water molecule. In total, 8 electrons are located in 6 molecular orbitals of a water molecule. v The interaction (overlap) of 1 s orbitals of two hydrogen atoms with 2 s and 2 px orbitals of an oxygen atom leads to the formation of 4 MOs, of which two are bonding (2 a 1, 1 b 1) and 2 are antibonding (4 a 1, 2 b 1). v AOs with lower energy (oxygen) make a greater contribution to the energy of bonding orbitals, and AO with higher energy (hydrogen) make a greater contribution to the energy of antibonding orbitals. v 2 pz-orbital of oxygen and 1 s-orbital of one of the hydrogens weakly overlap and form a weakly bonding orbital (3 a 1). v The 2 py orbital of the oxygen atom is perpendicular to the plane of overlapping orbitals and does not overlap with the 1 s orbitals of the at. H and, therefore, forms a nonbonding 1 b 2 orbital.

Structure of a water molecule (MO method) In total, 2 binding, 2 non-binding and 2 antibonding MOs are formed in a water molecule. 8 electrons are located in pairs in 2 bonding and 2 nonbonding orbitals.

Structure of water (MO method) The electrons of MO 2 a 1, 1 b 1 form O – H bonds, and ē, located on non-bonding orbitals (3 a 1, 1 b 2), correspond to free electron electrons of the water molecule. However, MOs 3 a 1 and 1 b 2 differ in localization and energy. 1 b 2 is localized on the oxygen atom and has a purely p-character. 3 a 1 has lower energy and is delocalized, since it is formed with the participation of AO at. H and O. Localization of 1 b 2 on the oxygen atom leads to the fact that the negative charge in the water molecule is concentrated near the oxygen atom, and positive charge near the hydrogen atom. Hence the polarity of the water molecule (μ = 1.84 D).

Hydrogen bonds in a water molecule Consequences of the polarity of water molecules: 1) The ability to form intermolecular hydrogen bonds; 2) High melting and boiling points (0 °C and 100 °C) 3) Strong surface tension

Boiling point of chalcogens H 2 O Bp. , °С H 2 S 100 -60 H 2 Se H 2 Te -42 -2

Hydrogen bonds in an ice molecule Ice crystals have a hexagonal structure. Each water molecule is connected to three neighboring molecules through hydrogen bonds. Each water molecule forms 4 hydrogen bonds using both LEPs. When melting, one hydrogen bond is broken (at 0 °C 15% of the hydrogen bond is broken). In this case, some of the molecules get inside the frame, this explains the fact: at 4 °C the density of water is maximum.

Oxides In accordance with the nature of the element in a positive oxidation state, the nature of the oxides in the periods and groups of the periodic system naturally changes. During the periods, the effective negative charge on the oxygen atoms decreases and a posterior transition from basic through amphoteric oxides to acidic oxides occurs.

Interaction of oxides with water In the series of acidic oxides, the acidic properties and ability to interact with water increase during the period, which is confirmed by a decrease in the potential ∆G 0298.

Superoxides The addition of one electron to an oxygen molecule causes the formation of superoxide ion O 2 -. O2 derivatives are called superoxides, which are known for the most active alkali metals (potassium, rubidium, cesium). Superoxides are formed by the direct interaction of simple substances: K + O 2 = KO 2. The unpaired electron of the O 2 - ion determines the paramagnetism of superoxides and their color. Superoxides are very strong oxidizing agents. They react violently with water, releasing hydrogen.

Send your good work in the knowledge base is simple. Use the form below

Students, graduate students, young scientists who use the knowledge base in their studies and work will be very grateful to you.

Posted on http://www.allbest.ru

CHEMISTRY OF ELEMENTSVIGROUPS

16th group of the periodic table of chemical elements D.I. Mendeleev consists of Chalcohemny (from the Greek chblkpt - copper (in the broad sense), ore (in the narrow sense) and genpt - giving birth). These chemical elements include oxygen (O), sulfur (S), selenium (Se), tellurium (Te), radioactive polonium (Po), and artificially produced livermorium (Lv). All of these elements occur naturally (except Lv), and each has several naturally occurring isotopes (except Po). All of them are p-elements. The structure of the valence level of their electron shells corresponds to the formula ns 2 n.p. 4 .

oxygen sulfur selenium tellurium polonium

1. Oxygen

The most common element on Earth: in the air - 21% by volume (located in the form of O 2 and O 3), in the earth's crust - 49% by weight, in the hydrosphere - 89% by weight, in living organisms - up to 65% by weight mass. In the earth's crust it is contained in the form of various minerals, which are oxides (for example, Al 2 O 3, SiO 2, Cu 2 O, SnO 2) and salts (for example, FeCO 3, CaSO 4, Ca 3 (PO 4) 2 ).

Atom

Ordinal number 8, electronic structure: 1s 2 2s 2 2p 4

Valence

II, in CO - valence III due to DAS.

Oxidation states

1) -2 - oxides, hydroxides, salts;

2) -1 - peroxides;

3) -1/2, -1/3 - superoxides, ozonides;

4) +2 - ОF 2, +1 - O 2 F 2.

Electronegativity

Isotopes of natural oxygen

O (99.76%), O (0.04%), O (0.20%).

Physical properties

Oxygen is a colorless, odorless, and tasteless gas, slightly heavier than air. Poorly soluble in water. Liquid oxygen is a bluish liquid, boiling at -183 0 C. Attracted by a magnet. Solid oxygen is blue crystals that melt at -218.7 0 C.

Molecule structure

A molecule consists of two atoms connected by a double bond. The bond is covalent nonpolar.

Allotropic modifications (changes)

O 2 and O 3 (ozone).

Methods of obtaining

1) Industrial method: distillation of liquid air;

2) Laboratory method: decomposition (exposure to t 0 C) of some oxygen-containing substances

2KMnO 4 = K 2 MnO 4 + MnO 2 + O 2

2KClO 3 = 2KCl + 3O 2 (k - MnO 2)

2H2O2 = 2H2O + O2 (k - MnO 2)

2HgO = 2Hg + O2

2KNO 3 = 2KNO 2 + O 2

Methods for collecting oxygen

Chemical properties

1) Interaction with HeMet (except for fluorine and noble gases): OXIDES are formed - binary compounds with oxygen in oxidation state -2:

Si + O 2 = SiO 2 (t=400-500 0 C)

· When interacting with carbon, phosphorus and arsenic, depending on the amount of oxygen, different oxides are formed:

C + O 2(g) = CO2; C + O 2 (week) = CO

· When interacting with sulfur, the latter is oxidized to sulfur dioxide (further exposure to oxygen leads to the formation of a higher oxide - SO 3):

· When interacting with fluorine, oxygen FLUORIDE is formed:

2F 2 + O 2 = 2OF 2

2) Interaction with Met: basic and amphoteric oxides are formed:

4Al + 3O 2 = 2AL 2 O 3

· When interacting with sodium, peroxide is formed:

Na + O 2 = Na 2 O 2

· When interacting with potassium, rubidium and cesium, superoxides are formed:

· When interacting with iron, a mixture of oxides is formed:

3Fe + 2O 2 = Fe 3 O 4 (Fe 2 O 3 *FeO)

· When interacting with manganese, manganese dioxide is formed:

3) Interaction with complex substances:

· Combustion and roasting of sulfides, hydrogen compounds:

4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2

2H 2 S + 3O 2 = 2SO 2 + 2H 2 O

CH 4 + 2O 2 = CO 2 + 2H 2 O

Oxidation of lower oxides to higher ones:

FeO + O 2 = Fe 2 O 3

· Oxidation of hydroxides and salts in aqueous solutions: if the substance is unstable in air:

2HNO2 + O2 = 2HNO3

4Fe(OH) 2 + O 2 + 2H 2 O = 4Fe(OH) 3

Oxidation in the presence of catalysts:

ammonia: NH 3 + O 2 = NO + H 2 O

organic substances: C 2 H 5 OH + O 2 = CH 3 -COH, etc. (k - Cu; t 0 C)

Application of oxygen

Oxygen is used in medicine, in blasting operations, for welding metals, for cutting metals, in aviation for breathing, in aviation for engines, and in metallurgy.

CH 4 + 2O 2 = CO 2 + 2H 2 O

C 6 H 10 O 5 + 6O 2 = 5CO 2 + 6H 2 O

In all these processes carbon monoxide (IV) is formed. The only natural process for the binding of carbon monoxide (IV) is the process of photosynthesis, which takes place in green plants under the influence of sunlight:

6CO 2 + 6H 2 O = C 5 H 12 O 6 + O 2 (k - hv)

In this case, glucose is formed - the basis for the construction of plant tissues.

If a substance reacts slowly with oxygen, then such oxidation is called slow. These are, for example, the processes of food decomposition and rotting.

Ozone

Ozone is an allotropic modification of oxygen.

Physical properties

The gas, with the smell of fresh pine needles, is colorless.

Receipt

1) Air ozonation: 3O 2 2O 3

2) During a thunderstorm (in nature);

3) In the laboratory - in an ozonizer.

Chemical properties

1) Unstable, easily decomposes: O 3 O 2 + O. . This produces ATOMIC oxygen, a very strong oxidizing agent. It decolorizes dyes, reflects UV rays, and destroys microorganisms;

2) Strong oxidizing agent, stronger than oxygen:

6NO 2 + O 3 = 3N 2 O 5

3PbS + 4O 3 = 3PbSO 4

3) Qualitative reaction to ozone: reaction with potassium iodide, a yellow-brown color of iodine appears:

2KI + O 3 = 2KOH + I 2 + O 2

Hydrogen peroxide

It is worth paying special attention to this substance, since of all oxygen compounds, hydrogen peroxide is most often used as a sterilizing and antiseptic agent.

Molecule structure

H 2 O 2 or N-O-O-N

Physical properties

It is a colorless, unstable liquid. The density is 1.45 g/cm3. Its concentrated solution (30%) is explosive and is called perhydrol.

Receipt

Hydrolysis of metal peroxides with water or acids:

BaO 2 + H 2 SO 4 = H 2 O 2 + BaSO 4

Chemical properties

1) Decomposition:

2H 2 O 2 = 2H 2 O + O 2 (t 0 C, k - MnO 2)

2) Hydrogen peroxide can be both an oxidizing agent and a reducing agent:

· Oxidizing properties are more characteristic - turns into H 2 O or OH -:

Na 2 S +4 O 3 + H 2 O 2 = Na 2 S +6 O 4 + H 2 O

(non-metals in an intermediate oxidation state are oxidized to a higher oxidation state)

2Fe(OH) 2 + H 2 O 2 = 2Fe(OH) 3

2KI + H 2 O 2 + H 2 SO 4 = I 2 + K 2 SO 4 + 2H 2 O

PbS + 4H 2 O 2 = PbSO 4 + 4H 2 O

(sulfides turn into sulfates)

2Cr +3 Cl 3 + 3H 2 O 2 + 10KOH = 2K 2 Cr +6 O 4 + 6KCl + 8H 2 0

(any chromium compound +3 oxidizes to +6)

· Reducing properties - goes into O 2:

CaOCl 2 + H 2 O 2 = CaCl 2 + O 2 + H 2 O

2KMnO 4 + 5H 2 O 2 + 3H 2 SO 4 = 2MnSO 4 + K 2 SO 4 + 5O 2 + 8H 2 O

K 2 Cr 2 O 7 + 3H 2 O 2 + 4H 2 SO 4 = Cr 2 (SO 4) 3 + K 2 SO 4 + 3O 2 + 7H 2 O

Metal peroxides and superoxides

The chemical properties of these compounds have their own peculiarity: oxygen is formed in the ORR, as well as in decomposition reactions with water:

Chemical properties

1) Decomposition by water:

Na 2 O 2 + H 2 O = H 2 O 2 + NaOH

K 2 O 2 + H 2 O = H 2 O 2 + O 2 + KOH (at t 0 C)

2) They are strong oxidizing agents:

KO 2 + Al = KAlO 2

3) In some reactions they exhibit reducing properties:

2KMnO 4 + 5Na 2 O 2 + 8H 2 SO 4 = 2MnSO 4 + K 2 SO 4 + 5Na 2 SO 4 + 5O 2 + 8H 2 O

4) Interact with carbon dioxide:

Na 2 O 2 + CO 2 = Na 2 CO 3 + O 2

Receipt

1) Me (alkaline and Ba) + oxygen:

Na + O 2 = Na 2 O 2

2) Metal oxides + oxygen:

2K 2 O + 3O 2 = 4KO 2

2. Sulfur

Sulfur is an element of the 16th group, the third period of the periodic table of chemical elements D.I. Mendeleev, with atomic number 16. Exhibits non-metallic properties. Denoted by the symbol S (Latin sulfur).

Atom

Ordinal number 16, electronic structure: 1s 2 2s 2 2p 6 3s 2 3p 4.

Valence

Oxidation states

1) -2 - sulfides;

2) 0 - atom;

3) +4 - salts, acids;

4) +6 - salts, acids.

Electronegativity

Physical properties

A yellow crystalline solid, insoluble in water, not wetted by water (floats on the surface), boiling point = 445 0 C.

Allotropic modifications

1) Crystalline:

· Rhombic sulfur - S 8 - the most stable modification;

· Monoclinic sulfur - dark yellow needles. Stable at temperatures above 960C, under normal conditions it turns into rhombic;

2) Plastic sulfur - a brown rubber-like (amorphous) mass.

Being in nature

1) Native sulfur;

2) Sulfides: zinc ZnS, mercury PbS (cinnabar), iron FeS 2 (pyrite);

3) Sulfates: gypsum CaSO 4 *2H 2 O, Glauber’s salt Na 2 SO 4 *10H 2 O.

Obtaining sulfur

1) Method of thermal decomposition of sulfides:

2) Reduction of sulfur dioxide (IV) with carbon: SO 2 - a by-product in the smelting of metals from sulfur ores

SO 2 + C = S + CO 2

3) Oxidation of hydrogen sulfide by a lack of oxygen or sulfur dioxide (IV) in a mixture of SO 2 and H 2 O (waste vapors from metallurgical and coke oven batteries):

2H 2 S + O 2 = 2S + 2H 2 O

2H 2 S + SO 2 = 3S + 2H 2 O

(the latter reaction also occurs in nature during volcanic eruptions)

4) Isolation from natural gas, oil and related petroleum products.

Chemical properties

Under normal conditions, the chemical activity of sulfur is low, but when heated, sulfur is very active and can be both an oxidizing agent and a reducing agent.

1) Interaction with Me:

· With alkaline Me without heating:

2Na + S = Na 2 S

· With the rest of Me (except Au, Pt) - at elevated temperatures:

2Al + 3S = Al 2 S 3

2) Interaction with NeMe:

· With hydrogen: H 2 + S = H 2 S;

· With phosphorus: 2P + 3S = P 2 S 3 ;

· With oxygen: S + O 2 = SO 4;

· With halogens: the chemical activity of sulfur in relation to halogens decreases in the order from fluorine to iodine: sulfur interacts with fluorine at room temperature, with bromine and chlorine - when heated, no compounds were obtained with iodine: S + Cl 2 = SCl 2

With carbon: C + S = CS 2

3) Interaction with oxidizing acids (when heated):

S + 6HNO 3 (conc) = H 2 SO 4 + 6NO 2 + 2H 2 O

4) Interaction with alkali (disproportionation):

3S 0 + 6KOH = K 2 S +4 O 3 + 2K 2 S -2 + 3H 2 O (boiling)

Application of sulfur

Sulfur is used to produce sulfuric acid, for the vulcanization of rubber, in the production of matches and black powder, as a medicine included in ointments against skin diseases.

In agriculture, sulfur in the form of a fine powder is used to combat diseases of plants, bees and other domestic animals.

In everyday life, sulfur is often used to bind toxic mercury into mercury sulfide and then remove it.

Hydrogen sulfide

The hydrogen sulfide molecule has an angular shape, like a water molecule, with a sulfur atom in the center. However, unlike water, hydrogen sulfide molecules are not capable of forming hydrogen bonds with each other.

Hydrogen sulfide is a colorless gas with a characteristic odor of rotten eggs.

t boil = -60 0 C.

A solution of hydrogen sulfide in water is called hydrosulfide acid - it is a very weak dibasic acid (weaker than carbonic acid). Hydrogen sulfide is very toxic.

Hydrogen sulfide is part of volcanic gases and is also found in the water of some mineral springs.

Receipt

1) In nature, hydrogen sulfide is formed during the decay of proteins;

2) It is obtained in the laboratory in the following ways:

Direct synthesis from simple substances:

· By displacement from sulfides, in the voltage series standing to the left of iron:

FeS + 2HCl = FeCl 2 + H 2 S

Hydrolysis of aluminum sulfide in the cold:

Chemical properties

Hydrogen sulfide exhibits the properties of a strong reducing agent in reactions.

1) Interaction with air oxygen when heated to 700 0 C (with explosion):

2H 2 S + 3O 2 (g) = 2SO 2 + 2H 2 O + Q

2H 2 S + O 2 (wk) = 2H 2 O + 2S

2) Interaction with halogens:

H 2 S + I 2 = 2HI + S

H 2 S + 4Cl 2 + 4H 2 O = H 2 SO 4 + 8HCl

3) Interaction with alkalis: forms two series of salts - medium (sulfides) and acidic (hydrosulfides):

H 2 S + 2NaOH = Na 2 S + 2H 2 O

H 2 S + NaOH = NaHS + H 2 O

4) Interaction with soluble salts of heavy metals: copper, silver, lead, mercury, forming BLACK very slightly soluble sulfides:

5) Interaction with oxidizing agents of medium activity with the formation of sulfur and with strong oxidizing agents - oxidizes to sulfuric acid:

H 2 S + Br 2 = S + 2HBr

H 2 S + 2FeCl 3 = 2FeCl 2 + S + 2HCl

H 2 S + 4Cl 2 + H 2 O = H 2 SO 4 + 8HCl

3H 2 S + HNO 3 (conc) = 3H 2 SO 4 + 8NO + 4H 2 O

H 2 S + 3H 2 SO 4 (conc) = 4SO 2 + 4H 2 O

H 2 S + 4PbO 2 = H 2 SO 4 + 4PbO

6) When heated, it decomposes to simple substances:

7) Interaction with silver:

2H 2 S + 4Ag + O 2 = 2Ag 2 S + 2H 2 O

8) Dissociation in water: mainly in the first step:

HS - H + + S 2- (rare)

9) Qualitative reaction to hydrogen sulfide and soluble sulfides - the formation of a dark brown (almost black) PbS precipitate:

H 2 S + Pb(NO 3) 2 = PbS + 2HNO 3

Na 2 S + Pb(NO 3) 2 = PbS + 2NaNO 3

Sulfides

Classification of sulfides

Soluble in water.	Insoluble in water, but soluble in mineral acids (hydrochloric, phosphoric, dilute sulfuric).	Insoluble neither in water nor in mineral acids - only in oxidizing acids.	Hydrolyzed by water, not existing in aqueous solutions.
Alkaline Me and ammonium sulfides.	White and colored sulfides: ZnS, MnS, FeS, CdS.	Black sulfides: CuS, HgS, PbS, Ag 2 S, NiS, CoS.	Sulfides of aluminum, chromium (III), iron (III)
You can displace hydrogen sulfide using hydrochloric acid: FeS + 2HCl = FeCl 2 + H 2 S	It is IMPOSSIBLE to obtain hydrogen sulfide from these sulfides!	Completely decomposed by water: Al 2 S 3 + 6H 2 O = 2Al(OH) 3 + H 2 S

Receipt

1) Heating Me with sulfur:

2Cr + 3S = Cr 2 S 3

2) Soluble sulfides are obtained by the action of hydrogen sulfide on alkali:

H 2 S + 2KOH = K 2 S + 2H 2 O

3) Insoluble sulfides are obtained by exchange reactions:

H 2 S + Pb(NO 3) 2 = PbS + 2HNO 3

(only acid-insoluble sulfides)

ZnSO 4 + Na 2 S = Na 2 SO 4 + ZnS

(any water-insoluble sulfides)

Chemical properties

1) Soluble sulfides are hydrolyzed by anion, alkaline medium:

K 2 S + H 2 OKHS + KOH

S 2- + H 2 OHS - + OH -

2) Me sulfides, standing in the series of stresses to the left of iron (inclusive), are soluble in strong mineral acids:

FeS + 2HCl = FeCl 2 + H 2 S

3) Insoluble sulfides can be converted into a soluble state (oxidized) by the action of concentrated nitric acid:

3CuS + 14HNO 3 = 3Cu(NO 3) 2 + 3H 2 SO 4 + 8NO + 4H 2 O

4) Sulfides can be converted to sulfates with hydrogen peroxide:

CuS + H 2 O 2 = CuSO 4 + 4H 2 O

5) When sulfides are roasted with oxygen, oxides are formed:

2ZnS + 3O 2 = 2ZnO + 2SO 2

Oxides sulfur

Sulfur oxide (IV)

SO 2 - sulfur dioxide, sulfur dioxide (colorless with a pungent odor, highly soluble in water).

Oxidation state of sulfur in this compound: +4

Since the sulfur atom in the compound has an intermediate oxidation state, in reactions (depending on the redox properties of other reagents) it plays the role of both a reducing agent and an oxidizing agent. In addition, it can react without changing the oxidation state.

Receipt

1) When burning sulfur in oxygen:

2) Oxidation of sulfides:

2ZnS + 3O 2 = 2ZnO + 2SO 2

3) Treatment of sulfurous acid salts with mineral acids:

Na 2 SO3 + 2HCl = 2NaCl + SO 2 + H 2 O

4) When treating certain metals with sulfuric acid (conc.):

Cu + 2H 2 SO 4 = CuSO 4 + SO 2 + 2H 2 O

Chemical properties

1) Sulfur dioxide is an acidic oxide. Reacts with water, basic oxides and alkalis:

SO 2 + H 2 O = H 2 SO 3

BaO + SO 2 = BaSO 3

Ba(OH) 2 + SO 2 = BaSO 3 + H 2 O

Ba(OH) 2 + 2SO 2 = Ba(HSO 3) 2

2) Oxidation reactions (S +4 -23S +6):

SO 2 + Br 2 + 2H 2 O = H 2 SO 4 + 2HBr

5SO 2 + 2KMnO 4 + 2H 2 O = K 2 SO 4 = 2MnSO 4 + 2H 2 SO 4

3) Reduction reactions (S +4 +4зS 0):

SO 2 + C = S + CO 2

SO 2 + 2H 2 S = 3S + 2H 2 O

Sulfur oxide (VI)

SO 3 - sulfuric anhydride. Colorless volatile liquid with a suffocating odor (t boil = 43 0 C); “smoke” in air, hygroscopic: SO 3 + H 2 O = H 2 SO 4 + Q.

In this compound, sulfur has the highest oxidation state of +6. Therefore, sulfuric anhydride is an active oxidizing agent.

The chemical activity of sulfur (VI) oxide is very high. Interacts with water, basic and amphoteric oxides, alkalis.

Receipt:

1) Catalytic oxidation of SO 2 with atmospheric oxygen:

2SO 2 + O 2 = 2SO 3 + Q (t 0 C, k - V 2 O 5)

2) Thermal decomposition of sulfates:

Fe 2 (SO 4) 3 = Fe 2 O 3 + 3SO 3 (t 0 C)

3) Interaction of SO 2 with ozone:

SO 2 + O 3 = SO 3 + O 2

4) Interaction of SO 2 and NO 2:

SO 2 + NO 2 = SO 3 + NO

Chemical properties:

1) Interaction with water - the formation of strong dibasic sulfuric acid:

SO 3 + H 2 O = H 2 SO 4 + Q.

2) Interaction with bases:

2NaOH (g) + SO 3 = Na 2 SO 4 + H 2 O

NaOH + SO 3 (g) = NaHSO 4

3) Interaction with basic oxides:

CaO + SO 3 = CaSO 4

4) Dissolution in concentrated sulfuric acid - formation of OLEUM:

H 2 SO 4 (conc) + SO 3 = H 2 S 2 O 7

H 2 SO 4 (conc) + 2SO 3 = H 2 S 3 O 10

Application:

Sulfuric anhydride is used mainly to produce sulfuric acid.

Sulfur acids and their salts

Sulfurous acid and its salts

H 2 SO 3 - is formed by the reaction of sulfur (IV) oxide with water and exists only in the form of a solution. It is a weak, volatile, unstable dibasic acid. H 2 SO 3 forms medium (sulfites) and acidic (hydrosulfites) salts. Like sulfur dioxide, sulfurous acid and its salts are strong reducing agents, although in the presence of stronger reducing agents they can exhibit oxidizing properties.

Chemical properties:

1) Participation in the OVR reaction:

Manifestation of the properties of the reducing agent:

2Na 2 S +4 O 3 + O 2 0 = 2Na 2 S +6 O 4 -2

Manifestation of oxidizing properties:

Na 2 SO 3 + S = Na 2 S 2 O 3 - sodium hyposulfite (t 0 C)

2) Thermal decomposition of sulfites:

4Na 2 SO 3 - Na 2 S + 3Na 2 SO 4 (t 0 C)

3) Hydrolysis of water-soluble sulfites - hydrolysis by anion - alkaline medium:

K2SO3 + HOH = KHSO3 + KOH

Sulfuric acid and its salts

Physical properties:

Sulfuric acid (100%) is a colorless heavy oily liquid (“oil of vitriol”); density = 1.84 g/cm 3, t pl = 10.3 0 C, t boil = 296 0 C; odorless, non-volatile, highly soluble in water. Sulfuric acid mixes with water in any ratio. It is very hygroscopic - it actively absorbs water vapor, so it is used as a desiccant both in the laboratory and in everyday life.

Receipt:

Currently, the contact method is used to produce sulfuric acid. This method makes it possible to obtain very pure sulfuric acid of any concentration, as well as oleum.

The starting materials for the production of sulfuric acid can be sulfur, hydrogen sulfide, and metal sulfides. I will consider the production of sulfuric acid by the contact method, in which the starting raw material is iron pyrite FeS 2

Schematic diagram of the production of sulfuric acid.

The process consists of three stages:

	Processes
1. Roasting iron pyrites to produce sulfur oxide (IV). Furnace gas purification.	First stage reaction equation: 4FeS 2 +11O 2 =Fe 2 O 3 +8SO 2 +Q Crushed, purified pyrite is poured into a furnace for firing in a “fluidized bed.” Air enriched with oxygen is passed from below (counterflow principle) for more complete firing of pyrite. The temperature for firing reaches 800 0 C. Furnace gas purification Furnace gas comes out of the furnace, the composition of which is: SO 2, O 2, water vapor and tiny particles of iron oxide. Such furnace gas must be purified from impurities. Furnace gas purification is carried out in two stages - in a cyclone(centrifugal force is used, solid particles fall down) and in emailektrofilters(electrostatic attraction is used, cinder particles stick to the electrified plates of the electrostatic precipitator). Furnace gas is dried in the drying tower- furnace gas rises from bottom to top, and concentrated sulfuric acid pours from top to bottom.
2. OxidationSO 2 VSO 3 oxygen. Leaks in the contact apparatus.	The reaction equation for this stage is: 2SO 2 +O 2 2SO 3 +Q The complexity of the second stage lies in the fact that the process of oxidation of one oxide into another is reversible. Therefore it is necessary to choose optimal conditions for the forward reaction to occur(receiving SO 3): a)temperature: the optimal temperature for a direct reaction to occur with maximum SO 3 formation is temperature 400-500 0 WITH. In order to increase the reaction rate at such a low temperature, a catalyst is introduced into the reaction - vanadium oxide -V 2 O 5 ; b)pressure: a direct reaction occurs with a decrease in gas volumes. The process is carried out at elevated pressure. Heating of the mixture of SO 2 and O 2 to a temperature of 400-500 0 C begins in the heat exchanger. The mixture passes between the heat exchanger tubes and is heated by these tubes. As soon as the mixture of sulfur oxide and oxygen reaches the catalyst layers, the process of oxidation of SO 2 into SO 3 begins. The resulting sulfur oxide SO 3 leaves the contact apparatus and enters the absorption tower through a heat exchanger.
3. ReceiptH 2 SO 4 . Leaks in the absorption tower.	If water is used to absorb sulfur oxide, sulfuric acid is formed in the form of a mist consisting of tiny droplets of sulfuric acid. To prevent the formation of sulfuric acid fog, use 98% concentrated sulfuric acid. Sulfur oxide dissolves very well in such acid, forming oleum: H 2 SO 4 * nSO 3 . The reaction equation for this process is: nSO 3 + H 2 SO 4 = H 2 SO 4 * nSO 3 The resulting oleum is poured into metal tanks and sent to a warehouse. Then tanks are filled with oleum, a train is formed and sent to the consumer.

Chemical properties:

H 2 SO 4 is a strong dibasic acid.

1) Dissociation: in the first stage there is complete dissociation, in the second - sulfuric acid behaves like an acid of medium strength:

H 2 SO 4 H + + HSO4 - (a = 1)

HSO4 - H + + SO4 - (a 1)

2) Interaction with Me:

a) Dilute sulfuric acid:

· Forms the corresponding salts with Me, which are in the voltage series to the left of hydrogen (except for Pb):

H 2 SO 4 + Fe = FeSO 4 + H 2

· Displaces volatile acids from their salts:

H 2 SO 4 + 2NaCl = Na 2 SO 4 + 2HCl

· Displaces weaker acids from their salts:

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

· Interacts with basic and amphoteric oxides:

H 2 SO 4 + CaO = CaSO 4 + H 2 O

· Interacts with bases:

When neutralizing sodium hydroxide with sulfuric acid, the formation of medium and acidic salts is possible, depending on the ratio of the reagents.

When there is an excess of alkali, a medium salt is formed:

2NaOH (g) + H 2 SO 4 = Na 2 SO 4 + 2H 2 O

Full ionic equation:

2Na + + 2OH - + 2H + + SO 4 2- = 2Na + + SO 4 2- + 2H 2 O

When there is an excess of acid, an acid salt is formed:

NaOH + H 2 SO 4 (g) = NaHSO 4 + 2H 2 O

Full ionic equation:

Na + + OH - + H + + HSO 4 - = Na + + HSO 4 - + 2H 2 O

When aluminum hydroxide reacts with sulfuric acid, the formation of medium and basic salts is possible, since Al(OH)3 is a weak base, and H2SO4 is a strong acid. It depends on the ratio of the components.

When there is an excess of sulfuric acid, a medium salt is formed:

2Al(OH) 3 + 3H 2 SO 4 (g) = Al 2 (SO 4) 3 + 6H 2 O

Full ionic equation:

2Al(OH) 3 + 6H + + 3SO 4 2- = 2Al +3 + 3SO 4 2- + 6H 2 O

With a ratio of 1 mol Al(OH)3 to 1 mol H 2 SO 4, a monobasic salt is obtained:

Al(OH)3 + H2SO4 = Al(OH)SO4 + 2H2O

Full ionic equation:

Al(OH) 3 + 2H + + SO 4 2- = Al(OH) 2+ + SO 4 2- + 2H 2 O

With a ratio of 2 mol Al(OH)3 per 1 mol H 2 SO 4, a dibasic salt is obtained:

2Al(OH) 3 + H 2 SO 4 = 2 SO 4 + 2H 2 O

Full ionic equation:

2Al(OH) 3 + 2H + + SO 4 2- = 2 + + SO 4 2- + 2H 2 O

b) Concentrated sulfuric acid is a strong oxidizing agent: when reacting with:

· Inactive Me - reduced to SO 2:

2Ag + 2H2SO4 = Ag2SO4 + SO2 + H2O

Alkaline earth Me and magnesium - up to S:

3Mg + 4H 2 SO 4 = 3MgSO 4 + S + 4H 2 O

· Alkaline Me and zinc - up to H 2 S:

8Na + 5H 2 SO 4 = 4Na 2 SO 4 + H 2 S + 4H 2 O

Al, Fe and Cr are passivated (become inactive due to the formation of a protective film) with concentrated sulfuric acid in the cold (therefore H 2 SO 4 with a concentration above 75% is transported in iron containers), but when heated they are oxidized by it with the formation of sulfates of these metals.

The oxidizing power of dilute and concentrated acids is different:

a) In reactions of dilute sulfuric acid with metals, Me is a reducing agent, and the acid (more precisely, hydrogen) is an oxidizing agent:

Mg + H 2 SO 4 (diluted) = MgSO 4 + H 2

b) In reactions of concentrated sulfuric acid with metals, Me is a reducing agent, and the acid (more precisely, sulfur) is an oxidizing agent:

Mg + 2H 2 SO 4 (conc) = MgSO 4 + SO 2 + 2H 2 O

3) Interaction with HeMe: oxidizes the nonmetal to an acid in the highest oxidation state or to an oxide (if the acid is unstable), and is itself reduced to SO 2:

C + 2H 2 SO 4 (conc) = CO 2 + 2SO 2 + 2H 2 O

S + 2H 2 SO 4 (conc) = 3SO 2 + 2H 2 O

2P + 5H 2 SO 4 (conc) = 5SO 2 + 2H 3 PO 4 + 2H 2 O

4) Concentrated sulfuric acid oxidizes many complex substances:

2KBr + 2H 2 SO 4 (conc) = SO 2 + Br 2 + K 2 SO 4 + 2H 2 O

8KI + 5H 2 SO 4 (conc) = H 2 S + 4I 2 + 4K 2 SO 4 + 4H 2 O

H 2 S + H 2 SO 4 (conc) = SO 2 + S + H 2 O

5) Interaction with bases and amphoteric hydroxides:

H 2 SO 4 + 2NaOH (g) = Na 2 SO 4 + 2H 2 O

H 2 SO 4 (conc) + NaOH = NaHSO 4 + 2H 2 O

H 2 SO 4 + Zn(OH) 2 = ZnSO 4 + 2H 2 O

6) Interaction with basic and amphoteric oxides:

CuO + H 2 SO 4 = CuSO 4 + H 2 O

Al 2 O 3 + 3H 2 SO 4 = Al 2 (SO 4) 3 + 3H 2 O

7) Enters into exchange reactions with medium, acidic and basic salts if a gas, precipitate or slightly dissociating substance is formed:

CaCO 3 + H 2 SO 4 = CaSO 4 + CO 2 + H 2 O

NaHCO 3 + H 2 SO 4 = Na 2 SO 4 + CO 2 + H 2 O

(CuOH) 2 CO 3 + 2H 2 SO 4 = 2CuSO 4 + CO 2 + 3H 2 O

8) Can turn medium salts into acidic ones (or sour ones into more acidic ones):

CaSO 4 + H 2 SO 4 = Ca(HSO 4) 2

CaHPO 4 + H 2 SO 4 = Ca(H 2 PO 4) 2 + CaSO 4

9) 100% sulfuric acid chars organic matter:

C 12 H 22 O 11 (tv) + H 2 SO 4 = 12C (tv) + H 2 SO 4 * 11H 2 O

10) Qualitative reaction to sulfates and sulfuric acid:

BaCl 2 + H 2 SO 4 = BaSO 4 + 2HCL

(formation of a white, acid-insoluble precipitate of barium sulfate)

3. Selenium

Selenium is a chemical element of group 16, period 4 in the periodic table, has atomic number 34, denoted by the symbol Se (lat. Selenium).

Atom

Ordinal number 34, electronic structure: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4.

Valence

Oxidation states

Electronegativity

Physical properties

Brittle, shiny, black non-metal

Allotropic modifications

Solid selenium has several allotropic modifications:

1) Gray selenium (g-Se, “metallic selenium”) is the most stable modification with a hexagonal crystal lattice

2) Red crystalline selenium - three monoclinic modifications: orange-red b-Se, dark red b-Se, red g-Se

3) Red amorphous selenium

4) Black glassy selenium

When gray selenium is heated, it gives a gray melt, and upon further heating it evaporates to form brown vapors. When the vapor is cooled sharply, selenium condenses in the form of a red allotrope.

Being in nature

The selenium content in the earth's crust is about 500 mg/t. The main features of the geochemistry of selenium in the earth's crust are determined by the proximity of its ionic radius to the ionic radius of sulfur. Selenium forms 37 minerals, among which first of all should be noted ashavalite FeSe, clausthalite PbSe, timannite HgSe, guanajuatite Bi 2 (Se, S) 3, hastite CoSe 2, platinite PbBi 2 (S, Se) 3, associated with various sulfides . Native selenium is rarely found

Getting Selenium

Significant quantities of selenium are obtained from copper-electrolyte production sludge, in which selenium is present in the form of silver selenide. There are several methods of obtaining:

1) oxidative roasting with sublimation of SeO 2;

2) heating the sludge with concentrated sulfuric acid, oxidizing selenium compounds to SeO 2 with its subsequent sublimation;

3) oxidative sintering with soda, conversion of the resulting mixture of selenium compounds to Se(IV) compounds and their reduction to elemental selenium by the action of SO 2.

Chemical properties

Selenium is an analogue of sulfur and exhibits oxidation states of 2 (H 2 Se), +4 (SeO 2) and +6 (H 2 SeO 4). However, unlike sulfur, selenium compounds in the +6 oxidation state are the strongest oxidizing agents, and selenium compounds (-2) are much stronger reducing agents than the corresponding sulfur compounds.

The simple substance selenium is much less chemically active than sulfur. Thus, unlike sulfur, selenium is not capable of burning in air on its own. Selenium can be oxidized only with additional heating, during which it slowly burns with a blue flame, turning into SeO 2 dioxide. Selenium reacts (very violently) with alkali metals only when molten.

Unlike SO 2, SeO 2 is not a gas, but a crystalline substance, highly soluble in water. Obtaining selenous acid (SeO 2 + H 2 O > H 2 SeO 3) is no more difficult than sulfurous acid. And by acting on it with a strong oxidizing agent (for example, HClO 3), they obtain selenic acid H 2 SeO 4, almost as strong as sulfuric acid.

Application

One of the most important areas of its technology, production and consumption is the semiconductor properties of both selenium itself and its numerous compounds (selenides), their alloys with other elements in which selenium began to play a key role.

The stable isotope selenium-74 made it possible to create a plasma laser with colossal amplification in the ultraviolet region (about a billion times).

The radioactive isotope selenium-75 is used as a powerful source of gamma radiation for flaw detection.

Potassium selenide together with V 2 O 5 is used in the thermochemical production of hydrogen and oxygen from water (selenium cycle)

4. Tellurium

Tellurium is a chemical element of group 16, period 5 in the periodic table, has atomic number 52; denoted by the symbol Te (lat. Tellurium) .

Atom

Ordinal number 52, electronic structure: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 4.

Valence

Oxidation states

Electronegativity

Being in nature

Physical properties

Tellurium is a brittle silvery-white substance with a metallic luster. In thin layers in the light it is red-brown, in vapors it is golden-yellow. When heated, it becomes plastic. The crystal lattice is hexagonal.

Obtaining tellurium

The main source is sludge from electrolytic refining of copper and lead. The sludge is fired, the tellurium remains in the cinder, which is washed with hydrochloric acid. Tellurium is isolated from the resulting hydrochloric acid solution by passing sulfur dioxide SO 2 through it.

Sulfuric acid is added to separate selenium and tellurium. In this case, tellurium dioxide TeO 2 falls out, and H 2 SeO 3 remains in solution.

Tellurium is reduced from TeO2 oxide with coal.

To purify tellurium from sulfur and selenium, its ability, under the influence of a reducing agent (Al, Zn) in an alkaline medium, to transform into soluble disodium ditelluride Na 2 Te 2 is used:

6Te + 2Al + 8NaOH = 3Na 2 Te 2 + 2Na

To precipitate tellurium, air or oxygen is passed through the solution:

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

To obtain tellurium of special purity, it is chlorinated:

Te + 2Cl2 = TeCl 4

The resulting tetrachloride is purified by distillation or rectification. The tetrachloride is then hydrolyzed with water:

TeCl 4 + 2H 2 O = TeO 2 + 4HCl

and the resulting TeO 2 is reduced with hydrogen:

TeO 2 + H 2 = Te + 2H 2 O

Chemical properties

In chemical compounds, tellurium exhibits an oxidation state of -2; +2; +4; +6. It is an analogue of sulfur and selenium, but is chemically less active than sulfur. It dissolves in alkalis, is susceptible to the action of nitric and sulfuric acids, but is poorly soluble in dilute hydrochloric acid. Tellurium metal begins to react with water at 100 0 C.

With oxygen it forms compounds TeO, TeO 2, TeO 3. In powder form, it oxidizes in air even at room temperature, forming TeO 2 oxide. When heated in air, it burns, forming TeO 2 - a strong compound that is less volatile than tellurium itself. This property is used to purify tellurium from oxides, which are reduced with flowing hydrogen at a temperature of 500-600°C. Tellurium dioxide is poorly soluble in water, but soluble in acidic and alkaline solutions.

In the molten state, tellurium is quite inert, so graphite and quartz are used as container materials when melting it.

Tellurium forms a compound with hydrogen when heated, easily reacts with halogens, and interacts with sulfur and phosphorus and metals. When reacting with concentrated sulfuric acid, it forms sulfite. Forms weak acids: hydrotelluric acid (H 2 Te), telluric acid (H 2 TeO 3) and telluric acid (H 6 TeO 6), most of whose salts are poorly soluble in water.

Application

Tellurium is used in the production of alloys, thermoelectric materials, chalcogenide glasses, and in the production of rubber.

5. Polonium

Polonium is a radioactive chemical element of group 16, period 6 in the periodic table of D.I. Mendeleev, with atomic number 84, is designated by the symbol Po (lat. Polonium)

Atom

Serial number 84, electronic structure:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 10 4f 14 6p 4 .

Valence

Oxidation states

Electronegativity

Physical properties

Under normal conditions it is a soft metal with a silvery-white color.

Isotopes

33 isotopes of polonium are known in the range of mass numbers from 188 to 220. In addition, 10 metastable excited states of polonium isotopes are known. It has no stable isotopes. The longest-lived isotopes are 209 Po and 208 Po.

Receipt

In practice, the polonium nuclide 210 Po is synthesized artificially in gram quantities by irradiating metal 209 Bi with thermal neutrons in nuclear reactors. The resulting 210 Bi turns into 210 Po due to β-decay. When the same isotope of bismuth is irradiated with protons according to the reaction

209 Bi + p > 209 Po + n

the longest-lived isotope of polonium, 209 Po, is formed.

Chemical properties

Polonium metal quickly oxidizes in air. Polonium dioxide (PoO 2) x and polonium monoxide PoO are known. Forms tetrahalides with halogens. When exposed to acids, it goes into solution with the formation of pink Po 2+ cations:

Po + 2HCl = PoCl 2 + H 2

When polonium is dissolved in hydrochloric acid in the presence of magnesium, hydrogen chloride is formed:

Po + Mg + 2HCl = MgCl 2 + H 2 Po

which is in a liquid state at room temperature (from? 36.1 to 35.3 ° C)

In indicator quantities, acidic polonium trioxide PoO 3 and salts of polonium acid, which does not exist in a free state - polonates K 2 PoO 4, were obtained. Forms halides of the composition PoX 2, PoX 4 and PoX 6. Like tellurium, polonium is capable of forming chemical compounds - polonides - with a number of metals.

Polonium is the only chemical element that, at low temperatures, forms a monatomic simple cubic crystal lattice

Application

Polonium-210 in alloys with beryllium and boron is used to manufacture compact and very powerful neutron sources that produce virtually no g-radiation.

Polonium-210 is often used to ionize gases (particularly air).

An important area of application for polonium-210 is its use in the form of alloys with lead, yttrium, or independently for the production of powerful and very compact heat sources for autonomous installations, for example, in space.

Polonium-210 can serve in an alloy with a light isotope of lithium (6 Li) as a substance that can significantly reduce the critical mass of a nuclear charge and serve as a kind of nuclear detonator

6. Livermorium

Livermorium (Latin: Livermorium, Lv), previously known as ununhexium (Latin: Ununhexium, Uuh) and eka-polonium - the 116th chemical element, belongs to the 16th group and the 7th period of the periodic table, atomic number - - 116, the atomic mass of the most stable isotope is 293. An artificially synthesized radioactive element, not found in nature.

Atom

Serial number 116, electronic structure:

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 10 4f 14 6p 6 7s 2 6d 10 5f 14 7p 4 .

Valence

Isotopes

Receipt

Isotopes of livermorium were obtained as a result of nuclear reactions:

and also as a result of the b-decay of 294 Uuo:

Chemical properties

Livermorium is a member of the chalcogen group, where it comes after polonium. However, the chemical properties of livermorium will be significantly different from those of polonium, so separating these elements will not be difficult.

It is assumed that the main and most stable oxidation state for livermorium will be +2. Livermorium will form livermorium oxide with oxygen (LvO), LvHal 2 halides.

With fluorine or under more severe conditions, livermorium will also be able to exhibit an oxidation state of +4 (LvF 4). Livermorium can exhibit this degree of oxidation both in cations and form, like polonium, livermoric acid or its salts - livermorites (or livermorates), for example K 2 LvO 3 - for example, potassium livermorite.

Livermorites, as well as other livermorium compounds with the +4 oxidation state, will exhibit strong oxidizing properties similar to permanganates. Unlike lighter elements, it is assumed that the +6 oxidation state for livermorium will probably be impossible due to the extremely high energy required to depair the 7s 2 electron shell, so the highest oxidation state of livermorium will be +4.

With strong reducing agents (alkali metals or alkaline earth metals), the oxidation state ?2 is also possible (for example, the compound CaLv would be called calcium livermoride). However, livermorides will be very unstable, and exhibit strong reducing properties, since the formation of the Lv 2- anion and the inclusion of two additional electrons is disadvantageous to the main shell of 7p electrons, and the proposed chemistry of livermorium makes the formation of cations much more advantageous than anions.

With hydrogen, the formation of the hydride H 2 Lv is expected, which will be called livermor hydrogen. Very interesting properties are expected for livermorium hydrogen, for example, the possibility of “superhybridization” is assumed - uninvolved 7s 2 electron clouds of livermorium will be able to form an additional mutual bond with each other, and such a bond will somewhat resemble a hydrogen bond, therefore the properties of livermorium hydrogen may differ from the properties of chalcogen hydrogens of lighter analogues . Livermor hydrogen, despite the fact that livermorium will definitely be a metal, will not fully replicate the properties of metal hydrides and will retain a largely covalent character.

The elements of group 6 of the main subgroup are called. Elements of group vi b

Ion radius

Ar

Send your good work in the knowledge base is simple. Use the form below

Similar documents

Password recovery